![General Chemistry 1 PDF [PDF]](https://pdfs.asia/img/200x200/general-chemistry-1-pdf.jpg)
33 0 19 MB
The Commission on Higher Education in collaboration with the Philippine Normal University
Teaching Guide for Senior High School
GENERAL CHEMISTRY 1 SPECIALIZED SUBJECT | ACADEMIC STEM
This Teaching Guide was collaboratively developed and reviewed by educators from public and private schools, colleges, and universities. We encourage teachers and other education stakeholders to email their feedback, comments, and recommendations to the Commission on Higher Education, K to 12 Transition Program Management Unit Senior High School Support Team at [email protected]. We value your feedback and recommendations.
Development Team Team Leader: Wyona C. Patalinghug, Ph.D. Writers: Wyona C. Patalinghug, Ph.D., Vic Marie I. Camacho, Ph.D., Fortunato B. Sevilla III, Ph.D., Maria Cristina D. Singson
Published by the Commission on Higher Education, 2016
Chairperson: Patricia B. Licuanan, Ph.D. Commission on Higher Education
K to 12 Transition Program Management Unit
Office Address: 4th Floor, Commission on Higher Education, C.P. Garcia Ave., Diliman, Quezon City
Telefax: (02) 441-0927 / E-mail Address: [email protected]
Technical Editors: Marlene B. Ferido, Ph.D., Janeth M. Fuentes Copy Reader: Patricia Marie W. Baun Illustrator: Juan Miguel M. Razon, Rachelle Ann J. Bantayan, Danielle Christine Quing Cover Artists: Paolo Kurtis N. Tan, Renan U. Ortiz
Senior High School Support Team
CHED K to 12 Transition Program Management Unit Program Director: Karol Mark R. Yee
Consultants THIS PROJECT WAS DEVELOPED WITH THE PHILIPPINE NORMAL UNIVERSITY.
University President: Ester B. Ogena, Ph.D.
VP for Academics: Ma. Antoinette C. Montealegre, Ph.D.
VP for University Relations & Advancement: Rosemarievic V. Diaz, Ph.D. Ma. Cynthia Rose B. Bautista, Ph.D., CHED
Bienvenido F. Nebres, S.J., Ph.D., Ateneo de Manila University
Carmela C. Oracion, Ph.D., Ateneo de Manila University
Minella C. Alarcon, Ph.D., CHED Gareth Price, Sheffield Hallam University
Stuart Bevins, Ph.D., Sheffield Hallam University
Lead for Senior High School Support:
Gerson M. Abesamis Lead for Policy Advocacy and Communications:
Averill M. Pizarro Course Development Officers:
John Carlo P. Fernando, Danie Son D. Gonzalvo, Stanley Ernest Yu Teacher Training Officers:
Ma. Theresa C. Carlos, Mylene E. Dones Monitoring and Evaluation Officer:
Robert Adrian N. Daulat Administrative Officers:
Ma. Leana Paula B. Bato, Kevin Ross D. Nera, Allison A. Danao, Ayhen Loisse B. Dalena Printed in the Philippines by EC-TEC Commercial, No. 32 St. Louis Compound 7, Baesa, Quezon City, [email protected]
This Teaching Guide by the Commission on Higher Education is licensed under a Creative Commons AttributionNonCommercial-ShareAlike 4.0 International License. This means you are free to: Share — copy and redistribute the material in any medium or format Adapt — remix, transform, and build upon the material. The licensor, CHED, cannot revoke these freedoms as long as you follow the license terms. However, under the following terms: Attribution — You must give appropriate credit, provide a link to the license, and indicate if changes were made. You may do so in any reasonable manner, but not in any way that suggests the licensor endorses you or your use. NonCommercial — You may not use the material for commercial purposes. ShareAlike — If you remix, transform, or build upon the material, you must distribute your contributions under the same license as the original.
Table of Contents DepEd Curriculum Guide
i
Lesson 18: Emission Spectrum of Hydrogen, and Dual
Lesson 1: Matter and Its Properties
1
Nature of Matter
Lesson 2: Matter and Its Various Forms
14
Lesson 19: Flame Test (Laboratory)
158
Lesson 3: Measurements
19
Lesson 20: Electronic Structure of the Atom
162
Lesson 4: Measurements (Laboratory)
25
Lesson 21: Electron Configuration
175
Lesson 5: Atoms, Molecules, and Ions (Lecture)
30
Lesson 22: Periodic Relationships among the Elements
190
Lesson 6: Atoms, Molecules, and Ions (Laboratory)
47
Lesson 23: Periodic Relationships of Main Group
203
Lesson 7: Atomic Mass
52
Elements (Laboratory)
Lesson 8: The Mole Concept and Molar Mass (Lecture)
60
Lesson 24: Ionic Bonds
208
Lecture 9: The Mole Concept and Molar Mass (Laboratory)
70
Lesson 25: Covalent Bonds and Lewis Structures
220
Lesson 10: Percent Composition and Chemical Formulas
76
Lesson 26: Geometry of Molecules and Polarity
236
Lesson 11: Chemical Reactions and Chemical Equations (Lecture)
82
of Compounds
Lesson 12: Chemical Reactions and Chemical Equations (Laboratory)
89
Lesson 27: Geometry of Molecules and Polarity
Lesson 13: Mass Relationships in Chemical Reactions (Lecture)
94
of Molecules (Laboratory)
Lesson 14: Mass Relationships in Chemical Reactions (Laboratory)
105
Lesson 28: Carbon Compounds
254
Lesson 15: Gases (Lecture)
110
Lesson 29: Polymers
284
Lesson 16: Gases (Laboratory)
128
Lesson 30: Biomolecules
297
Lesson 17: Electromagnetic Waves, Planck’s Quantum Theory, and
132
Biographical Notes
314
Additional Images
317
Photoelectric Effect
144
250
Introduction As the Commission supports DepEd’s implementation of Senior High School (SHS), it upholds the vision and mission of the K to 12 program, stated in Section 2 of Republic Act 10533, or the Enhanced Basic Education Act of 2013, that “every graduate of basic education be an empowered individual, through a program rooted on...the competence to engage in work and be productive, the ability to coexist in fruitful harmony with local and global communities, the capability to engage in creative and critical thinking, and the capacity and willingness to transform others and oneself.” To accomplish this, the Commission partnered with the Philippine Normal University (PNU), the National Center for Teacher Education, to develop Teaching Guides for Courses of SHS. Together with PNU, this Teaching Guide was studied and reviewed by education and pedagogy experts, and was enhanced with appropriate methodologies and strategies. Furthermore, the Commission believes that teachers are the most important partners in attaining this goal. Incorporated in this Teaching Guide is a framework that will guide them in creating lessons and assessment tools, support them in facilitating activities and questions, and assist them towards deeper content areas and competencies. Thus, the introduction of the SHS for SHS Framework.
SHS for SHS Framework
The SHS for SHS Framework, which stands for “Saysay-Husay-Sarili for Senior High School,” is at the core of this book. The lessons, which combine high-quality content with flexible elements to accommodate diversity of teachers and environments, promote these three fundamental concepts:
SAYSAY: MEANING
HUSAY: MASTERY
SARILI: OWNERSHIP
Why is this important?
How will I deeply understand this?
What can I do with this?
Through this Teaching Guide, teachers will be able to facilitate an understanding of the value of the lessons, for each learner to fully engage in the content on both the cognitive and affective levels.
Given that developing mastery goes beyond memorization, teachers should also aim for deep understanding of the subject matter where they lead learners to analyze and synthesize knowledge.
When teachers empower learners to take ownership of their learning, they develop independence and selfdirection, learning about both the subject matter and themselves.
Parts of the
Teaching Guide
This Teaching Guide is mapped and aligned to the DepEd SHS Curriculum, designed to be highly usable for teachers. It contains classroom activities and pedagogical notes, and is integrated with innovative pedagogies. All of these elements are presented in the following parts: 1. Introduction • Highlight key concepts and identify the essential questions • Show the big picture • Connect and/or review prerequisite knowledge • Clearly communicate learning competencies and objectives • Motivate through applications and connections to real-life 2. Motivation • Give local examples and applications • Engage in a game or movement activity • Provide a hands-on/laboratory activity • Connect to a real-life problem 3. Instruction/Delivery • Give a demonstration/lecture/simulation/hands-on activity • Show step-by-step solutions to sample problems • Give applications of the theory • Connect to a real-life problem if applicable 4. Practice • Discuss worked-out examples • Provide easy-medium-hard questions • Give time for hands-on unguided classroom work and discovery • Use formative assessment to give feedback 5. Enrichment • Provide additional examples and applications • Introduce extensions or generalisations of concepts • Engage in reflection questions • Encourage analysis through higher order thinking prompts 6. Evaluation • Supply a diverse question bank for written work and exercises • Provide alternative formats for student work: written homework, journal, portfolio, group individual projects, student-directed research project
iii
On DepEd Functional Skills and CHED College Readiness Standards As Higher Education Institutions (HEIs) welcome the graduates of the Senior High School program, it is of paramount importance to align Functional Skills set by DepEd with the College Readiness Standards stated by CHED.
On the other hand, the Commission declared the College Readiness Standards that consist of the combination of knowledge, skills, and reflective thinking necessary to participate and succeed without remediation - in entry-level undergraduate courses in college.
The DepEd articulated a set of 21st century skills that should be embedded in the SHS curriculum across various subjects and tracks. These skills are desired outcomes that K to 12 graduates should possess in order to proceed to either higher education, employment, entrepreneurship, or middle-level skills development.
The alignment of both standards, shown below, is also presented in this Teaching Guide - prepares Senior High School graduates to the revised college curriculum which will initially be implemented by AY 2018-2019.
College Readiness Standards Foundational Skills
DepEd Functional Skills
Produce all forms of texts (written, oral, visual, digital) based on: 1. 2. 3. 4. 5.
Solid grounding on Philippine experience and culture; An understanding of the self, community, and nation; Visual and information literacies, media literacy, critical thinking Application of critical and creative thinking and doing processes; and problem solving skills, creativity, initiative and self-direction Competency in formulating ideas/arguments logically, scientifically, and creatively; and Clear appreciation of one’s responsibility as a citizen of a multicultural Philippines and a diverse world;
Systematically apply knowledge, understanding, theory, and skills for the development of the self, local, and global communities using prior learning, inquiry, and experimentation
Global awareness, scientific and economic literacy, curiosity, critical thinking and problem solving skills, risk taking, flexibility and adaptability, initiative and self-direction
Work comfortably with relevant technologies and develop adaptations and innovations for significant use in local and global communities
Global awareness, media literacy, technological literacy, creativity, flexibility and adaptability, productivity and accountability
Communicate with local and global communities with proficiency, orally, in writing, and through new technologies of communication
Global awareness, multicultural literacy, collaboration and interpersonal skills, social and cross-cultural skills, leadership and responsibility
Interact meaningfully in a social setting and contribute to the fulfilment of individual and shared goals, respecting the fundamental humanity of all persons and the diversity of groups and communities
Media literacy, multicultural literacy, global awareness, collaboration and interpersonal skills, social and cross-cultural skills, leadership and responsibility, ethical, moral, and spiritual values
v
K to 12 BASIC EDUCATION CURRICULUM SENIOR HIGH SCHOOL – SCIENCE, TECHNOLOGY, ENGINEERING AND MATHEMATICS (STEM) SPECIALIZED SUBJECT
Grade: 11 Subject Title: General Chemistry 1 & 2
Semester: 1st and 2nd No. of Hours/ Semester: 80 hours per semester
Subject Description: Composition, structure, and properties of matter; quantitative principles, kinetics, and energetics of transformations of matter; and fundamental concepts of organic chemistry CONTENT Quarter 1 – General Chemistry Matter and its properties 1. the particulate nature of matter 2. states of matter a. the macroscopic b. microscopic view 3. Physical and chemical properties 4. Extensive and intensive properties 5. Ways of classifying matter a. pure substances and mixtures b. elements and compounds c. homogeneous and heterogeneous mixtures 6. Methods of separating mixtures into their component substances
CONTENT STANDARD
PERFORMANCE STANDARD
LEARNING COMPETENCIES
CODE
1
The learners demonstrate an understanding of: the properties of matter and its various forms
The learners: design using multimedia, demonstrations, or models, a representation or simulation of any of the following: a. atomic structure b. gas behavior c. mass relationships in d. reactions
K to 12 Senior High School STEM Specialized Subject – General Chemistry 1 and 2 December 2013
The learners:
1. recognize that substances are made up of smaller particles 2. describe and/or make a representation of the arrangement, relative spacing, and relative motion of the particles in each of the three phases of matter 3. distinguish between physical and chemical properties and give examples 4. distinguish between extensive and intensive properties and give examples 5. use properties of matter to identify substances and to separate them 6. differentiate between pure substances and mixtures 7. differentiate between elements and compounds 8. differentiate between homogenous and heterogenous mixtures 9. recognize the formulas of common chemical substances 10. describe separation techniques for mixtures and compounds 11. compare consumer products on the basis of their components for use, safety, quality and cost 12. (LAB) apply simple separation techniques such as distillation, chromatography
STEM_GC11MP-Ia-b-1
STEM_GC11MP-Ia-b-2 STEM_GC11MP-Ia-b-3 STEM_GC11MP-Ia-b-4 STEM_GC11MP-Ia-b-5 STEM_GC11MP-Ia-b-6 STEM_GC11MP-Ia-b-7 STEM_GC11MP-Ia-b-8 STEM_GC11MP-Ia-b-9 STEM_GC11MP-Ia-b-10 STEM_GC11MP-Ia-b-11 STEM_GC11MP-Ia-b-12 Page 1 of 17
K to 12 BASIC EDUCATION CURRICULUM SENIOR HIGH SCHOOL – SCIENCE, TECHNOLOGY, ENGINEERING AND MATHEMATICS (STEM) SPECIALIZED SUBJECT
CONTENT
CONTENT STANDARD
Measurements 1. Accuracy and precision 2. Significant figures in calculations 3. Density measurement
1. the difference between accuracy and precision 2. different sources of errors in measurements
Atoms, Molecules, and Ions 1. Dalton’s atomic theory 2. Basic laws of matter 3. Atomic structure 4. Subatomic particles (protons, electrons, neutrons) 5. Molecules and Ions 6. Chemical Formulas 7. Naming Compounds
1. atomic structure 2. formulas and names of compounds
PERFORMANCE STANDARD
K to 12 Senior High School STEM Specialized Subject – General Chemistry 1 and 2 December 2013
LEARNING COMPETENCIES
CODE
1. differentiate between precision and accuracy
STEM_GC11MT-Ib-13
2. (LAB) Determine the density of liquids & solids
STEM_GC11MT-Ib-14
1. explain how the basic laws of matter (law of conservation of mass, law of constant composition, law of multiple proportion) led to the formulation of Dalton’s Atomic Theory
STEM_GC11AM-Ic-e-15
2. describe Dalton’s Atomic Theory
STEM_GC11AM-Ic-e-16
3. differentiate among atomic number, mass number, and isotopes, and which of these distinguishes one element from another
STEM_GC11AM-Ic-e-17
4. write isotopic symbols
STEM_GC11AM-Ic-e-18
5. recognize common isotopes and their uses.
STEM_GC11AM-Ic-e-19
6. differentiate among atoms, molecules, ions and give examples
STEM_GC11AM-Ic-e-20
7. represent compounds using chemical formulas, structural formulas and models
STEM_GC11AM-Ic-e-21
8. give the similarities and differences between the empirical formula and molecular formula of a compound
STEM_GC11AM-Ic-e-22 Page 2 of 17
K to 12 BASIC EDUCATION CURRICULUM SENIOR HIGH SCHOOL – SCIENCE, TECHNOLOGY, ENGINEERING AND MATHEMATICS (STEM) SPECIALIZED SUBJECT
CONTENT
CONTENT STANDARD
PERFORMANCE STANDARD
LEARNING COMPETENCIES
CODE
9. name compounds given their formula and write formula given the name of the compound
STEM_GC11AM-Ic-e-23
10. (LAB) Practice chemical nomenclature: writing the chemical formulas of ionic compounds; naming ionic compounds from formulas Stoichiometry 1. Atomic mass 2. Avogadro’s number 3. The mole concept
4. Percent composition and chemical formulas
1. the mole concept in relation to Avogadro’s number and mass
2. the relationship of percent composition and chemical formula
K to 12 Senior High School STEM Specialized Subject – General Chemistry 1 and 2 December 2013
STEM_GC11AM-Ic-e-24
1. explain relative atomic mass and average atomic mass
STEM_GC11S-Ie-25
2. define a mole
STEM_GC11S-Ie-26
3. illustrate Avogadro’s number with examples
STEM_GC11S-Ie-27
4. determine the molar mass of elements and compounds
STEM_GC11S-Ie-28
5. calculate the mass of a given number of moles of an element or compound or vice versa
STEM_GC11S-Ie-29
6. calculate the mass of a given number of particles of an element or compound or vice versa
STEM_GC11S-Ie-30
1. calculate the percent composition of a compound from its formula
STEM_GC11PC-If-31
2. calculate the empirical formula from the percent composition of a compound
STEM_GC11PC-If-32
Page 3 of 17
K to 12 BASIC EDUCATION CURRICULUM SENIOR HIGH SCHOOL – SCIENCE, TECHNOLOGY, ENGINEERING AND MATHEMATICS (STEM) SPECIALIZED SUBJECT
CONTENT
CONTENT STANDARD
PERFORMANCE STANDARD
LEARNING COMPETENCIES 3. calculate molecular formula given molar mass
5. Chemical reactions and chemical equations 6. Types of chemical reactions in aqueous solutions
3. the use of chemical formulas to represent chemical reactions
4. write equations for chemical reactions and balance the equations 5. interpret the meaning of a balanced chemical reaction in terms of the law of conservation of mass 6. describe evidences that a chemical reaction has occurred 7. (LAB) Perform exercises on writing and balancing chemical equations
7. Mass relationships in chemical reactions
Gases 1. Pressure of a gas a. Units of pressure 2. The Gas laws
4. the quantitative relationship of reactants and products in a chemical reaction
5. the mathematical relationship between pressure, volume, and temperature of
K to 12 Senior High School STEM Specialized Subject – General Chemistry 1 and 2 December 2013
1. construct mole or mass ratios for a reaction in order to calculate the amount of reactant needed or amount of product formed in terms of moles or mass 2. Calculate percent yield and theoretical yield of the reaction 3. explain the concept of limiting reagent in a chemical reaction; identify the excess reagent(s) 4. calculate reaction yield when a limiting reagent is present 5. (LAB) Determine mass relationship in a chemical reaction 1. define pressure and give the common units of pressure 2.
express the gas laws in equation form
CODE
STEM_GC11PC-If-33
STEM_GC11CR-If-g-34 STEM_GC11CR-If-g-35 STEM_GC11CR-If-g-36
STEM_GC11CR-If-g-37
STEM_GC11MR-Ig-h-38
STEM_GC11MR-Ig-h-39 STEM_GC11MR-Ig-h-40 STEM_GC11MR-Ig-h-41 STEM_GC11MR-Ig-h-42 STEM_GC11G-Ih-i-43 STEM_GC11G-Ih-i-44 Page 4 of 17
K to 12 BASIC EDUCATION CURRICULUM SENIOR HIGH SCHOOL – SCIENCE, TECHNOLOGY, ENGINEERING AND MATHEMATICS (STEM) SPECIALIZED SUBJECT
CONTENT
CONTENT STANDARD
a. Boyle’s Law b. Charles’ Law c. Avogadro’s Law 3. Ideal Gas Equation
a gas
4. Dalton’s Law of partial pressures
6. the partial pressures of gases in a mixture
5. Gas stoichiometry
7. quantitative relationships of reactants and products in a gaseous reaction
6. Kinetic molecular theory of gases
PERFORMANCE STANDARD
LEARNING COMPETENCIES use the gas laws to determine pressure, volume, or temperature of a gas under certain conditions of change 4. use the ideal gas equation to calculate pressure, volume, temperature, or number of moles of a gas 5. use Dalton’s law of partial pressures to relate mole fraction and partial pressure of gases in a mixture
CODE
3.
8. the behavior and properties of gases at the molecular level
STEM_GC11G-Ih-i-45 STEM_GC11G-Ih-i-46
STEM_GC11DL-Ii-47
6. apply the principles of stoichiometry to determine the amounts (volume, number of moles, or mass) of gaseous reactants and products
STEM_GC11GS-Ii-j-48
7. explain the gas laws in terms of the kinetic molecular theory of gases
STEM_GC11KMT-Ij-49
8. relate the rate of gas effusion with molar mass
STEM_GC11KMT-Ij-50
9. (LAB) Demonstrate Graham’s law of effusion in an experiment
STEM_GC11KMT-Ij-51
Quarter 2 – General Chemistry 1 Electronic Structure of Atoms 1. Quantum mechanical description of the atom 2. Schrodinger’s model of the hydrogen atom and wave functions 3. Main energy levels, sublevels and orbitals
the quantum mechanical description of the atom and its electronic structure
illustrate the reactions at the molecular level in any of the following: 1. enzyme action 2. protein denaturation 3. separation of components in coconut milk
K to 12 Senior High School STEM Specialized Subject – General Chemistry 1 and 2 December 2013
1. describe the quantum mechanical model of the atom 2. describe the electronic structure of atoms in terms of main energy levels, sublevels, and orbitals, and relate this to energy 3. use quantum numbers to describe an electron in an atom 4. (LAB) Perform exercises on quantum numbers
STEM_GC11ES-IIa-b-52 STEM_GC11ES-IIa-b-53 STEM_GC11ES-IIa-b-54 STEM_GC11ES-IIa-b-55
Page 5 of 17
K to 12 BASIC EDUCATION CURRICULUM SENIOR HIGH SCHOOL – SCIENCE, TECHNOLOGY, ENGINEERING AND MATHEMATICS (STEM) SPECIALIZED SUBJECT
CONTENT
CONTENT STANDARD
PERFORMANCE STANDARD
4. Quantum numbers 5. Electron Configuration a. Aufbau Principle b. Pauli Exclusion Principle c. Hund’s Rule d. Diamagnetism and Paramagnetism e. Orbital diagrams Electronic Structure and Periodicity 1. The Electron Configuration and the Periodic Table 2. Periodic Variation in Atomic Properties a. Atomic Radius and effective nuclear charge; the shielding effect in many-electron atoms b. Ionic radius c. Ionization energy d. Electron affinity
Chemical Bonding Ionic Bonds 1. The stability of noble gases 2. Forming ions 3. Ionic bonding 4. Ionic compounds 5. Formulas 6. Structure 7. Properties
the arrangement of elements in the periodic table and trends in the properties of the elements in terms of electronic structure
1. ionic bond formation in terms of atomic properties 2. the properties of ionic compounds in relation to their structure
K to 12 Senior High School STEM Specialized Subject – General Chemistry 1 and 2 December 2013
LEARNING COMPETENCIES
CODE
5. write the electronic configuration of atoms
STEM_GC11ES-IIa-b-56
6. determine the magnetic property of the atom based on its electronic configuration 7. draw an orbital diagram to represent the electronic configuration of atoms 8. (LAB) Perform exercises on writing electronic configuration 1. explain the periodic recurrence of similar properties among elements in the periodic table in terms of electronic structure 2. relate the number of valence electrons of elements to their group number in the periodic table 3. compare the properties of families of elements 4. predict the properties of individual elements based on their position in the periodic table 5. describe and explain the trends in atomic properties in the periodic table 6. (LAB) Investigate reactions of ions and apply these in qualitative analysis 7. (LAB) Determine periodic properties of the main group elements 1. relate the stability of noble gases to their electron configuration
STEM_GC11ES-IIa-b-57 STEM_GC11ES-IIa-b-58 STEM_GC11ES-IIa-b-59 STEM_GC11ESP-IIc-d-60 STEM_GC11ESP-IIc-d-61 STEM_GC11ESP-IIc-d-62 STEM_GC11ESP-IIc-d-63 STEM_GC11ESP-IIc-d-64 STEM_GC11ESP-IIc-d-65 STEM_GC11ESP-IIc-d-66 STEM_GC11CB-IId-g-67
2. state the octet rule
STEM_GC11CB-IId-g-68
3. determine the charge of the ions formed by the representative elements and relate this to their ionization energy or electron affinity, valence electron configuration and position in the periodic table
STEM_GC11CB-IId-g-69
Page 6 of 17
K to 12 BASIC EDUCATION CURRICULUM SENIOR HIGH SCHOOL – SCIENCE, TECHNOLOGY, ENGINEERING AND MATHEMATICS (STEM) SPECIALIZED SUBJECT
CONTENT
Covalent Bonds 1. Formation of covalent bonds 2. Formulas of molecular compounds 3. Lewis structure of molecules 4. Molecules of elements 5. Molecules of compounds 6. Structure and properties of molecular compounds 7. Strength of covalent bonds 8. Electronegativity and bond polarity 9. Geometry of molecules 10. Polarity of compounds
CONTENT STANDARD
PERFORMANCE STANDARD
1. covalent bond formation in terms of atomic properties 2. the properties of molecular covalent compounds in relation to their structure
K to 12 Senior High School STEM Specialized Subject – General Chemistry 1 and 2 December 2013
LEARNING COMPETENCIES
CODE
4. draw the Lewis structure of ions
STEM_GC11CB-IId-g-70
5. predict the formula of the ionic compound formed by a metal and non-metal among the representative elements
STEM_GC11CB-IId-g-71
6. Lewis structure of ionic compounds
STEM_GC11CB-IId-g-72
7. list the properties of ionic compounds and explain these properties in terms of their structure 8. (LAB) Perform exercises on writing Lewis structures of ions/ionic compounds and molecules 9. describe covalent bonding in terms of electron sharing 10. apply the octet rule in the formation of molecular covalent compounds 11. write the formula of molecular compounds formed by the nonmetallic elements of the representative block 12. draw Lewis structure of molecular covalent compounds 13. explain the properties of covalent molecular compounds in terms of their structure. 14. determine the polarity of a bond based on the electronegativities of the atoms forming the bond 15. describe the geometry of simple compounds
STEM_GC11CB-IId-g-73 STEM_GC11CB-IId-g-74 STEM_GC11CB-IId-g-75 STEM_GC11CB-IId-g-76 STEM_GC11CB-IId-g-77 STEM_GC11CB-IId-g-78 STEM_GC11CB-IId-g-79 STEM_GC11CB-IId-g-80 STEM_GC11CB-IId-g-81
16. determine the polarity of simple molecules
STEM_GC11CB-IId-g-82
17. (LAB) Determine and/or observe evidence of molecular polarity
STEM_GC11CB-IId-g-83 Page 7 of 17
K to 12 BASIC EDUCATION CURRICULUM SENIOR HIGH SCHOOL – SCIENCE, TECHNOLOGY, ENGINEERING AND MATHEMATICS (STEM) SPECIALIZED SUBJECT
CONTENT Organic compounds 1. The carbon atom 2. Bonding patterns in hydrocarbons 3. Properties and reactivities of common functional groups 4. Polymers 5. Biomolecules
CONTENT STANDARD
PERFORMANCE STANDARD
the properties of organic compounds and polymers in terms of their structure
LEARNING COMPETENCIES 1.
describe the special nature of carbon
2.
list general characteristics of organic compounds describe the bonding in ethane, ethene(ethylene) and ethyne(acetylene) and explain their geometry in terms of hybridization and σ and ¶ carbon-carbon bonds describe the different functional groups cite uses of representative examples of compounds bearing the different functional groups describe structural isomerism; give examples describe some simple reactions of organic compounds: combustion of organic fuels, addition, condensation, and saponification of fats describe the formation and structure of polymers
3.
4. 5.
6. 7.
8. 9.
STEM_GC11OC-IIg-j-84 STEM_GC11OC-IIg-j-85
STEM_GC11OC-IIg-j-86
STEM_GC11OC-IIg-j-87 STEM_GC11OC-IIg-j-88 STEM_GC11OC-IIg-j-89
STEM_GC11OC-IIg-j-90
STEM_GC11OC-IIg-j-91
give examples of polymers
STEM_GC11OC-IIg-j-92
10.
explain the properties of some polymers in terms of their structure
STEM_GC11OC-IIg-j-93
11.
describe some biomolecules: proteins, nucleic acids, lipids, and carbohydrates describe the structure of proteins, nucleic acids, lipids, and carbohydrates, and relate them to their function
12.
K to 12 Senior High School STEM Specialized Subject – General Chemistry 1 and 2 December 2013
CODE
STEM_GC11OC-IIg-j-94
STEM_GC11OC-IIg-j-95
Page 8 of 17
K to 12 BASIC EDUCATION CURRICULUM SENIOR HIGH SCHOOL – SCIENCE, TECHNOLOGY, ENGINEERING AND MATHEMATICS (STEM) SPECIALIZED SUBJECT
CONTENT
CONTENT STANDARD
PERFORMANCE STANDARD
LEARNING COMPETENCIES 13. 14. 15.
Third Quarter – General Chemistry 2 Intermolecular Forces and 1. the properties of Liquids and Solids liquids and solids to 1. Kinetic molecular model of the nature of forces liquids and solids between particles 2. Intermolecular Forces 2. phase changes in 3. Dipole-dipole forces terms of the 4. Ion-dipole forces accompanying 5. Dispersion forces changes in energy 6. Hydrogen bonds and forces between 7. Properties of liquids and particles IMF 8. Surface Tension 9. Viscosity 10. Vapour pressure, boiling point 11. Molar heat of vaporization 12. Structure and Properties of Water 13. Types and properties of solids 14. Crystalline and amorphous solids 15. Types of Crystals – ionic,
design a simple investigation to determine the effect on boiling point or freezing point when a solid is dissolved in water
K to 12 Senior High School STEM Specialized Subject – General Chemistry 1 and 2 December 2013
(LAB) Perform exercises on the structure of organic compounds using of models (LAB) Prepare selected organic compound and describe their properties (LAB) Perform laboratory activities on enzyme action, protein denaturation, separation of components in coconut milk
CODE STEM_GC11OC-IIg-j-96 STEM_GC11OC-IIg-j-97
STEM_GC11OC-IIg-j-98
1.
use the kinetic molecular model to explain properties of liquids and solids
STEM_GC11IMF-IIIa-c-99
2.
describe and differentiate the types of intermolecular forces
STEM_GC11IMF-IIIa-c100
3.
predict the intermolecular forces possible for a molecule
STEM_GC11IMF-IIIa-c101
4.
describe the following properties of liquids, and explain the effect of intermolecular forces on these properties: surface tension, viscosity, vapor pressure, boiling point, and molar heat of vaporization explain the properties of water with its molecular structure and intermolecular forces
5.
STEM_GC11IMF-IIIa-c102
STEM_GC11IMF-IIIa-c103
6.
describe the difference in structure of crystalline and amorphous solids
STEM_GC11IMF-IIIa-c104
7.
describe the different types of crystals and their properties: ionic, covalent, molecular, and metallic.
STEM_GC11IMF-IIIa-c105 Page 9 of 17
K to 12 BASIC EDUCATION CURRICULUM SENIOR HIGH SCHOOL – SCIENCE, TECHNOLOGY, ENGINEERING AND MATHEMATICS (STEM) SPECIALIZED SUBJECT
CONTENT
CONTENT STANDARD
PERFORMANCE STANDARD
covalent, molecular, metallic 16. Phase Changes - phase diagrams of water and carbon dioxide
LEARNING COMPETENCIES 8.
9. 10. 11. Physical Properties of Solutions 1. Types of Solutions 2. Energy of solution formation 3. Concentration Units and comparison of concentration units a. percent by mass, by volume b. mole fraction c. molality d. molarity e. percent by volume, percent by mass, ppm 4. Solution stoichiometry 5. Factors affecting Solubility 6. Colligative Properties of Nonelectrolyte and electrolyte solutions
properties of solutions, solubility, and the stoichiometry of reactions in solutions
K to 12 Senior High School STEM Specialized Subject – General Chemistry 1 and 2 December 2013
CODE
describe the nature of the following phase changes in terms of energy change and the increase or decrease in molecular order: solid-liquid, liquidvapor, and solid-vapor
STEM_GC11IMF-IIIa-c106
interpret the phase diagram of water and carbon dioxide
STEM_GC11IMF-IIIa-c107
(LAB) Measure and explain the difference in the viscosity of some liquids (LAB) Determine and explain the heating and cooling curve of a substance
STEM_GC11IMF-IIIa-c108 STEM_GC11IMF-IIIa-c109
1.
describe the different types of solutions
STEM_GC11PP-IIId-f-110
2.
use different ways of expressing concentration of solutions: percent by mass, mole fraction, molarity, molality, percent by volume, percent by mass, ppm
STEM_GC11PP-IIId-f-111
3.
perform stoichiometric calculations for reactions in solution
STEM_GC11PP-IIId-f-112
4.
explain the effect of temperature on the solubility of a solid and of a gas
STEM_GC11PP-IIId-f-113
5.
explain the effect of pressure on the solubility of a gas
STEM_GC11PP-IIId-f-114
6.
describe the effect of concentration on the colligative properties of solutions
STEM_GC11PP-IIId-f-115
7.
differentiate the colligative properties of nonelectrolyte solutions and of electrolyte solutions
STEM_GC11PP-IIId-f-116 Page 10 of 17
K to 12 BASIC EDUCATION CURRICULUM SENIOR HIGH SCHOOL – SCIENCE, TECHNOLOGY, ENGINEERING AND MATHEMATICS (STEM) SPECIALIZED SUBJECT
CONTENT
CONTENT STANDARD
PERFORMANCE STANDARD
LEARNING COMPETENCIES 8.
Calculate boiling point elevation and freezing point depression from the concentration of a solute in a solution
STEM_GC11PP-IIId-f-117
9.
calculate molar mass from colligative property data
STEM_GC11PP-IIId-f-118
10.
(LAB) Perform acid-base titration to determine concentration of solutions
STEM_GC11PP-IIId-f-119
11.
(LAB) Determine the solubility of a solid in a given amount of water at different temperatures (LAB) Determine the molar mass of a solid from the change of melting point or boiling point of a solution explain the energy changes during chemical reactions distinguish between exothermic and endothermic processes explain the first law of thermodynamics
12. Thermochemistry 1. Energy Changes in Chemical Reactions: exothermic and endothermic processes 2. First Law of Thermodynamics 3. Enthalpy of a Chemical Reaction - thermochemical equations 4. Calorimetry 5. Standard Enthalpy of Formation and Reaction Hess’ Law
energy changes in chemical reactions
1. 2. 3. 4.
explain enthalpy of a reaction.
5.
Write the thermochemical equation for a chemical reaction Calculate the change in enthalpy of a given reaction using Hess Law (LAB) Do exercises on thermochemical calculations (LAB)Determine the heat of neutralization of an acid describe how various factors influence the rate of a reaction write the mathematical relationship between the rate of a reaction, rate constant, and concentration of the
6. 7. 8.
Chemical Kinetics 1. The rate of a 1. The Rate of a Reaction reaction and the 2. Factors that influence various factors that reaction rate influence it 3. The Rate Law and its 2. the collision theory K to 12 Senior High School STEM Specialized Subject – General Chemistry 1 and 2 December 2013
CODE
1. 2.
STEM_GC11PP-IIId-f-120 STEM_GC11PP-IIId-f-121 STEM_GC11TC-IIIg-i-122 STEM_GC11TC-IIIg-i-123 STEM_GC11TC-IIIg-i-124 STEM_GC11TC-IIIg-i-125 STEM_GC11TC-IIIg-i-126 STEM_GC11TC-IIIg-i-127 STEM_GC11TC-IIIg-i-128 STEM_GC11TC-IIIg-i-129 STEM_GC11CK-IIIi-j-130 STEM_GC11CK-IIIi-j-131 Page 11 of 17
K to 12 BASIC EDUCATION CURRICULUM SENIOR HIGH SCHOOL – SCIENCE, TECHNOLOGY, ENGINEERING AND MATHEMATICS (STEM) SPECIALIZED SUBJECT
CONTENT 4. 5.
CONTENT STANDARD
PERFORMANCE STANDARD
components Collision theory Catalysis
LEARNING COMPETENCIES 3. 4. 5. 6. 7. 8. 9. 10.
Fourth Quarter – General Chemistry 2 Chemical Thermodynamics spontaneous change, 1. Spontaneous processes entropy, and free energy 2. Entropy 3. The Second Law of Thermodynamics 4. Gibbs Free Energy and Chemical Equilibrium
Chemical Equilibrium 1. The equilibrium condition
Chemical equilibrium and Le Chatelier’s
prepare a poster on a specific application of one of the following: a. Acid-base equilibrium b. Electrochemistry Include in the poster the concepts, principles, and chemical reactions involved, and diagrams of processes and other relevant materials
K to 12 Senior High School STEM Specialized Subject – General Chemistry 1 and 2 December 2013
reactants differentiate zero, first-, and secondorder reactions write the rate law for first-order reaction discuss the effect of reactant concentration on the half-time of a first-order reaction explain the effect of temperature on the rate of a reaction explain reactions qualitatively in terms of molecular collisions explain activation energy and how a catalyst affects the reaction rate cite and differentiate the types of catalysts (LAB)Determine the effect of various factors on the rate of a reaction
CODE
STEM_GC11CK-IIIi-j-132 STEM_GC11CK-IIIi-j-133 STEM_GC11CK-IIIi-j-134 STEM_GC11CK-IIIi-j-135 STEM_GC11CK-IIIi-j-136 STEM_GC11CK-IIIi-j-137 STEM_GC11CK-IIIi-j-138 STEM_GC11CK-IIIi-j-139
1. predict the spontaneity of a process based on entropy
STEM_GC11CT-IVa-b-140
2. determine whether entropy increases or decreases if the following are changed: temperature, phase, number of particles
STEM_GC11CT-IVa-b-141
3. explain the second law of thermodynamics and its significance
STEM_GC11CT-IVa-b-142
4. use Gibbs’ free energy to determine the direction of a reaction
STEM_GC11CT-IVa-b-143
1. describe reversible reactions
STEM_GC11CE-IVb-e-144
Page 12 of 17
K to 12 BASIC EDUCATION CURRICULUM SENIOR HIGH SCHOOL – SCIENCE, TECHNOLOGY, ENGINEERING AND MATHEMATICS (STEM) SPECIALIZED SUBJECT
CONTENT 2. Writing the reaction quotient/equilibrium constant expression 3. Predicting the direction of a reaction 4. Significance of the equilibrium constant 5. Le Chatelier’s Principle
CONTENT STANDARD
PERFORMANCE STANDARD
Principle
LEARNING COMPETENCIES
CODE
2. explain chemical equilibrium in terms of the reaction rates of the forward and the reverse reaction
STEM_GC11CE-IVb-e-145
3. write expressions for the reaction quotient/equilibrium constants
STEM_GC11CE-IVb-e-146
4. explain the significance of the value of the equilibrium constant.
STEM_GC11CE-IVb-e-147
5. calculate equilibrium constant and the pressure or concentration of reactants or products in an equilibrium mixture 6. state the Le Chatelier’s principle and apply it qualitatively to describe the effect of changes in pressure, concentration and temperature on a system at equilibrium
Acid-Base Equilibria and Salt Equilibria 1. Bronsted acids and bases 2. The acid-base properties of water 3. pH- a measure of acidity 4. Strength of acids and bases 5. Weak acids/weak bases and
1. acid-base equilibrium and its applications to the pH of solutions and the use of buffer solutions 2. solubility equilibrium and its applications
K to 12 Senior High School STEM Specialized Subject – General Chemistry 1 and 2 December 2013
STEM_GC11CE-IVb-e-148
STEM_GC11CE-IVb-e-149
7. (LAB) Describe the behavior of reversible reactions
STEM_GC11CE-IVb-e-150
8. (LAB) Describe the behavior of a reaction mixture when the following takes place: a. change in concentration of reactants or products b. change in temperature
STEM_GC11CE-IVb-e-151
9. (LAB) Perform calculations involving equilibrium of gaseous reactions
STEM_GC11CE-IVb-e-152
1. define Bronsted acids and bases
STEM_GC11AB-IVf-g-153
2. discuss the acid-base property of water
STEM_GC11AB-IVf-g-154
3. define pH
STEM_GC11AB-IVf-g-155
4. calculate pH from the concentration of hydrogen ion or hydroxide ions in aqueous solutions
STEM_GC11AB-IVf-g-156
Page 13 of 17
K to 12 BASIC EDUCATION CURRICULUM SENIOR HIGH SCHOOL – SCIENCE, TECHNOLOGY, ENGINEERING AND MATHEMATICS (STEM) SPECIALIZED SUBJECT
CONTENT
CONTENT STANDARD
PERFORMANCE STANDARD
ionization constants 6. Relationship between the ionization constants of acids and their conjugate bases 7. The Common Ion Effect 8. Buffer solutions 9. Solubility equilibria
Electrochemistry 1. Redox reactions 2. Galvanic cells 3. Standard reduction potentials 4. Spontaneity of redox reactions 5. Batteries 6. Corrosion 7. Electrolysis
LEARNING COMPETENCIES 5. determine the relative strength of an acid or a base, from the value of the ionization constant of a weak acid or base 6. determine the pH of a solution of weak acid or weak base 7. explain the Common Ion Effect
Redox reactions as applied to galvanic and electrolytic cells
K to 12 Senior High School STEM Specialized Subject – General Chemistry 1 and 2 December 2013
8. describe how a buffer solution maintains its pH 9. calculate the pH of a buffer solution using the Henderson-Hasselbalch equation 10. explain and apply the solubility product constant to predict the solubility of salts 11. describe the common ion effect on the solubility of a precipitate 12. explain the effect of pH on the solubility of a precipitate 13. (LAB) Determine the pH of solutions of a weak acid at different concentrations and in the presence of its salt 14. (LAB)Determine the behavior of the pH of buffered solutions upon the addition of a small amount of acid and base 1. define oxidation and reduction reactions
CODE STEM_GC11AB-IVf-g-157 STEM_GC11AB-IVf-g-158 STEM_GC11AB-IVf-g-159 STEM_GC11AB-IVf-g-160 STEM_GC11AB-IVf-g-161 STEM_GC11AB-IVf-g-164 STEM_GC11AB-IVf-g-165 STEM_GC11AB-IVf-g-166 STEM_GC11AB-IVf-g-167 STEM_GC11AB-IVf-g-168 STEM_GC11AB-IVf-g-169
2. balance redox reactions using the change in oxidation number method
STEM_GC11AB-IVf-g-170
3. draw the structure of a galvanic cell and label the parts
STEM_GC11AB-IVf-g-171
4. identify the reaction occurring in the different parts of the cell
STEM_GC11AB-IVf-g-172 Page 14 of 17
K to 12 BASIC EDUCATION CURRICULUM SENIOR HIGH SCHOOL – SCIENCE, TECHNOLOGY, ENGINEERING AND MATHEMATICS (STEM) SPECIALIZED SUBJECT
CONTENT
CONTENT STANDARD
PERFORMANCE STANDARD
LEARNING COMPETENCIES
CODE
5. write the half-equations for the reactions occurring in the electrodes
STEM_GC11AB-IVf-g-173
6. write the balanced overall cell reaction 7. give different examples of galvanic cell 8. define reduction potential, oxidation potential, and cell potential 9. describe the standard hydrogen electrode 10. calculate the standard cell potential 11. relate the value of the cell potential to the feasibility of using the cell to generate an electric current 12. describe the electrochemistry involved in some common batteries: a. leclanche dry cell b. button batteries c. fuel cells d. lead storage battery
K to 12 Senior High School STEM Specialized Subject – General Chemistry 1 and 2 December 2013
STEM_GC11AB-IVf-g-174 STEM_GC11AB-IVf-g-175 STEM_GC11AB-IVf-g-176 STEM_GC11AB-IVf-g-177 STEM_GC11AB-IVf-g-178
STEM_GC11AB-IVf-g-179
STEM_GC11AB-IVf-g-180
13. apply electrochemical principles to explain corrosion
STEM_GC11AB-IVf-g-181
14. explain the electrode reactions during electrolysis
STEM_GC11AB-IVf-g-182
15. describe the reactions in some
STEM_GC11AB-IVf-g-183 Page 15 of 17
K to 12 BASIC EDUCATION CURRICULUM SENIOR HIGH SCHOOL – SCIENCE, TECHNOLOGY, ENGINEERING AND MATHEMATICS (STEM) SPECIALIZED SUBJECT
CONTENT
CONTENT STANDARD
PERFORMANCE STANDARD
LEARNING COMPETENCIES
CODE
commercial electrolytic processes 16. (LAB) Determine the potential and predict the cell reaction of some assembled electrochemical cells 17. (LAB) Describe the reactions at the electrodes during the electrolysis of water; cite the evidence for your conclusion
K to 12 Senior High School STEM Specialized Subject – General Chemistry 1 and 2 December 2013
STEM_GC11AB-IVf-g-184
STEM_GC11AB-IVf-g-185
Page 16 of 17
K to 12 BASIC EDUCATION CURRICULUM SENIOR HIGH SCHOOL – SCIENCE, TECHNOLOGY, ENGINEERING AND MATHEMATICS (STEM) SPECIALIZED SUBJECT
Code Book Legend Sample: STEM_GC11AB-IVf-g-183
DOMAIN/ COMPONENT
LEGEND
SAMPLE
Learning Area and Strand/ Subject or Specialization
Science, Technology, Engineering and Mathematics General Chemistry
Uppercase Letter/s
Domain/Content/ Component/ Topic
Grade 11
STEM_GC11AB
Acid-Base Equilibria and Salt Equilibria
Roman Numeral
*Zero if no specific quarter
Quarter
Fourth Quarter
IV
Lowercase Letter/s
*Put a hyphen (-) in between letters to indicate more than a specific week
Week
Weeks six to seven
f-g
Arabic Number
Competency
describe the reactions in some commercial electrolytic processes
Matter and Its Properties
MP
Measurements
MT
Atoms, Molecules and Ions
AM S
Stoichiometry
First Entry Grade Level
CODE
183
K to 12 Senior High School STEM Specialized Subject – General Chemistry 1 and 2 December 2013
Percent Composition and Chemical Formulas
PC
Mass Relationships in Chemical Reactions
MR
Chemical reactions and chemical equations
CR
Gases
G
Dalton’s Law of partial pressures
DL
Gas stoichiometry
GS KMT
Kinetic molecular theory of gases Electronic Structure of Atoms
ES
Electronic Structure and Periodicity
ESP
Chemical Bonding
CB
Organic compounds
OC
Intermolecular Forces and Liquids and Solids
MF
Physical Properties of Solutions
PP
Thermochemistry
TC
Chemical Kinetics
CK
Chemical Thermodynamics
CT
Chemical Equilibrium
CE
Acid-Base Equilibria and Salt Equilibria
AB Page 17 of 17
General Chemistry 1
120 MINS
Lesson 1: Matter and its properties
Content Standard
Lesson Outline
The learners demonstrate an understanding of the properties of matter and its various forms.
Introduction
Presentation of Learning Objectives and Important Keywords
Performance Standards
Motivation
Application of the Particulate State of Matter through Syringe Test
15
Instruction
Matter and its Properties
60
Enrichment
Demonstration on the Visualization of Matter
30
Evaluation
Written Task
10
The learners shall be able to: 1. Make a representation of the particulate nature of the three phases of matter; 2. Discuss the difference between: a. Pure substances and mixtures b. Elements and compound c. Homogeneous and heterogeneous mixtures; 3. Classify the properties of matter as: a. Physical or chemical b. Intensive or extensive; and 4. Perform simple separation procedures.
5
Materials Projector, Computer, Flip charts Resources (1) Chang, R. & Goldsby, K. (2016). Chemistry. (12th ed.). New York: McGraw-Hill.
Learning Competencies At the end of the lesson, the learners: 1. Recognize that substances are made of smaller particles (STEM_GGC11-MP-Ia-b-1); 2. Describe and make a representation of the arrangements, relative spacing, and relative motion of the particles in the three phases of matter (STEM_GGC11-MP-Ia-b-2); 3. Distinguish between physical and chemical properties and give examples (STEM_GGC11-MP-Ia-b-3); 4. Distinguish between extensive and intensive properties and give examples (STEM_GGC11-MP-Ia-b-4); 5. Use properties of matter to identify substances and to separate them (STEM_GGC11-MP-Ia-b-5);
6. Differentiate between pure substances and mixtures (STEM_GGC11-MP-Ia-b-6); 7. Differentiate between elements and compounds (STEM_GGC11-MP-Ia-b-7); 8. Differentiate between homogenous and heterogeneous mixtures (STEM_GGC11-MP-Ia-b-8); 9. Recognize the formula of common chemical substances (STEM_GGC11-MP-Ia-b-9); 10. Describe separation techniques for mixtures and compounds (STEM_GGC11-MP-Ia-b-10); and 11. Compare consumer products on the basis of their components for use, safety, quality, and cost (STEM_GGC11-MP-Ia-b-11).
1
INTRODUCTION (5 minutes)
Teacher Tip Display the objectives prominently on the board, so that the learners can track the progress of their learning.
2. Present the keywords for the concepts to be learned: a. Atoms b. Chemical properties c. Compounds d. Distillation e. Elements f. Extensive properties g. Filtration h. Gas i. Heterogeneous mixtures j. Homogeneous mixtures k. Intensive properties l. Ions m. Liquid n. Magnetic separation o. Mixtures p. Molecules q. Physical properties r. Pure substances s. Solid
Teacher Tip List these keywords on the board or through PowerPoint slides. Alternatively, you can write them on flip charts. The learners will be asked to complete a concept map using words from this list.
1. Introduce the learning objectives by using the suggested protocol (Read-aloud): a. I will be able to describe the particulate nature of the different forms of matter b. I will be able to classify the properties of matter c. I will be able to differentiate pure substance and mixtures; elements and compounds; homogeneous and heterogeneous mixtures d. I will be able to recognize the formulas of some common substances e. I will be able to discuss methods to separate the components of a mixtures f. I will be able to recognize chemical substances present in some consumer products
Another approach is to write these keywords in meta cards of different colors.
2
MOTIVATION (15 minutes)
1. Present two 60-mL plastic syringes with the needle removed and replaced by a seal. One syringe contains a small block of wood, while the other contains entrapped air. The plunger is set to touch the wood block, as shown below:
2. Ask them what will happen if the plunger will be pushed down the syringe. 3. Make one learner push the plunger in the two syringes, and check if they have predicted the behavior of the plunger in the two syringes correctly. 4. Ask them to answer the question: Why is it easier to compress the entrapped air than the wood block? 5. Highlight that a particulate model for matter is very useful in explaining the properties of matter. Point out that some basic concepts on matter that have been introduced in junior high school will be reviewed in this lesson.
INSTRUCTION (60 minutes)
1. Construct the following block diagram and make the learners fill it up using the keywords listed in the board.
Teacher Tip This demonstration is meant to make them realize the usefulness of visualizing matter being made up of particles. It is likely that their answers will be based on what they will recall from experience and from what they learned from junior high school: that the plunger can be moved more easily in the syringe containing an entrapped gas than in the other syringe containing a solid. A gas is made up of particles that are far • apart from each other, which can be pushed closer towards each other; • A solid is made up of particles which are compact, so that it is no longer possible to push these particles closer to each other. Let them recognize that the keywords to be encountered in the lesson are commonly used to describe the things around them.
Teacher Tip Make them take turns in filling up each box with the correct keyword (or in placing the proper meta cards). The block diagram can be presented through PowerPoint slides projected on a white board. Alternatively, it can be prepared on flip charts or on manila paper. Answer for Number 1 The keywords to be placed are: atoms; ions; molecules.
2. Ask them to answer the question: How do the following particles differ from each other? a. Atoms b. Molecules c. Ions
Answer for Number 2 a. Atoms – the smallest particle b. Molecules – composed of atoms c. Ions – particles with charges In case they fail to recall the differences, a short discussion might be necessary. Also, refer them to read Chapter 1 of the resource book (Chang, R. & Goldsby, K., Chemistry).
3
3. Construct the following block diagram and make them fill it up using the keywords listed in the board.
Answer for Number 3 The block diagram can be presented through PowerPoint slides projected on a white board. Alternatively, it can be prepared on flip charts or on manila paper. The keywords to be placed are: solid; liquid; gas
Answer for Number 4 The arrangement of the particles for solid, liquid, and gas, respectively are:
4. For the bottom layer of boxes, ask them to illustrate how the particles are distributed or arranged in each state of matter using circles. 5. Ask them to answer the following questions: a. How separated are the particles in each state of matter? b. How free are the particles to move in each state of matter? 6. Ask them to classify the following substances according to the three states of matter: a. Iron nail b. Sugar c. Syrup d. Air e. Ice f. Alcohol
4
Answer for Number 5 • Solid: closely packed; restricted motion • Liquid: far apart; free movement • Gas: very far apart; very free (chaotic) movement Answer for Number 6 a. Solid b. Solid c. Liquid d. Gas e. Solid f. Liquid
7. Construct the following block diagram and make them fill it up using the keywords listed in the board.
Answer for Number 7 The block diagram can be presented through PowerPoint slides projected on a white board. Alternatively, it can be prepared on flip charts or on manila paper. The keywords to be placed are: physical properties and chemical properties (left cluster); and extensive properties and intensive properties (right cluster).
8. Ask them to answer the following questions: a. What is the difference between physical properties and chemical properties? b. How do the extensive properties differ from the intensive properties?
9. Ask them to classify the following examples as physical or chemical properties: a. Melting of ice b. Evaporation of water c. Rusting d. Digestion 5
Answer for Number 8 a. In physical properties, no change in composition takes place during the determination or measurement of these properties. On the other hand, in chemical properties, a change in composition occurs during the determination or measurement of these properties. b. Extensive properties change their value when the amount of matter or substance is changed. Meanwhile, intensive properties do not change their value when the amount of matter is changed. In case they fail to recall the differences, a short discussion might be necessary. Also, refer them to read Chapter 1 of the resource book (Chang, R. & Goldsby, K., Chemistry). Answer for Number 9 a. Physical property b. Physical property c. Chemical property d. Chemical property
10. Ask them to classify the following examples as intensive or extensive properties: a. Boiling point b. Weight c. Volume
Answer for Number 10 a. Intensive property b. Extensive property c. Extensive property d. Intensive property
d. Density 11. Construct the following block diagram and make them fill it up using the keywords listed in the board.
Answer for Number 11 The block diagram can be presented through PowerPoint slides projected on a white board. Alternatively, it can be prepared on flip charts or on manila paper. The keywords to be placed are: pure substances and mixtures (top cluster); elements and compounds (bottom left cluster); homogeneous mixtures and heterogeneous mixture (bottom right cluster).
12. Ask them to answer the question: How do pure substances differ from mixtures?
Answer for Number 12 Pure substances are composed of only one component, while mixtures are composed of several components. In case they fail to recall the differences, a short discussion might be necessary. Also, refer them to read Chapter 1 of the resource book (Chang, R. & Goldsby, K., Chemistry).
6
13. Present the following substances (or pictures of these substances), and ask them to answer the question: Which of the following are pure substances and which are mixtures? a. Table sugar b. Table salt c. Iodized salt d. Brown sugar e. Distilled water f.
Answer for Number 13 a. Pure substance b. Pure substance c. Mixture d. Mixture e. Pure substance f. Mixture g. Pure substance h. Mixture
Soft drinks
g. Oxygen gas (in tank) h. Human breath 14. Ask them to answer the question: What is the difference between elements and compounds? Give examples of each.
15. Ask them to answer the question: What is the difference between homogeneous and heterogeneous mixtures?
7
Answer for Number 14 Elements are pure substances that are • made up of only one kind of atoms. Possible examples: iron; gold; mercury Compounds are pure substances made • up of two or more kinds of atoms. Possible examples: salt; sugar; water In case they fail to recall the differences, a • short discussion might be necessary. Also, refer them to read Chapter 1 of the resource book (Chang, R. & Goldsby, K., Chemistry). Answer for Number 15 A homogeneous mixture has a uniform • composition and exhibits the same properties in different parts of the mixture. A heterogeneous mixture has a non• uniform composition and its properties vary in different parts of the mixture. In case they fail to recall the differences, a • short discussion might be necessary. Also, refer them to read Chapter 1 of the resource book (Chang, R. & Goldsby, K., Chemistry).
16. Present the following mixtures (or pictures of these mixtures), and ask them the question: Which of the following are homogeneous mixtures? Which are heterogeneous mixtures? a. Rubbing alcohol b. Mixture of water and oil c. Mixture of salt and pepper
Answer for Number 16 a. Homogeneous mixture b. Heterogeneous mixture c. Heterogeneous mixture d. Homogeneous mixture e. Homogeneous mixture
d. Carbonated soft drink e. Human breath 17. Construct the following block diagram and make them fill it up using the keywords listed in the board. 18. Learner prompt: Look at this diagram. Give three common ways to separate the components of a mixture?
19. Ask them to answer the question: When can each method be used in separating the components of a mixture?
20. Ask them to answer the question: How can the following components of the following mixtures be separated? a. Salt from salt water b. Salt from a mixture of iron and salt 8
Answer for Number 17 The block diagram can be presented through PowerPoint slides projected on a white board. Alternatively, it can be prepared on flip charts or on manila paper. Answer for Number 18 Some keywords that can placed are filtration; distillation; magnetic separation; decantation; sublimation. Answer for Number 19 Filtration: to separate a solid from a liquid • in a heterogeneous mixture using a filtering membrane, like paper or cloth Distillation: to separate a liquid in a • homogeneous mixture Magnetic separation: to separate a • magnetic solid from a heterogeneous mixture Decantation: to separate a solid from a • liquid in a heterogeneous mixture based on gravity Sublimation: to separate a volatile solid • from a non-volatile solid Answer for Number 20 a. Heating to evaporate the water b. By adding water to dissolve the salt, and filter or decant to separate the iron.
ENRICHMENT (30 minutes)
1. Present a demonstration for the visualization of matter. This will reinforce the concept on the differences between pure substances, mixtures, elements, and compounds. See attached sheet.
Teacher Tip This activity can be done at the end of the lecture session.
2. Then conduct the learner’s activity on Visualization and Classification of Matter. See the teacher’s guide and learner’s worksheet.
EVALUATION (10 minutes)
1. Make them do an activity wherein they will apply the visualization of matter to classify pure substances, mixtures, elements, and compounds. See attached sheet. 2. Written task (assignment): Classify some substances found in the kitchen and in the bathroom as pure substances or mixtures; elements or compounds; and homogeneous or heterogeneous mixture.
Teacher Tip This activity can be done at the end of the lecture session. In case there is no longer enough time, it can be done during the laboratory session.
EVALUATION EXCEEDS EXPECTATIONS
MEETS EXPECTATIONS
The learner classified six or more substances in Part 1 correctly, and ten or more substances in their list in Part 2.
The learner classified four to five substances in Part 1 correctly, and six to nine substances in their list for Part 2.
NEEDS IMPROVEMENT
NOT VISIBLE
The learner classified less than The learner did not do the four substances in Part 1 correctly, assigned tasks. and less than five substances in their list for Part 2.
9
VISUALIZATION AND CLASSIFICATION OF MATTER
Introduction
In this activity, physical models, such as balls or beads, will be used to illustrate that matter is made up of particles. A ball or a bead will represent an atom of an element, and a combination of balls or bead will represent a compound. A collection of single balls and/or combined balls will be used to show the difference between pure substances and mixtures. This activity was adapted from Chemistry with Charisma, published by Terrific Science Press.
Materials
a. A set of balls or beads of two or more colors b. Zip lock bags
Procedure
1. Assemble the following sets of balls and place them in unlabelled zip lock bags.
4. Ask them the following questions: a. Are the balls the same or different? b. Do the balls represent a pure substance of a mixture? c. Do the balls represent an element or a compound? 5. Repeat Step 3 with the mixture set. 6. Repeat Step 3 with the compound set. 7. Repeat Step 3 with the diatomic set. Highlight and discuss the answer to the last question. 8. Introduce the concept of formulas. Each ball of a certain color will be assigned a letter (e.g. A for the white, B for the black, and C for another color). 9. Ask them for a possible formula for: a. the monoatomic element (Answer: A) b. the diatomic element (Answer: A2) c. the compound (Answer: AB)
2. Show the bags to the learners and tell them that their task is to find out if the bag contains a pure substance or a mixture. If the content is a pure substance, they have to determine if it is a monoatomic element, a diatomic element, or a compound. 3. Show them the bag with monoatomic elements, and ask them if it is a pure substance or a mixture. Take out the contents from the bag one by one, and show them to the learners. 10
LEARNER WORKSHEET BAG LABEL
PURE SUBSTANCE or MIXTURE
ELEMENT(S) or COMPOUND(S)
A
B
C
D
11
FORMULA FOR EACH SUBSTANCE IN THE BAG
LEARNER’S ACTIVITY: TEACHER’S GUIDE
Introduction
In this activity, physical models, such as balls or beads, will be used to illustrate that matter is made up of particles. A ball or a bead will represent an atom of an element, and a combination of balls or bead will represent a compound. A collection of single balls and/or combined balls will be used to show the difference between pure substances and mixtures. This activity was adapted from Chemistry with Charisma, published by Terrific Science Press.
Materials •
A set of balls or beads of two or more colors
•
Zip lock bags
Procedure
1. Assemble the following sets of balls and place them in zip lock bags labeled only with the letters.
2. Distribute the set of bags and ask the learners to fill up the provided worksheet (see below) using the bags labeled A to H. 3. Ask them to compare their results. 4. For the bags containing models of compounds, ask them to write the formula of the compound represented by the model.
12
LEARNER’S ACTIVITY: LEARNER WORKSHEET: VISUALIZATION AND CLASSIFICATION OF MATTER
Introduction
In this activity, physical models, such as balls or beads, will be used to illustrate that matter is made up of particles. A ball or a bead will represent an atom of an element, and a combination of balls or bead will represent a compound. A collection of single balls and/or combined balls will be used to show the difference between pure substances and mixtures. This activity was adapted from Chemistry with Charisma, published by Terrific Science Press.
Materials
1. A set of balls or beads of two or more colors 2. Zip lock bags
Procedure
1. Obtain a set of bags with physical models of the particles of different substances from your teacher. 2. Examine the particles in each bag and classify them as pure substances or mixtures, monoatomic elements, or diatomic elements. Fill up the worksheet provided below using the bags labeled A to H. 3. For the bags with models of compounds, write the formula of the compound represented by the model. BAG LABEL
PURE SUBSTANCE or MIXTURE
ELEMENT(S) or COMPOUND(S)
A B C D E F G H
13
FORMULA FOR EACH SUBSTANCE IN THE BAG
General Chemistry 1
120 MINS
Lesson 2: Matter and its Various Forms Content Standard The learners demonstrate an understanding of the properties of matter and its various forms.
Lesson Outline Introduction
Pre-Laboratory Work
Motivation
Inquiry
The learners shall be able to
Instruction
Experiment
90
1. Perform simple separation procedures.
Enrichment
Discussion of Alternative Procedures for the Separation
15
Evaluation
Submission of the Report on the Experiment
Performance Standard
Learning Competency At the end of the lesson, the learners: 1. Apply simple separation techniques such as distillation, chromatography (STEM_GGC11-MP-Ia-b-12).
10 5
Materials Laboratory glassware or alternative containers Resources (1) Separation of a mixture [PDF file]. Retrieved from Princeton High School web site: http://phs.princetonk12.org/teachers/jgiammanco/ Chem%201/Labs/C2-SepMixtureLab.pdf (2) Solar still challenge [PDF file]. Retrieved from American Chemical Society web site: http://www.acs.org/content/dam/acsorg/global/ iyc2011/global-water-experiment-purification.pdf
14
INTRODUCTION (10 minutes)
1. This introduction can serve as a pre-laboratory discussion prior to the experiment proper. 2. Ask the learners to recall how to differentiate a pure substance from a mixture. 3. Point out that mixtures are common and that in some situations, it is necessary to separate the components or to isolate one component of a mixture.
Teacher Tip A laboratory experiment sheet has to be prepared and distributed to the learners. The experiment described in Annex 1 could be adopted or revised to suit the available facilities.
4. State the objective of the experiment they will be performing.
MOTIVATION (5 minutes)
1. Ask them how table salt is obtained from seawater. 2. As an alternative, you can ask how drinking water is obtained from seawater.
INSTRUCTION (90 minutes)
1. Provide each group with a prepared mixture of salt, sand, and iron filings. 2. Ask them to follow the procedure in the experiment sheet.
Teacher Tip This could be given as an assignment before the laboratory session. They will be asked to search the internet on how these processes are actually carried out.
Teacher Tip Low-cost (or zero-cost materials) can be used in place of the materials described in the experiment sheet: a. A vial can be used instead of the evaporating dish, and the watch glass can be omitted. A moistened filter paper can be used to cover the vial. b. A vial or a small bottle can be used in place of a beaker. c. A plastic funnel used at home can be a substitute for the glass funnel. They can be asked to make a flow diagram of the procedure. If desired, the experiment can be performed as a quantitative procedure wherein the isolated substances will be dried and weighed.
15
ENRICHMENT (15 minutes)
1. In the post-laboratory discussion, ask them what properties of the components were used to separate each from the other. 2. Discuss possible alternative procedures for the separation. 3. They can be asked to perform the Solar Still Challenge, as described in the following internet webpage: http://www.acs.org/content/dam/acsorg/global/iyc2011/global-water-experimentpurification.pdf
EVALUATION
1. Ask them to submit a report on the experiment. 2. They could be provided with a worksheet that they have to fill up, which could include some questions.
Teacher Tip The volatile nature of naphthalene enabled its sublimation. Point out that the odor of naphthalene is caused by the vapor it produces. The difference in the solubility of sodium chloride and sand (or silicon dioxide) in water was used in separating the two components. An alternative procedure could involve the differences in the solubility of the components in alcohol and in water. a. Naphthalene dissolves in ethanol but not in water. b. Sodium chloride dissolves in water, but not in alcohol. c. Silica does not dissolve in alcohol and in water. This experiment was conducted as part of the Global Experiment for the International Year of Chemistry in 2011. It could be done to motivate their innovative skills.
EVALUATION EXCEEDS EXPECTATIONS The learner: i. performed the experiment correctly; ii. described the results correctly; iii. discussed the results of the experiment very well; and iv. performed the Solar Still Challenge.
MEETS EXPECTATIONS
NEEDS IMPROVEMENT
The learner: The learner: i. performed the experiment i. performed the experiment correctly; correctly; ii. described the results correctly; ii. described the results correctly; iii. discussed the results of the but iii. did not discuss the results of experiment well, but iv. did not perform the Solar Still the experiment, and iv. did not perform the Solar Still Challenge. Challenge. 16
NOT VISIBLE The learner: i. did not do the assigned task.
SEPARATION OF THE COMPONENTS IN A MIXTURE Introduction
Several components, which retain their identity and characteristic properties, are present in a mixture. No chemical reactions occur between the components of a mixture. Many of the materials surrounding us are mixtures, such as soil, cement, soft drinks, and pharmaceuticals. In this experiment, the components of a mixture will be separated from each other. The techniques applied for this separation does not involve a chemical reaction, so that the isolated components will retain their identity.
Materials
1. A mixture containing the following: a. Sodium chloride, NaCl b. Naphthalene c. Silicon dioxide, SiO2 (sand) 2. Digital balance 3. Beaker
4. 5. 6. 7. 8. 9.
Funnel Watch glass Masking tape Evaporating dish Filter paper Hot plate
Procedure
1. Weigh 0.50 to 0.60 g of the mixture on the digital balance. 2. Place the mixture on an evaporating dish and cover it with the pre-weighed watch glass. 3. Seal the sides with masking tape. 4. Place a moist tissue paper over the watch glass, and gently heat the evaporating dish until white vapors are emitted. 5. Cool the setup and carefully remove the watch glass. Describe the solid adhering to the watch glass. 6. Pour distilled water into the mixture remaining in the evaporating dish and stir it carefully. 7. Filter the mixture and collect the filtrate in the pre-weighed beaker. Wash the residual solid in the filter paper with a small amount of water, combining the washing with the filtrate. 8. Gently heat to evaporate the water in the filtrate. 9. Cool the beaker. Describe the solid remaining in it 10. Dry the filter paper with the sand in an oven at 100oC. Describe the solid remaining in the filter paper. 17
Treatment of Results 1. Record the description of the substances isolated in the experiment. Tabulate your data below: DESCRIPTION Solid adhering to the watch glass Solid remaining in the beaker Solid remaining in the filter paper
2. Knowing the substances present in the mixture, identify the isolated solids. IDENTITY Solid adhering to the watch glass Solid remaining in the beaker Solid remaining in the filter paper 3. Devise another procedure to separate the components of the mixture used in the experiment.
18
General Chemistry 1
60 MINS
Lesson 3: Measurements Content Standard The learners demonstrate an understanding of measurement and the difference between accuracy and precision.
Lesson Outline Introduction
Communicating Learning Objectives
3
Motivation
Why is Measurement Important?
7
1. Discuss the need and describe the result of a measurement, in general;
Instruction
Demonstration
30
2. Differentiate between the accuracy and precision of a measurement;
Enrichment
Laboratory Experiment
15
Evaluation
Take-home Activity
Performance Standards The learners shall be able to:
3. Point out possible sources of errors in a measurement; and 4. Carry out a measurement and report the results correctly.
Materials Projector, Computer, Flip charts
Learning Competencies
Resources (1) Chang, R. & Goldsby, K. (2016). Chemistry. (12th ed.). New York: McGraw-Hill.
At the end of the lesson, the learners: 1. Explain the need for measurements; 2. Describe how to carry out measurements of length, mass, and volume; and 3. Dfferentiate between precision and accuracy (STEM_GC11MT-Ib-13).
19
5
INTRODUCTION (3 minutes)
1. Introduce the following learning objectives using the suggested protocol (Read-aloud): a. I will be able to describe the need for measurement
Teacher Tip The lesson is essentially a review of some concepts presented and used in junior high school.
b. I will be able to carry out simple measurements of length, volume, and mass c. I will be able to differentiate the accuracy and the precision of a measurement Teacher Tip List these keywords on the board or through PowerPoint slides. Alternatively, you can write them on flip charts.
2. Present the keywords for the concepts to be learned: a. Measurements b. Units of measurement c. Accuracy d. Precision e. Significant figures f.
Errors
MOTIVATION (7 minutes)
1. Present to two plastic bottles containing different amounts of water, and ask the learners to describe and differentiate the two objects. Make them realize the need to use a number (the volume of the water content or the weight of the bottles and their contents) to describe the objects more clearly and to differentiate them. 2. Make them realize the need for a quantitative or a numerical description of some properties of matter, and how this is applied in their daily lives. Ask them to cite some situations in daily life where a measurement is important.
INSTRUCTION (30 minutes)
1. After the motivation, they will see the importance of a quantitative description of some parameters, such as length, mass, and volume. 2. Call two learners separately. Ask each one to measure the length of a table without using a ruler, meter stick, or tape measure. Make them write their measurements on the board (number, unit: e.g., 3 hand spans). 3. Ask the class to compare the results and explain for differences or similarities. Ask them to answer the question: Why is there a need to use a common unit for measurement? 20
Teacher Tip Alternatively, a small and a big ball of the same color and material can be used. Another option is to use a small and a long plastic ruler. Sample Responses a. Measuring the ingredients during cooking (or baking) b. Measuring the weight of salt being purchased
Teacher Tip It is expected that the learners will use the span of their fingers, hands, or arms for the measurement.
4. Introduce the concept of unit of measurement, which is a means for a quantitative description of a property. Highlight the need for a common or universally accepted unit of measurement. 5. Point out that for scientific measurements, a common system has been agreed upon and is used by all scientists anywhere and all the time. Ask them to answer the question: What is the measurement system adopted in scientific measurements? 6. Post this table on the board and ask them to supply the unit for each property. PROPERTY
Teacher Tip The results of the measurements will be different because of the difference in the length of their finger, hand, or arms. Make them recall from their Science class in junior high school that the International System (or SI) of Measurement is being used in measurements in science.
SI Units Teacher Tip The table can be presented through PowerPoint slides projected on a white board. Alternatively, it can be prepared on flip charts or on manila paper.
Length Mass
It is expected that they will be able to fill up the table, recalling what they have learned from junior high school.
Volume Time Temperature
PROPERTY
SI Units
Length 7. Ask them to cite some examples where these units of measurements are used in real life.
Mass
Sample responses:
Volume
•
Length – in measuring the height of a person; distances; the size of cloths
•
Mass – in measuring the weight of a person; the amount of salt or sugar being bought
Time
•
Volume – in measuring the amount of a liquid (e.g. soft drinks)
•
Time – in measuring the duration of an event (e.g. to run through a distance)
•
Temperature – in measuring the body temperature of a person or of the atmosphere.
8. Ask them to group into pairs. Tell them to measure the length, width, and thickness of a book, and record their results on the following table (to be shown on the board).
Temperature In case they fail to recall the correct units of measurement, a short discussion might be necessary. Also, refer them to read Chapter 1 of the resource book (Chang, R. & Goldsby, K., Chemistry).
All pairs should measure the same book.
21
MEASUREMENTS
TRIAL 1
TRIAL 2
TRIAL 3
Length Mass Volume
Teacher Tip The correct results will include two decimal units.
Time Temperature 9. When the pairs have completed the measurements and recorded their results on the table, ask them to answer the question: How many significant figures did you use in reporting your measurements? 10. Explain that based on the calibration of the ruler, the measurement is certain until the first decimal unit and that the result can include one insignificant or uncertain figure.
!
The concept of significant figures has been presented in junior high school, but it might not have been fully understood. Therefore, reviewing it would be worthwhile. For the guidelines for using significant figures, see Chapter 1 of the resource book (Chang, R. & Goldsby, K., Chemistry). Let them examine the ruler they used. At the end of this short activity, you should address misconceptions that they have on the concepts presented. The concepts of accuracy and precision have been presented in junior high school. It would be worth reviewing these concepts.
11. Ask them to examine the results of the three measurements that they made on the length, width, and thickness of the book. Ask them to answer the following questions: a. Are the results of each measurement (length, width or thickness) close to each other? b. Were the measurements accurate or precise? 12. Write the actual length, width, and thickness of the book on the board, and ask them to compare their results with this value. 22
Point out that the closeness of the results of a measurement to each other is expressed by its precision. It is not suggested that they should be made to quantify precision in terms of standard deviation. This will be done in their course on Mathematics (or Statistics).
Ask them to answer the following questions: a. Are the results of each measurement (length, width, or thickness) close to the true value? b. Were the measurements accurate or precise? 13. Let them recall the difference between accuracy and precision. Then, state the definitions of accuracy and precision as used in measurement.
Answer Key It is likely that: a. The results will be close to the true value; b. The measurements were accurate Point out that the closeness of the results of a measurement to the true value is expressed by its accuracy.
14. Evaluating the accuracy of a measurement will require the true value. However, the true value for the dimensions of the book is not available. Point out that if twenty or more measurements were done, the mean value can be taken as the true value. This is an assumption in statistics. 15. Draw the following dot plots on the board, and explain that each dot is the result of a measurement whose value is indicated in the horizontal (or x-) axis. Tell them that the plot presents the results of six measurements of the weight of a pebble whose true weight is 8.0 g. Ask them to determine whether each measurement is accurate or inaccurate, and precise or imprecise.
Teacher Tip The dot plot can be drawn on a manila paper before class, or presented through a PowerPoint slide.
Answer Key (A) Accurate and precise (B) Accurate and imprecise (C) Inaccurate and imprecise (D) Inaccurate and precise
23
16. Highlight that the measurement they made could have errors, which could: I.
Cause the result to be far from the true value (low accuracy). These errors are known as systematic errors.
II. Cause the results to be different from each other (low precision). These errors are known as random errors. Ask them to answer the question: What possible errors did the person who made the measurements commit to lower the accuracy of the results? To lower the precision of the results?
ENRICHMENT (15 minutes)
Make the learners perform a laboratory experiment on the determination of density. This activity will reinforce the concept of measurements, the units used, and the concept of significant figures.
EVALUATION (5 minutes)
1. Assign them to read the labels of some canned or bottled goods in the kitchen, and report the mass or volume of the contents. 2. Let them classify the following measurement data as high precision or low precision: a. Volume of a liquid: 11.0 cm3, 11.3 cm3, 10.9 cm3, 11.1 cm3 b. Mass of a solid: 25.0 g, 23.0 g, 20.0 g, 28.0 g
24
Teacher Tip Point out that the errors could be due to the measuring instrument or due to the person doing the measurement
Teacher Tip Refer to the Teacher’s Guide for this laboratory activity.
General Chemistry 1
120 MINS
Lesson 4: Measurements (Laboratory) Content Standard
Lesson Outline
The learners demonstrate understanding of basic measurement skills. Introduction
State the Objectives of the Experiment
5
The learners shall be able to:
Motivation
Application of Density Data
5
1. Carry out a measurement and report correctly the results.
Instruction
Experiment
90
Enrichment
Discussion of the Interpretation of the Graph
20
Evaluation
Report
Performance Standard
Learning Competency At the end of the lesson, the learners: 1. determine the density of a liquid (STEM_GC11MT-Ib-14).
Materials Simple laboratory glassware or low-cost alternatives Resources (1) Laboratory experiment in Annex 1
25
INTRODUCTION (5 minutes)
1. State the objective of the experiment that the learners will be performing. 2. Ask them to recall the definition of density and the formula for calculating it. 3. Review the methods for measuring weight and volume.
Teacher Tip A laboratory experiment sheet has to be prepared and distributed to the learners. The experiment found in the Annex makes use of low-cost materials. Density is used as a means to obtain the concentration of a solution.
MOTIVATION (5 minutes)
1. Point out some application of density data in industry.
INSTRUCTION (90 minutes)
Each group should be provided with different concentrations so that the relationship between density and concentration can be shown. Sugar solution can be used instead of salt solution.
ENRICHMENT (20 minutes)
This relationship can be used as a means to determine the concentration of a solution.
1. Provide each group with a salt solution of a given concentration. 2. Ask them to follow the procedure in the experiment sheet.
1. Discuss the interpretation of the graph between density and the concentration of the solution. 2. Assign them internet research on the density of the following: a. Regular soda in can b. Light soda in can c. Soda with aspartame in can 3. Ask them to explain the difference in density of these soft drinks.
EVALUATION
Point out that this relationship is used in industry to monitor the concentration of some solutions. The different drinks contain different concentrations of sugar, so their density will vary. They could be provided with a worksheet that they have to fill up. It could include some questions.
1. Ask them to submit a report on the experiment. EXCEEDS EXPECTATIONS
MEETS EXPECTATIONS
NEEDS IMPROVEMENT
The learner:
The learner:
The learner:
i.
i.
i.
performed the experiment correctly;
performed the experiment correctly;
performed the experiment correctly;
ii. described the results correctly; and
ii. described the results correctly; and
ii. described the results correctly; but
iii. discussed the results of the experiment very well.
iii. discussed the results of the experiment well.
iii. did not discuss the results of the experiment. 26
NOT VISIBLE The learner: i. did not do the assigned task.
DENSITY OF AN AQUEOUS SOLUTION
Introduction
Density is an important property of matter. It expresses the weight of a unit volume of a substance, is used to characterize substances, and can provide a means for the identification of a solid, a liquid, or a gas.
In this experiment, the density of an aqueous solution will be determined by measuring the weight of different volumes of these solutions. Several solutions containing different concentration of a solute will be assigned to different groups, and the variation of the density of the solutions with the solute concentration will be studied. The behavior that you will observe has important applications in industrial and in health monitoring.
Materials
1. NaCl solution, in 5%, 10%, 15%, and 20% concentrations 2. Digital balance 3. Syringe, 1 mL 4. Plastic mini tray
Procedure
1. Place the plastic mini tray on the stage of the digital balance and measure its weight. 2. Measure 1 mL of the test solution into the syringe, making sure that no air bubbles are trapped. 3. Slowly transfer the liquid in the syringe onto the mini tray. Measure the weight of the tray with the solution in it. 4. Repeat Steps 1 to 3 to provide a duplicate measurement. This will be used to check the repeatability of the results. 5. Repeat the whole procedure using 2 mL and 3 mL of the solution.
27
Treatment of results
1. Record the weight of the mini tray at the beginning of the experiment. Record the weight after each addition of 1 mL, 2 mL, and 3 mL of the sample solution. MEASUREMENTS
TRIAL 1
TRIAL 2
Weight of empty container Weight of empty container + 1 mL solution Weight of empty container Weight of empty container + 2 mL solution Weight of empty container Weight of empty container + 3 mL solution
2. From the data above, calculate the weight of each of the different volumes that you have added to the plastic mini tray by subtracting the weight before the addition from the weight after the addition. Calculate the average value of the measured weights.
MEASUREMENTS
TRIAL 1
Weight of 1 mL solution Weight of 2 mL solution Weight of 3 mL solution
28
TRIAL 2
3. From the data in the previous table, calculate the density of the solution. Calculate the average value of the density. MEASUREMENTS
DENSITY OF SOLUTION
Based on 1 mL solution Based on 2 mL solution Based on 3 mL solution AVERAGE
4. Obtain the results from the other groups who used different concentrations of the solution. Tabulate the density of the various solutions studied.
CONCENTRATION
5%
10%
15%
20%%
Density, g/mL
5. Plot the concentration of the solution (in the x-axis) against its density (in the y-axis). Infer how the density varies based on the concentration of the solution.
29
General Chemistry 1
160 MINS
Lesson 5: Atoms, Molecules, and Ions (Lecture) Lesson Outline
Content Standard The learners demonstrate understanding of the structure of an atom and the formula and the name of compounds.
Introduction
Presentation of Learning Objectives and Important Keywords
5
Performance Standards
Motivation
The Particles that Make Up an Atom
5
Instruction
The Laws of Chemical Changes
Enrichment
Laboratory Session
10
Evaluation
Check Up Quiz
20
The learners shall be able to: 1. Describe the structure of an atom of an element; 2. Recognize and differentiate atoms, molecules, and ions; and 3. Write the formula and give the name of simple compounds. Learning Competencies
120
Materials Projector, Computer, Flip charts
At the end of the lesson, the learners:
1. Explain how the basic laws of matter (Law of Conservation of Mass, Law of Constant Composition, and Law of Multiple Proportion) led to the formulation of Dalton’s Atomic Theory (STEM_GC11AM-Ic-e-15); 2. Describe Dalton’s Atomic Theory (STEM_GC11AM-Ic-e-16); 3. Differentiate among atomic number, mass number, and isotopes, and which of these distinguishes one element from another (STEM_GC11AM-Ic-e-17); 4. Write isotopic symbols (STEM_GC11AM-Ic-e-18); 5. Recognize common isotopes and their uses (STEM_GC11AMIc-e-19); 6. Differentiate among atoms, molecules, ions, and give examples (STEM_GC11AM-Ic-e-20);
Resources (1) Chang, R. & Goldsby, K. (2016). Chemistry (12th ed.). New York: McGraw-Hill.
7. Represent compounds using chemical formulas, structural formulas, and models (STEM_GC11AM-Ic-e-21); 8. Give the similarities and differences between the empirical formula and molecular formula of a compound (STEM_GC11AM-Ic-e-22); and 9. Name compounds given their formula and write formulas given the name of the compound (STEM_GC11AM-Ic-e-23).
30
INTRODUCTION (5 minutes)
1. Introduce the following learning objectives using the suggested protocol (Read-aloud): a. I will be able to describe and discuss the basic laws of chemical change
Teacher Tip Display the objectives prominently on the board, so that the learners can track the progress of their learning.
b. I will be able to discuss how Dalton’s Atomic Theory could explain the basic laws of chemical changes c. I will be able to give the information provided by the atomic number and mass number of an atom and its isotopes d. I will be able to differentiate atoms, molecules, and ions e. I will be able to write the chemical formula of some molecules f.
I will be able to differentiate a molecular formula and an empirical formula
g. I will be able to give the name of a compound, given its chemical formula Teacher Tip List these keywords on the board. They will be asked to complete a concept map based on words on this list.
2. Present the keywords for the concepts to be learned: a. Law of Conservation of Matter b. Law of Definite Proportion c. Law of Multiple Proportion d. Dalton’s Atomic Theory e. Atomic number f.
Mass number
g. Isotope h. Atom i.
Molecule
j.
Ion
k. Chemical formula l.
Molecular formula
m. Empirical formula
31
MOTIVATION (5 minutes)
1. Call one of the learners to the front and give him/her a piece of paper. Ask him/her to cut the paper in half, and then cut one of the halves again in half, and again and again. Let him/her proceed as long as s/he can cut a piece into half. 2. Ask him/her the question: Can you go on cutting the paper into half? 3. Tell him/her that though the cutting can go on and on mentally, there is a physical limit to this process. It is impossible to cut the paper into half forever. There is a limit – a point where the piece can no longer be divided. 4. Highlight that the limit is an indivisible piece, which was called by the Greek philosopher Democritus as the atom. 5. However, beginning in the late 1800s, experiments have indicated that atoms are made up of smaller particles. 6. Ask them the question: What are these particles that make up the atom? 7. Point out that the science of chemistry is based on the concept of the atom and molecules. Knowledge of the atoms and molecules in the environment and in biological systems has provided an understanding of the changes occurring in them. It has also allowed the prediction of their behavior and the solution to any problem observed in their behavior.
INSTRUCTION (120 minutes)
1. Present the laws of chemical changes. These laws were inferred from several experiments conducted during the 18th century using a balance for the measurements: a. Law of Conservation of Mass b. Law of Definite Proportion c. Law of Multiple Proportion 2. Introduce the Law of Conservation of Mass: In a chemical reaction, no change in mass takes place. The total mass of the products is equal to the total mass of the reactant. 3. Antoine Lavoisier, a brilliant French chemist, formulated this law by describing one of his experiments involving mercuric oxide. He placed a small amount of mercuric oxide, a red solid, inside a retort and sealed the vessel tightly.
32
Teacher Tip The law might have been presented in the Science course in junior high school. In this case, ask a learner to state the law.A PowerPoint slide can be prepared for this part.
He weighed the system, and then subjected it to high temperature. During the heating, the red solid turned into a silvery liquid. This observation indicated that a chemical reaction took place. After which, the setup was cooled and then weighed. The weight of the system was found to be the same as before heating. Illustrate an application of this law through the following problems. Ask them to solve the problems in their seats, and ask one learner to write his/her solution on the board: a. How many grams of water will be formed if 1.00 g hydrogen gas reacts with 8.00 g oxygen? The reaction can be represented by the following word equation: hydrogen + oxygen ! water b. 5.58 g iron reacted with 3.21 g sulfur. How many grams of iron (II) sulfide were produced? The reaction involved was: iron + sulfur ! iron(II) sulfide c. Magnesium burns in air to form magnesium oxide, as represented by the following word equation: magnesium + oxygen ! magnesium oxide When 2.43 g magnesium was burned, 4.03 g magnesium oxide was produced. How many grams of oxygen reacted with the magnesium? d. Ammonia is produced by the reaction of nitrogen with hydrogen: nitrogen + hydrogen ! ammonia 33
Teacher Tip The law might have been presented in the Science course in junior high school. In this case, ask a learner to state the law.A PowerPoint slide can be prepared for this part.
How many grams of nitrogen combined with 50.0 g hydrogen is needed to yield 283.3 g ammonia? 4. State the Law of Definite Proportion: A compound always contains the same constituent elements in a fixed or definite proportion by mass. If water samples coming from different sources are analyzed, all the samples will contain the same ratio by mass of hydrogen to oxygen.
Teacher Tip The law might have been presented in the Science course in junior high school. In this case, ask a learner to state the law.
This experiment can be best described using a PowerPoint slide. A picture of the burning magnesium can be included in the slide.
5. Illustrate the application of this law using the previous example of magnesium reacting with oxygen: a. Describe an experiment wherein different amounts of magnesium powder are heated in air. b. Magnesium burns brightly in air and reacts with oxygen. During the reaction, the gray powder turns into a white substance. The reaction causes the weight of the solid to increase. c. The following data were collected:
WEIGHTS OF MAGNESIUM
WEIGHT OF PRODUCT
WEIGHT OF OXYGEN COMBINED WITH MAGNESIUM
Length Mass Volume Time Temperature
34
RATIO OF MASS OF OXYGEN TO MASS OF MAGNESIUM
Magnesium
Product
Oxygen
Ratio
3.00
7.56
4.56
1.52
5.00
12.60
7.60
1.52
7.00
17.64
10.64
1.52
d. Ask them to complete the third column by applying the Law of Conservation of Mass. e. Ask them to fill up the fourth column by dividing the mass of oxygen (third column) by the mass of the magnesium (first column). 6. Ask them to solve the following problems: a. In the first problem given earlier, it was given that 1.00 g hydrogen combines with 8.00 g oxygen. How many grams of hydrogen will react with 10.00 g oxygen? b. In the previous set of problem, it was seen that 5.58 g iron reacted with 3.21 g sulfur. Based on this information, calculate how many grams of iron will combine with 80.0 g sulfur.
Teacher Tip Ask them to solve the problem in their seats. Call one learner to write his/her solution on the board. Answer Key 1. 1.25 g Solution: 2.
139 g Solution:
7. Present the Law of Multiple Proportions: If two elements can combine to form more than one compound, the masses of one element that will combine with a fixed mass of the other element are in a ratio of small whole numbers. 8. Illustrate the application of this law using the example of carbon which reacts with oxygen to form carbon monoxide and carbon dioxide. a. In carbon monoxide, 1.00 g carbon combines with 1.33 g oxygen; whereas, in carbon dioxide, 1.00 g carbon combines with 2.66 g oxygen. b. It can be seen that the ratio is 1:2. 9. Remind them that laws are derived from experimental results. A theory is formulated to provide an explanation to the laws.
35
The law might have been presented in the Science course in junior high school. In this case, ask a learner to state the law. Pictures or meta cards with chemical formulas may be posted on the board and used to facilitate discussion. It is highly encouraged to use pictures of actual substances.
Dalton’s Atomic Theory, proposed by John Dalton, can be used to explain the laws of chemical change. This theory is based on the following set of postulates: 1. Elements are made up of very small particles known as atoms. 2. All the atoms of an element are identical in mass and size, and are different from the atoms of another element. Dalton used the different shapes or figures to represent different elements, as follows:
Teacher Tip Draw atoms to clarify each postulate, particularly Postulates 2, 3, and 4. Drawing the Dalton symbols for the element will facilitate the understanding of Postulates 2 and 3.
• Oxygen
Hydrogen
Carbon
Nitrogen
Phosphorus
Sulfur
3. Compounds are composed of atoms of more than one element, combined in definite ratios with whole number values.
Carbon monoxide
Carbon dioxide
Nitric oxide
4. During a chemical reaction, atoms combine, separate, or rearrange. No atoms are created and no atoms disappear.
+""""2 Carbon
Oxygen
Carbon dioxide
5. Ask them which postulate could provide an explanation for the:
Answer Key a. Postulate 4 b. Postulate 3
a. Law of Conservation of Mass b. Law of Definite Proportion 6. Remind them that during the time of Dalton, the atom was believed to be the smallest particle comprising substances. However, before the end of the 19th century, experiments provided proof of the existence of smaller particles within the atom.
36
7. Ask them to recall the particles contained in an atom (or the subatomic particles) and differentiate the particles in terms of location, charge, and relative mass by filling up the following table: PARTICLE
LOCATION
CHARGE
RELATIVE MASS
Ask them to recall the information about the composition of an atom provided by the following: a. Atomic number b. Mass number
Teacher Tip This has been presented in the Science course in junior high school.
PARTICLE
LOCATION
CHARGE
RELATIVE MASS
PROTON
Nucleus
+1
1
ELECTRON
Outside nucleus
-1
0.0006
NEUTRON
Nucleus
0
1
As enrichment, assign them to read and make a report on the discovery of the existence of the electron, proton, and nucleus. The concepts of atomic number and mass number have been presented in the Science course in junior high school.
Confirm that the above numbers are defined by the following equations: a. Atomic number = number of protons = number of electrons in a neutral atom b. Mass number = number of protons + number of neutrons
The table can be presented through PowerPoint slides projected on a white board. Alternatively, it can be prepared on a flip chart or on manila paper.
8. To apply these concepts, ask them to fill up the following table: ATOMIC NUMBER
MASS NUMBER
4
9
14
28
NUMBER OF PROTONS
NUMBER OF ELECTRONS
8
NUMBER OF NEUTRON
9
11
12 52
24 19
20 37
Atomic Number
Mass number
Number of Protons
Number of electrons
Number of neutrons
4
9
4
4
5
14
28
14
14
14
8
7
8
8
9
11
23
11
11
12
24
52
24
24
28
19
39
19
19
20
9. Introduce the concept of isotopes – atoms of an element having the same atomic number but different mass number. The existence of isotopes was shown by mass spectroscopy experiments, wherein elements were found to be composed of several types of atoms, each with different masses. a. The atomic number identifies an element. The atoms of isotopes of an element have the same number of protons and electrons. b. The atoms of isotopes of an element differ in the number of neutrons. 10. To apply the concept of isotopes, ask them to complete the following table containing information about the isotopes of hydrogen: PROTIUM (Hydrogen)
DEUTERIUM
TRITIUM
Atomic Number
1
1
1
Mass number
1
2
3
ISOTOPE
Number of protons Number of electrons Number of neutrons The common hydrogen atom is protium, while deuterium is found in heavy water. Ask them to recall the difference between the following particles: a. Atom b. Molecule c. Ion
38
Teacher Tip For better understanding of the concept of isotopes, they can be assigned to read about mass spectroscopy. Make them refer to General Chemistry books instead of the internet, because the latter might lead them to complicated description of this technique. Ask them to answer the following questions afterwards: 1. What does a mass spectrometer do? 2. How does the mass spectro-meter separate isotopes of different masses? The table can be presented in PowerPoint slides projected on a white board. Alternatively, it can be prepared in flip charts or on manila paper. ISOTOPE
PROTIUM
DEUTERIUM
TRITIUM
Atomic Number
1
1
1
Mass number
1
2
3
Number of protons
1
1
0
Number of electrons
1
1
1
Number of neutrons
1
1
2
Let them complete the following concept map showing the relationship of these particles: Atoms Gain of electrons
?
Loss of electrons
Answer Key:
?
11. Emphasize that each element has a characteristic atom. a. Dalton differentiated the elements and their atoms through drawings. b. However, in present day, elements are differentiated and represented through symbols. i.
Assign them to find information from the internet on useful isotopes. These concepts might have been presented in the Science course in junior high school.
The concepts of characteristic atoms and ions might have been presented in the Science course in junior high school.
Many symbols are abbreviations derived from the name of the element.
ii. Some symbols are derived from their Latin names.
Atoms Gain of electrons
Call five or more learners to write some elements and their names and symbol on the board. Make them recall that the difference between an ion and an atom is the presence of charges. The simple ions are derived from atoms through the gain or loss of an electron. Let them complete the following concept map showing the relationship of these particles:
Ions can be made up of only one atom (monoatomic) or more than one type of atom (polyatomic). 39
CATIONS (Positive Ions)
An alternative diagram could be:
Loss of electrons
ANIONS (Negative Ions)
12. Monoatomic ions are named based on the element. a. For cations, the name of the element is unchanged. If an element can form two ions of different charges, the name, which is usually derived from its Latin name, is modified by the suffix –ic for the ion with the higher charge, and –ous for that with the lower charge.
Teacher Tip The naming of the compound or molecule will be discussed later.
b. For anions, the name of the element is modified by the suffix –ide. Answers for Number 13 a. Zn2+ – zinc ion b. Mg2+ – magnesium ion c. K+ – potassium ion d. Fe2+ – ferrous ion or iron (II) ion e. Fe3+ – ferric ion or iron(III) ion
13. Ask them to name the following cations: a. Zn2+ b. Mg2+ c. K+ d. Fe2+ e.
Answers for Number 14 a. Br- – bromide ion b. S2- – sulfide ion c. O2- – oxide ion d. I- – iodide
Fe3+
14. Ask them to name the following anions: a. Brb.
Teacher Tip Provide them with a list of the common anions, together with their names.
S2-
c. O2d. ISeveral anions are polyatomic and are named based on the atomic constituents and the suffix – ide. 15. The most common examples are: a. OH- – hydroxide ion b. CN- – cyanide ion
40
16. A number of polyatomic anions containing oxygen atoms are named based on the root word of the central (or non-oxygen) atom and the suffix –ate for the one with more oxygen atoms and –ite for the one with less oxygen atom. a. NO3- – nitrate ion b. NO2- – nitrite ion c. SO32- – sulfite ion d. SO42- – sulfate ion e. PO43- – phosphate ion 17. Some anions have common names ending with the suffix –ate. a. C2H3O2- – acetate ion b. C2O42- – oxalate ion Point out that the composition of a molecule or an ion can be represented by a chemical formula. The formula consists of the symbols of the atoms making up the molecule. If there is more than one atom present, a numerical subscript is used. Examples are the following: a. O2
– oxygen gas
b. H2O
– water
c. NaOH – sodium hydroxide (liquid Sosa) d. HCl
– hydrochloric acid (muriatic acid)
18. Discuss that there are two types of chemical formulas: a. Molecular formula – gives the composition of the molecule, in terms of the actual number of atoms present. Examples are the following: i. C6H12O6 ii. K3PO4 iii. Na2C2O4 41
Teacher Tip They might be able to recall some compounds that have been presented in the Science course in junior high school, such as sodium chloride and carbon dioxide.
b. Empirical formula – gives the composition of the molecule, in terms of the smallest ratio of the number of atoms present. Examples are the following: i. CH2O ii. NaCO2 19. The naming of compounds follows a set of rules. Start the lesson with the rule of naming of binary compounds. Binary compounds – made up of two elements. Discuss the rules for naming in two groups of binary compounds:
Answers for Number 19 – sodium iodide i. NaI ii. MgCl2 – magnesium chloride – iron (II) sulfide iii. FeS iv. K2O – potassium oxide
a. Ionic compounds – made up of a cation and an anion. They are named by giving the name of the cation first, followed by the name of the anion. Ask them to name the following compounds: i. NaI ii. MgCl2 iii. FeS iv. K2O b. Molecular compounds – made up of two non-metals. They are named by giving the name of the first nonmetal and then that of the second nonmetal modified by the ending ide. Molecular compounds are usually gases. Ask them to name the following compounds: i. HCl ii. CO2 iii. SO3 20. After they have learned how to name binary compounds, discuss the rules for naming ternary compounds – made up of three elements. The naming of ternary compounds follows the same rule as that of the binary ionic compound: the name of the cation is given first, followed by the name of the anion.
42
Answers for Number 20 i. HCl – hydrogen chloride ii. CO2 – carbon dioxide iii. SO3 – sulfur trioxide
Answers for Number 20 i. NaNO3 – sodium nitrate ii. BaCrO4 – barium chromate iii. K2SO4 – potassium sulfate
Ask them to name the following compounds: i. NaNO3 ii. BaCrO4 iii. K2SO4 21. Discuss next the naming of acids. Acids – yield hydrogen ions in aqueous solutions. a. Binary acids – composed of hydrogen and another element, usually a nonmetal. The first part of the name starts with the prefix hydro- followed by the name of the element, modified by the ending –ic. The second part consists of the word ‘acid’. Name = hydro(root name of element) -ic + acid
Answers for Number 21.a i. HCl – hydrochloric acid ii. H2S – hydrosulfuric acid iii. HI – hydroiodic acid
Ask them to name the following binary acids: i. HCl ii. H2S iii. HI b. Ternary acids – made up of hydrogen and an anion, usually containing oxygen. The first part of the name consists of the root word of the name of the element, modified by the ending –ic. The second part consists of the word ‘acid’. If there is another acid with the same atoms, the suffix –ous is used to denote the one with less number of atoms. Name = (root name of element) -ic (or –ous) + acid Ask them to name the following ternary acids: i. HNO3 ii. HNO2 iii. H2SO4 iv. H2SO3 v. H3PO4 43
Answers for Number 21.b i. HNO3 – nitric acid ii. HNO2 – nitrous acid iii. H2SO4 – sulfuric acid iv. H2SO3 – sulfurous acid v. H3PO4 – phosphoric acid
22. After they have become familiar with the naming of compounds, it would be easy to write the formula of the compound. Emphasize that in writing the formula, the total positive charges of the cations should be equal to the total of the negative charges of the anion. The net charge should be zero.
Answers for Number 22 – AgNO3 i. Silver nitrate ii. Potassium iodide – KI iii. Nitrogen dioxide – NO2 iv. Barium chloride – BaCl2 v. Hydrobromic acid – HBr
Ask them to write the formula of the following compounds, given the name of the compound: i. Silver nitrate ii. Potassium iodide iii. Nitrogen dioxide iv. Barium chloride v. Hydrobromic acid
ENRICHMENT
1. Conduct a laboratory session on the naming of compounds and on formula writing.
EVALUATION (20 minutes) Check-up Quiz
Choose the best answer from among the choices given: 1. In one experiment, 0.558 g element X was found to react with 0.320 g element Y to form only one product, compound Z. How many grams of compound Z were formed? A. 0.238 g
C. 0.558 g
B. 0.320 g
D. 0.878 g
2. When 24.3 g magnesium reacts completely with 16.0 g oxygen, exactly 40.3 g magnesium oxide is formed. Which of the following laws is illustrated by this observation? A. Law of Definite Proportion
C. Law of Conservation of Mass
B. Law of Multiple Proportion
D. Law of Conservation of Energy 44
Teacher Tip Refer to the laboratory teaching guide of this lesson as well as the Formula Writing and Naming of Compounds data table.
3. Which of the following statements is consistent with Dalton’s Atomic Theory? A. The atoms of element A are identical with the atoms of another element D. B. The atoms of element A have the same mass as the atoms of another element D. C. The atoms of element A are different from the atoms of another element D. D. The atoms of element A have the same properties as the atoms of another element D. 4. According to Dalton’s atomic theory, which of the following is involved in a chemical reaction? A. The conversion of one atom into another
C. The formation of a new atom
B. The combination of atoms
D. The disappearance of an atom
5. Which of the following subatomic particles has the smallest mass? A. Electron
C. Proton
B. Neutron
D. Nucleus
6. In which of the following quantities will two isotopes of an element have different values? A. Atomic number
C. Number of protons
B. Mass number
D. Number of electrons
7. Which of the following information on the number of protons (p), electrons (e) and neutrons (n) is correct for 92 U238? A. 92 p, 92 n, 92 e
C. 238 p, 146 n, 238 e
B. 92 p, 146 n, 92 e
D. 146 p, 82 n, 92 e
8. What is the mass number of an atom which has 11 protons, 11 electrons, and 12 neutrons? A. 11
C. 22
B. 12
D. 23 45
9. Which of the following data is correct for the Mg2+ ion (atomic number = 12)? A. 12 protons and 13 electrons
C. 14 protons and 12 electrons
B. 12 protons and 10 electrons
D. 12 protons and 14 electron
10. Which of the following symbols corresponds to the element tin? A. Ti
C. Pb
B. Zn
D. Sn
11. Which of the following takes place when a monovalent cation is formed from an atom? A. One electron is gained.
C. Two electrons are gained.
B. One electron is lost.
D. Two electrons are shared.
12. Which of the following anions is polyatomic? A. Iodide
C. Sulfide
B. Nitrite
D. Bromide
13. Which of the following is the correct formula of copper (II) nitrate? A. CuNO3
C. Cu(NO3)2
B. Cu2NO3
D. Cu2(NO3)2
14. Which of the following is a binary compound? A. Sodium nitrate
C. Sodium hydroxide
B. Sodium oxide
D. Sodium carbonate
46
General Chemistry 1
90 MINS
Lesson 6: Atoms, Molecules, and Ions (Laboratory) Content Standard The learners demonstrate an understanding of the formula and the name of compounds. Performance Standard The learners shall be able to:
Lesson Outline
Introduction
Review
Motivation
Names and Formulas of Compounds
Practice
Activity
Enrichment
Discussion of Answers
Materials Exercise sheets
1. Write the formula and give the name of simple compounds. Learning Competency
Resources (1) Chang, R. & Goldsby, K. (2016). Chemistry (12th ed.). New York: McGraw-Hill.
At the end of the lesson, the learners: 1. Write the chemical formulas of ionic compounds and name ionic compounds from their formulas (STEM_GC11AM-Ic-e-24).
47
15 5 70
INTRODUCTION (15 minutes)
1. Reiterate to the learners the importance of the names and formulas of compounds. Make them recall the basic rules involved in formula writing and chemical nomenclature. 2. Review the symbols of the common elements encountered in compound. 3. State the objective of the exercise that they will work on for the laboratory period.
MOTIVATION (5 minutes)
1. Point out that the names and formulas of compounds will be needed in the succeeding lessons, particularly in writing chemical equations.
PRACTICE (70 minutes)
1. Provide each of them a copy of the exercise worksheet, and ask them to answer the exercise.
ENRICHMENT
Teacher Tip Point out that the formula gives qualitative and quantitative information about the composition of a compound. It shows what elements make up the compound (qualitative information) and the mole ratio of the elements (quantitative information). Call the learners one by one and ask him/her to give the symbol of an element which you will name. Teacher Tip The exercise worksheet given in Annex 1 could be adopted or revised. Each learner will work independently. It might be best to keep the exercise as a closed-book activity, and discourage them from consulting one another. At the end of the exercise, let them check the answers of their fellow learners who are seated away from them.
1. Ask them to identify where they committed mistakes. Discuss the correct answers.
EVALUATION EXCEEDS EXPECTATIONS
MEETS EXPECTATIONS
The learner answered more than 90% of the items correctly.
The learner answered 70% to 89% of the items correctly.
NEEDS IMPROVEMENT The learner answered less than 70% of the items correctly.
48
NOT VISIBLE The learner did not answer any item correctly.
FORMULA WRITING AND NAMING OF COMPOUNDS Section 1: Ion names
Section 2: Ions from formulas
Complete the table by writing the name or formula of the ionic species.
Complete the chart by writing the formula of the ions and of the compounds.
ION
NAME
COMPOUND
Na+
KCl
Ca2+
Ba(NO3)2 magnesium ion
FeSO4
manganese (II) ion
Li2CO3
Fe3+
Na2O chromium (III) ion
(NH4)2SO4
Ba2+
Al(OH)3
ClNO3phosphate ion OHchromate ion C2O42permanganate ion
49
POSITIVE ION
NEGATIVE ION
Section 3: Writing formulas from chemical names
Section 4: Chemical names from formulas
Write the formula of the ions expected from the following compounds.
Write the chemical name of the ions expected from the following compounds.
COMPOUND
POSITIVE ION
NEGATIVE ION
FORMULA
FORMULA
Calcium sulfate
ZnCl2 K3PO4
Potassium chloride
Cu(NO3)2 Na2CrO4
Tin (IV) oxide
Ni(OH)2
Lead iodide
BaO
Bismuth nitrate
(NH4)2C2O4
Sodium carbonate Strontium chromate
50
POSITIVE ION
NEGATIVE ION
NAME
Section 5: Binary covalent compounds
Section 6: Acids and bases
Complete the table below by filling up the missing formula or chemical name.
Complete the table below by filling up the missing formula or chemical name.
FORMULA
NAME
FORMULA
NO2
NAME hydroiodic acid
phosphorus trichloride
potassium hydroxide
carbon monoxide
HClO
SbBr5
H2S sulfur tetraiodide
perchloric acid
hydrogen peroxide
Zn(OH)2
P2O5
H3PO4 silicon dioxide
nickel(II) hydroxide
nitrogen trifluoride
sulfuric acid
CI4
HNO2 Mg(OH)2 carbonic acid
51
General Chemistry 1
60 MINS
Lesson 7: Atomic Mass Content Standard
Lesson Outline
The learners demonstrate an understanding of the mole concept in relation to Avogadro’s number and mass.
Introduction
Communicating Learning Objectives
Performance Standards
Motivation
Activity: Counting by Weighing
10
The learners shall be able to design, using multimedia, demonstrations, or models, a representation or simulation of any of the following:
Instruction
Average Atomic Mass
35
1. Atomic structure
Enrichment
Vitamins and Minerals
2. Mass relationships in reactions
Evaluation
Check Up Quiz
Learning Competency At the end of the lesson, the learners: 1. Explain relative atomic mass and average atomic mass (STEM_GC11SIe-25). Specific Learning Outcomes At the end of the lesson, the learners shall be able to: 1. Define atomic mass unit; 2. Calculate the average atomic mass of elements; 3. Determine the average molecular mass of molecules; and 4. Determine the average formula mass of ionic compounds.
52
5
10
Resources (1) Burdge, J & Overby, J. (2012). Chemistry: Atoms first. New York: McGraw-Hill. (2) Chang, R. &Goldsby, K. (2016). Chemistry. (12th ed.). New York: McGraw-Hill. (3) Isotopes and atomic mass [Simulation]. Retrieved from Phet Interactive Simulations website: https://phet.colorado.edu/en/simulation/ isotopes-and-atomic-mass (4) Moore, J.W., Stanitski, C.L. &Jurs, P.C. (2012). Chemistry: The molecular science (4th ed.). Belmont, CA: Brooks Cole/Cengage Learning. (5) Zumdahl, SS. & Zumdahl, S.A. (2012). Chemistry: An atoms first approach. Belmont, CA: Brooks/Cole Cengage Learning.
INTRODUCTION (5 minutes)
1. Introduce the following learning objectives using any of the suggested protocol (Verbatim, Own Words, or Read-aloud): a. Define atomic mass unit b. Calculate the average atomic mass of elements c. Determine the average molecular mass of molecules d. Determine the average formula mass of ionic compounds 2. Present the keywords for the concepts to be learned: a. Atomic mass unit (amu) b. Average atomic mass c. Molecular mass d. Formula mass e. Avogadro’s number f.
Mole
3. Review isotopes a. What are isotopes? b. Give examples of isotopes. c. What is the similarity between Mg-24 and Mg-25? What is their difference?
MOTIVATION (10 minutes) Activity: Counting by Weighing
1. Ask the learners if they can count objects by weighing them. Present to them this situation: Ms. Lilia sells shelled peanuts in a store. But she meets customers asking for 150 peanuts, another for 750 peanuts, and another for 2,000 peanuts. Obviously, it will take Ms. Lilia a very long time to count the peanuts. What would be another way to count them?
53
Answers for Number 3 a. Isotopes are atoms that have the same number of protons but different number of neutrons.) b. Here are some examples of isotopes: i. U-235 and U-238 ii. O-17 and O-18 iii. Kr-80, Kr-82, Kr-83 c. Mg-24 and Mg-25 both have 12 protons. However, Mg-24 has 12 neutrons while Mg-25 has 13 neutrons.
Ms. Lilia takes 20 peanuts and weighs them. She finds out that 20 peanuts weigh 32 g. How much then will each peanut weigh?
Teacher Tip Display the question clearly. Present the problem to the class. This activity can be done as a class or in groups. Give them about five minutes to reflect on the problem. Then, guide them to the process and the answer using the example given.
Hence the weight of 150 peanuts would be:
Take note and emphasize that not all the peanuts will have the same mass of 1.6 g. They are not all identical. Some will be heavier while some will be lighter. What was done was to get the average mass of the peanut and 1.6 g is the average mass of a peanut. However, for purposes of counting, what is needed is only the average mass.
It will be easier to weigh the peanuts than to count them. Now, 960 g is appropriately how many peanuts?
This method of counting by weighing is useful for counting very small objects, e.g. small candies, beans, etc.
Ask them to draw a conclusion. Is it possible to count objects by weighing? Summarize the procedure done with the peanuts. This can be done with other objects like mongo beans, marbles, etc.
54
Answer Key The procedure is as follows: 1. Count a given number of peanuts and weigh them. 2. Get the average mass of a peanut. This assumes that the objects are identical. 3. Divide the mass of a sample of peanuts by the average mass to get the number of peanuts in the sample.
INSTRUCTION (35 minutes) What is the Atomic Mass and the Atomic Mass Unit? Relate the exercise on counting peanuts by weighing to counting atoms. Ask them if it is possible to use the same procedure to count atoms. Why or why not? Whether it is peanuts or mongo beans or candies or atoms, the procedure should be the same. The problem, however, is atoms are very, very small and it is not possible to see them and count them individually to get the average mass. We need to look for another way to get the average mass of the atom.
Teacher Tip Ask them to check the meaning of the word relative when used as an adjective. Ask them to provide their source of information. Then, let them express the meaning in English and in Filipino. Briefly, relative, when used as an adjective, means ‘compared to something or to someone’. Emphasize that they should always use the appropriate unit in calculations. For atomic mass, the unit is amu.
Additional information: A mass spectrometer is used to experimentally compare and determine the masses of atoms to a very high degree of accuracy.
Experiments have shown that atoms have different masses relative to one another. For example, a Mg atom is experimentally reported to be twice as heavy as a carbon atom; a silicon atom is twice the mass of a nitrogen atom. It is possible to make a relative scale if one atom is chosen as the reference or standard atom against which the masses of the other atoms are measured.
Answer key 6.410 x 12 amu = 76.92 amu
By international agreement, the reference atom chosen is the C-12 isotope which contains six protons and six neutrons. By definition, one atom of C-12 has a mass of exactly 12 atomic mass units (amu). One amu, therefore, is one-twelfth (1/12) the mass of a C-12 atom. The atomic mass of Cu-63 is 62.93 amu. This means that relative to C-12, one atom of Cu-63 is 62.93/12 or 5.244 times the mass of a C-12 atom. Ask them to answer this example: One atom of Se-77 is 6.410 times as heavy as an atom of C-12. What is the atomic mass of Se-77? 55
Average Atomic Mass Now, ask them to look up the atomic mass for carbon in the periodic table. The expected answer is 12.01 amu. Then, proceed to explaining the average atomic mass.
Teacher Tip They should all have the same version of the periodic table so that average atomic masses are reported with the same number of significant figures.
If C has six protons and six neutrons, why is the relative atomic mass of carbon given as 12.01 amu and not 12 amu? There are no individual atoms of carbon with a mass of 12.01 amu. The periodic table provides the average atomic mass which takes into account the different isotopes of an element and their relative abundances. It is not a simple average that is taken but a weighted average.
Ask them to look up the atomic masses of other elements to familiarize them with using the periodic table.
Illustrate a weighted average using final grade calculation: For the class in Chem 345, the teacher informs the class that the final grade will be based on Exam 1 (15%), Exam 2 (15%), Problem Sets (30%), and Final Exam (40%). To pass the course, the learner must get a final grade of 75% or higher. Calculate the final grade of learner Ms. Julita if she got the following scores: COMPONENTS OF FINAL GRADE
WEIGHT
SCORES OF MS. JULITA
Exam 1
15.0%
83%
Exam 2
15.0%
95%
Problem Sets
30.0%
65%
Final Exam
40.0%
88%
The final grade will be computed as follows: (.150 x .83) + (.150 x .95%) + (.300 x .65) + (.400 x .88) = 81% Therefore, Ms. Julita passes the course! 56
You may want to show the difference between simple average and weighted average using the same values in the example given.
Always observe the use of significant figures in calculations.
Isotopes of elements occur in different abundances. Some are more abundant than others. Chlorine has two isotopes. The natural abundance of Cl-35 is 75% while that of Cl-37 is 25%. This means that if you have 100 atoms of chlorine, 75 of them will be Cl-35 and 25 of them will be Cl-37. Magnesium, on the other hand, has three isotopes with varying abundances: Mg-24, Mg-25, and Mg-26, 11.01 have 78.99%, 10.00%, and 11.01% abundance, respectively.
Teacher Tip Note that the atomic mass of C-12 is exactly 12 amu. In calculations, this is treated as an exact number.
For carbon, the natural abundance of C-12 is 98.90% while that of C-13 is 1.10%. The atomic mass of C-13 has been determined to be 13.00335 amu while that of C-12 is exactly 12 amu. Now, we calculate the average atomic mass of carbon:
Misconception They may think that there is a carbon atom with a mass of 12.01 amu. There is none. There are only atoms of C-12 and C-13. The value 12.01 amu is an average atomic mass.
Review how exact numbers are treated in calculations.
= (atomic mass of C-12) (% abundance of C-12) + (atomic mass of C-13) (% abundance of C-13)
= (12.0000 amu) (.9890) + (13.00335 amu) (.0110) = 12.01 amu
Ask them to answer these practice exercises: 1. From the periodic table, look up the average atomic mass of the following elements: Co, Be, Al, Zn. 2. Copper has two stable isotopes with the following masses and % abundances: Cu-63 (62.93 amu, 69.09% abundance) and Cu-65 (64.9278 amu, 30.91% abundance). Calculate the average atomic mass of copper. 3. An element consists of an isotope with mass of 10.0129 amu and 19.91% abundance, and another isotope with mass of 11.0093 amu and 80.09% abundance. Calculate the average atomic mass of this element. Refer to the periodic table and identify the element.
57
Answer Key 1. Co (58.93 amu), Be (9.012 amu), Al (26.98 amu), Zn (65.39 amu) 2. 63.55 amu 3. 10.81 amu; the element is boron Ensure that they observe the proper use of significant figures in all their calculations.
Average Molecular Mass (also referred to as molecular mass) The molecular mass is the sum of the average atomic masses of the atoms in the molecule. Ask them to answer the following examples: 1. What is the molecular mass of carbon dioxide, CO2? 2. Determine the molecular mass of the following molecules:
Teacher Tip Note the difference between molecular mass and formula mass. Molecular mass is used for covalent compounds while formula mass is used for ionic compounds. For brevity, many books refer to the average molecular mass as simply molecular mass.
a. Water, H2O b. Methane, CH4
Answer Key 1. Molecular mass of CO2 = atomic mass of C + 2 (atomic mass of O) = 12.01 amu + 2 (16.00 amu) = 44.01 amu 2.
Average Formula Mass (also referred to as formula mass) The formula mass is the sum of the atomic masses of the atoms in the ionic compound.
What is the formula mass of sodium chloride, NaCl?
2.
What is the formula mass of magnesium chloride, MgCl2?
Teacher Tip Keep the examples simple. It is the concept that needs to be introduced. This will be taken up again in the next lesson. Answer key 1. Formula mass of NaCl = atomic mass of Na + atomic mass of Cl) = 22.99 amu + 35.45 amu = 58.44 amu 2. Formula mass of MgCl2 = 95.21 amu
Ask them to answer the following examples: 1.
a. molecular mass of water = 18.02 amu b. molecular mass of methane = 16.04 amu
58
Vitamins and minerals 1. Vitamins and minerals are nutrients for the body. An example of a vitamin is Vitamin C. Look up the molecular formula of Vitamin C and determine its average molecular mass. What is another common name for Vitamin C? Give at least one important use of Vitamin C in the body.
Teacher Tip This can be given as an assignment.
2. Minerals include potassium, calcium, iron, and zinc. Look up the average atomic mass of calcium, Ca. Give at least one important use of Ca in the body.
EVALUATION (10 minutes) Check-up quiz
Answer the following questions. Place the answers in the space provided. Show calculations where applicable. Observe the use of significant figures for calculations and indicate the appropriate units. Learners can use the periodic table to answer the questions. ______1. From the periodic table, look up the average atomic mass of bromine, Br. ______2. How much heavier is an atom of Br relative to an atom of carbon? ______3. Which element in the periodic table has an average atomic mass that is about ten times that of fluorine? Element A consists of isotope A-6 with natural abundance of 7.5% and a mass of 6.0151 amu, and isotope A-7 with natural abundance 92.5% and mass of 7.0160 amu. ______4. Calculate the average atomic mass of element A. ______5. Identify Element A. Naphthalene has the molecular formula C8H10. ______6. How many elements make up one molecule of naphthalene? What are they? ______7. What is the molecular mass of naphthalene? ______8. What is the formula mass of potassium chloride, KCl? 59
Answer Key 1. 79.90 amu 2. 6.653 times heavier 3. Osmium, Os 4. 6.94 amu 5. Lithium, Li 6. Two elements; Carbon and Hydrogen 7. 106.16 amu 8. 74.55 amu
General Chemistry 1
120 MINS
Lesson 8: The Mole Concept and Molar Mass (Lecture) Content Standard
Lesson Outline
The learners demonstrate an understanding of the mole concept in relation to Avogadro’s number and mass.
Introduction
Communicating Learning Objectives
Performance Standards
Motivation
Inquiry
Instruction
The Mole Concept and Molar Mass
75
Enrichment
Relating the Mole to Real Life Situations
15
Evaluation
Check Up Quiz
15
The learners shall be able to design, using multimedia, demonstrations, or models, a representation or simulation of any of the following: 1. Atomic structure 2. Mass relationships in reactions Learning Competencies At the end of the lesson, the learners: 1. Define a mole (STEM_GC11S-Ie-26); 2. Illustrate Avogadro’s number with examples (STEM_GC11S-Ie-27); 3. Determine the molar mass of elements and compounds (STEM_GC11SIe-28); 4. Calculate the mass of a given number of moles of an element or compound, or vice versa (STEM_GC11S-Ie-29); and 5. Calculate the mass of a given number of particles of an element or compound, or vice versa (STEM_GC11S-Ie-30). Specific Learning Outcomes
At the end of the lesson, the learners shall be able to:
Resources (1) Burdge, J & Overby, J. (2012). Chemistry: Atoms first. New York: McGraw-Hill. (2) Chang, R. & Goldsby, K. (2016). Chemistry. (12th ed.). New York: McGraw-Hill. (3) Moore, J.W., Stanitski, C.L. & Jurs, P.C. (2012). Chemistry: The molecular science (4th ed.). Belmont, CA: Brooks Cole/Cengage Learning. (4) Zumdahl, SS. & Zumdahl, S.A. (2012). Chemistry: An atoms first approach. Belmont, CA: Brooks/Cole Cengage Learning.
3. Define molar mass; 4. Determine the molar mass of elements and compounds; and 5. Perform calculations determining mass of a given number of particles of an element or compound, or vice versa.
1. State the value of Avogadro’s number; 2. Perform calculations converting moles to number of entities and vice versa;
60
12 3
INTRODUCTION (12 minutes)!
1. Introduce the following learning objectives using any of the suggested protocol (Verbatim, Own Words, or Read-aloud): a. State the value of Avogadro’s number b. Perform calculations converting moles to number of entities and vice versa c. Define molar mass d. Determine the molar mass of elements and compounds e. Perform calculations determining mass of a given number of particles of an element or compound, or vice versa 2. Present the keywords for the concepts to be learned: a. Avogadro’s number b. Mole c. Molar mass 3. Discuss the quiz given during the previous lesson. Show the answers with the corresponding calculations. 4. Discuss the enrichment assignment given during the last lecture.
MOTIVATION (3 minutes)
1. What do chemists observe every October 23, from 6:02 am to 6:02 pm? Teacher Tip Ask the learners to wait until the end of the lesson to find the significance of the date and time.
INSTRUCTION (75 minutes) The Mole
What is a mole? What is Avogadro’s number?
Mole Day is observed by chemists every October 23, from 6:02 am to 6:02 pm.
Atoms have very small masses. We expect that macroscopic samples will contain a very large number of atoms. A special unit of measure, called the mole, is used to deal with extremely large numbers. In the SI system, the mole is the amount of substance that contains as many entities as there are in exactly 12 g of C-12.
61
The number of atoms in 12 g of C-12 is experimentally determined to be 6.022 x 1023. This is called Avogadro’s number. Illustrate the mole with different counting units:
Misconception Avogadro’s number is not a defined value. It is an experimentally determined value. Mass spectroscopy techniques are used to determine the value of Avogadro’s number to a high degree of accuracy.
a. 1 dozen = 12 entities or units b. 1 dozen eggs = 12 eggs
Teacher Tip Ask the learners to write all the zeroes for Avogadro’s number.
c. 1 dozen papayas = 12 papayas d. 1 dozen cars = 12 cars
Emphasize that while a dozen always has 12 entities, one dozen eggs will not have the same mass as one dozen books, or one dozen oranges, or one dozen cars.
e. 1 dozen books = 12 books f.
1 pair = 2 entities
g. 1 gross = 12 dozens = 144 entities h. 1 ream = 500 entities i.
1 mole = 6.022 x
1023
entities = 6.022 x
1023
Teacher Tip These are sample exercises showing conversion of moles to number of atoms or molecules and vice versa.
of anything
Ask them to answer the following practice exercises:
Answer Key 1. 6.022 x 1023 eggs 2. 6.022 x 1023 mongo beans 3. 6.022 x 1023 Na atoms
1. How many eggs are there in one mole of eggs? 2. How many mongo beans are there in one mole of mongo beans? 3. How many sodium atoms are there in 1 mole of Na atoms? 4. Calculate the number of atoms of argon in 0.500 moles Ar?
5. How many moles of Co are there in 4.960 x 1025 atoms of Co?
62
6. How many molecules of H2O are there in 1 mole of water molecules? 7. How many molecules of carbon dioxide, CO2, are there in 2.648 moles CO2?
Answer Key 6. 6.022 x 1023 H2O molecules Ask them to use the unit factor method (also called dimensional analysis) in their calculations.
8. How many atoms of oxygen are there in 2.648 moles CO2?
="3.189"x"10"24"O"atoms" 9. Determine the number of moles of ammonia, NH3, in 8.254 x 1025 molecules of ammonia.
="137.1"moles"NH3" Molar Mass Recall that the mole is the amount of substance that contains Avogadro’s number of units or entities. But how much will one mole of a substance weigh? The molar mass is the mass in grams of one mole of a substance. One mole of C-12 has a mass of exactly 12 g and contains 6.022 x C-12 is called the molar mass.
1023
C-12 atoms. This mass of
Notes: 1. The appropriate unit for molar mass is g/mol 2. The molar mass in grams is numerically equal to the atomic mass in amu. The molar mass in grams is numerically equal to the molecular mass or the formula mass in amu.
63
Teacher Tip Emphasize the use of the appropriate units in calculations. For molar mass, the unit used is g/ mol. For brevity, atomic mass is often used instead of average atomic mass. It is understood that the value in the periodic table is the average atomic mass.
Ask them to answer the following examples: 1. What is the average atomic mass of Ca? What is the molar mass of Ca? 2. The atomic mass of Br is 79.90 amu. What is its molar mass? 3. The molecular mass of water, H2O, is 18.02 amu. What is the molar mass of water, H2O? 4. The formula mass of NaCl is 58.44 amu. What is the molar mass of NaCl? Illustrate the relationship of amu and grams: One mole of C-12 has a mass of exactly 12 g and one mole of C-12 has Avogadro’s number of atoms. Calculate the mass of one atom of C-12 in grams.
Teacher Tip Emphasize the use of the appropriate units in calculations. For molar mass, the unit used is g/ mol. For brevity, atomic mass is often used instead of average atomic mass. It is understood that the value in the periodic table is the average atomic mass. Answer Key 1. 40.08 amu; 40.08 g/mol 2. 79.90 g/mol 3. 18.02 g/mol 4. 58.44 g/mol Recall the previous lesson on how to get the molecular mass and the formula mass.
Calculate the mass in grams of 1 amu.
Therefore, 1 amu = 1.661 x 10-24 g. Illustrate Avogadro’s number and molar mass:
SAMPLE
NUMBER OF ATOMS in SAMPLE
MASS of 1 mole
1 mole of aluminium
6.022 x 1023 atoms
26.98 g
1 mole of copper
6.022 x 1023 atoms
63.55 g
1 mole of silver
6.022 x 1023 atoms
107.9 g
1 mole of gold
6.022 x 1023 atoms
197.0 g 64
Teacher Tip Emphasize that while 1 mole Al, 1 mole Cu, 1 mole Ag, and 1 mole Au will each contain the same number of atoms, they will not weigh the same. Similarly, one dozen apples and one dozen cars will each have 12 units but will not weigh the same.
Illustrate how to get the molar mass of elements and compounds through the following examples: 1. Determine the molar mass of silicon, Si. 2. Get the molar mass of zinc, Zn. 3. Which will have a higher mass: 0.500 mole zinc, Zn, or 0.250 mole lead, Pb?
Answer Key 1. 28.09 g/mol 2. 65.39 g/mol 3. Therefore, 0.250 mole of Pb has a higher mass than 0.500 mole of Zn. 4. 78.12 g/mol 5. 45.07 g/mol 6. 73.89 g/mol
Therefore, 0.250 mole of Pb has a higher mass than 0.500 mole of Zn. 4. What is the molar mass of benzene, C6H6 ? 5. Find the molar mass of ethanol which has the following structural formula:
Teacher Tip Emphasize that they should give the answers with the appropriate units.
6. What is the molar mass of lithium carbonate, Li2CO3? Using the above illustrations, ask them to do calculations involving moles, molar masses, and Avogadro’s number. 1. How many grams of silver, Ag, are there in 1.34 moles? (This example illustrates the conversion of moles ! grams)
65
Teacher Tip Always observe the proper use of significant figures in calculations. An Annex is included at the end of this module for enrichment and review of significant figures and rounding off in calculations. Show how the proper use of units will facilitate the solution of the problem through dimensional analysis. The units cancel out, leaving the correct unit required.
2. How many moles of copper, Cu, are there in 875 g Cu? (This example illustrates the conversion of grams ! moles)
3. A bottle of calcium supplements in tablet form contains 268 g Ca. How many atoms are present in 268 g calcium, Ca? (This example illustrates the conversion of grams ! moles ! number of atoms)
4. What is the mass in grams of 2.06 x 1023 atoms of potassium, K? (This example illustrates the conversion of number of atoms ! moles ! grams)
5. Which has more atoms? 3.68 g neon atoms or 1.10 g sodium atoms?
Therefore, 3.68 g Ne will have more atoms than 1.10 g Na. 66
Teacher Tip Allow them to analyze the way to solve the problem using relationships before doing the calculation. In Problem 3, for example, the grams need to be converted to moles, then the moles converted to number of atoms.
ENRICHMENT (15 minutes)
Teacher Tip Return to the motivation question and ask them why October 23, from 6:02 AM to 6:02 PM, is the chosen date for Mole Day.
1. What do chemists observe every October 23 from 6.02 am to 6.02 pm? 2. Relate the mole to real life situations: A. How many pesos are there in one mole of pesos? Do you think Manny Pacquiao will have one mole of pesos? Does Bill Gates have one mole of dollars? B. Ask them to check the Philippine national budget for one fiscal year. Does the Philippine national budget reach one mole of pesos? C. Ask them to look for the approximate age of the earth. Does the age of the earth approximate one mole of years?
EVALUATION (15 minutes) Check-up quiz
Answer the following questions. Place the answers in the space provided. Show calculations where applicable. Observe the use of significant figures for calculations and indicate the appropriate units. Learners can use the periodic table to answer the questions
This enrichment could be done for the more advanced learners. Answer Key for Enrichment B. The Philippine National Budget for 2016 is PHP 3.002 trillion or 3,002,000,000,000 or 3.002 x 1012 pesos. The Philippine national budget does not reach one mole of pesos. C. Current data show the earth to be about 4.54 billion years old. It is 4,540,000,000 years old or 4.54 x 109 years old. The age of the earth does not approximate one mole of years. Answer Key for Evaluation 1. 20 x 1024 molecules 2. 41 x 1024 C atoms
_____1. How many molecules of acetylene, C2H4, are there in 2.00 moles acetylene? _____2. How many atoms of carbon are there in 2.00 moles acetylene? Complete the following table: SUBSTANCE
MOLES
NO
2.88 moles
GRAMS
CCl4
MOLECULES
121.4 g
SO2
8.50 x 1024 molecules
67
Answer Key SUBSTANCE
MOLES
GRAMS
MOLECULES
NO
2.88 moles
86.4 g
1.73 x 1024 molecules
CCl4
0.7893 mole
121.4 g
4.753 x 1023 molecules
SO2
14.11 moles
904 g
8.50 x 1024 molecules
Guidelines for Using Significant Figures (from Chang, R. & Goldsby, K. (2016). Chemistry. (12thed.). New York: McGraw-Hill, Chapter 1, pp. 20-21) In scientific work, significant figures are always to be observed. Here are the rules on the use of significant figures: 1. Any digit that is not zero is significant. (Eg. 483 g has three significant figures; 2,578 m has four significant figures) 2. Zeros between nonzero digits are significant. (Eg. 6.06 kg has three significant figures; 60,804 cm has five significant figures) 3. Zeros to the left of the first nonzero digit are not significant. (Eg. 0.078 L has two significant figures; 0.004 kg has one significant figure) 4. A. If a number is greater than 1, the zeros after the decimal point are significant. (Eg. 4.0 mg has two significant figures; 20.04 g has four significant figures) B. If a number is less than 1, only the zeros after the first nonzero digit are significant. (Eg. 0.0750 m has three significant figures; 0.4006 g has four significant figures. 5. For numbers without decimal points, the zeroes at the end of nonzero digits may or may not be significant (ambiguous). For example, 600 g may have one or three significant figures. To avoid the ambiguity, we use scientific notation. We can say 6.00 g and this will have three significant figures. Or we can say 6 x 102 and this will have only one significant figure.
68
How do you handle significant figures in calculations? 1. In addition and subtraction, the answer cannot have more digits to the right of the decimal point than either of the original numbers. 45.112
!
three digits after the decimal point
- 6.02
!
two digits after the decimal point
39.092
!
round-off to 30.09 so the answer will have two digits after the decimal point
2. For multiplication and division, the number of significant figures in the final product or quotient is determined by the original number that has the smallest number of significant figures. 6.9 x 12.34 = 85.146
Round of the answer to 85, which has only two significant figures.
26.98/3.05 = 23.93
Round of the answer to 23.9, which has three significant figures because the smallest number of significant figures in the operation is 3.
3. Remember that exact numbers are considered to have infinite number of significant figures.
Rules for Rounding Off: 1. To round off a number at a certain point, drop the digits that follow if the first of them is less than 5. 8.143 rounded off to only two significant figures becomes 8.1. 2.
To round off a number at a certain point, add 1 to the preceding digit if the number that follows is 5 or greater than 5. 7.378 rounded off to three significant digits becomes 7.38. 8.465 rounded off to three significant digits becomes 8.47. 0.575 rounded off to two significant digits becomes 0.58. 69
General Chemistry 1
120 MINS
Lesson 9: The Mole Concept and Molar Mass (Laboratory)
Lesson Outline
Content Standard The learners demonstrate an understanding of the mole concept in relation to Avogadro’s number and mass. Performance Standards The learners shall be able to design, using multimedia, demonstrations, or models, a representation or simulation of any of the following: 1. Atomic structure 2. Mass relationships in reactions Learning Competencies At the end of the lesson, the learners: 1. Define a mole (STEM_GC11S-Ie-26); 2. Illustrate Avogadro’s number with examples (STEM_GC11S-Ie-27); 3. Determine the molar mass of elements and compounds (STEM_GC11SIe-28); 4. Calculate the mass of a given number of moles of an element or compound, or vice versa (STEM_GC11S-Ie-29); and 5. Calculate the mass of a given number of particles of an element or compound, or vice versa (STEM_GC11S-Ie-30). Specific Learning Outcomes At the end of the lesson, the learners shall be able to: 1. Count the number of small objects by weighing; 2. Determine the number of moles in a given sample; and 3. Determine the number of atoms in a given sample 70
Introduction
Can you count objects by weighing them?
10
Instruction and Practice
Laboratory Work
80
Enrichment
Post-laboratory Discussion
20
Evaluation
Checking of Accomplished Data Tables
Materials (1) Balance (triple beam or electronic balance) (2) Paper cups (3) Samples (kidney beans, mongo beans, rice, dried sago) (4) Plastic spoons (5) Aluminium metal or foil (6) Iron (nails or filings) (7) Sodium chloride (table salt, NaCl) (8) Sucrose (table sugar, C12H22O11) Resources (1) Allan, Andy. The mole [PowerPoint presentation]. Retrieved from http://www.sciencegeek.net/APchemistry/FlashPPT/3_TheMole/ index.html (2) Burdge, J. & Overby, J. (2012). Chemistry: Atoms first. New York: McGraw-Hill. (3) Chang, R. & Goldsby, K. (2016). Chemistry (12th ed.). New York: McGraw-Hill. (4) Moore, J. W. & Stanitski, C.L. (2015). Chemistry: The Molecular Science (5th ed.). Belmont, CA: Brooks Cole/Cengage Learning. (5) Zumdahl, SS. &Zumdahl, S.A. (2012). Chemistry: An atoms first approach. Belmont, CA: Brooks/Cole Cengage Learning.
INTRODUCTION (10 minutes)
Can you count objects by weighing them? Ask the learners if they can count objects by weighing them. Present to them the following situations: 1. Ms. Lilia sells shelled peanuts in a store. But she meets customers asking for ten peanuts, another for 750 peanuts, and another for 2,000 peanuts. Obviously, it will take Ms. Lilia a very long time to count the peanuts. What would be another way to count them?
Teacher Tip 1. Prepare the classroom or laboratory, the materials, the laboratory sheets to be used. 2. Distribute the laboratory sheets at the start of the lesson. 3. After the introduction and motivation, explain the procedures of the activity. 4. Explain the safety precautions.
2. Mr. Jose goes to a hardware store and asks for 400 pieces of nails. What is an easier way to approximate 400 pieces of nails without counting them one by one? 3. A candy factory sells chocolate chips in a bag. Each bag should have the same number of chips. How does the candy factory count the number of chocolate chips in each bag?
INSTRUCTION and PRACTICE (80 minutes) Safety Precautions
Teacher Tip The activity can be performed individually or in groups. All materials are household materials. Nevertheless, caution must be observed in handling any material in the lab. Instruct them how to behave in the laboratory.
1. Never taste anything during a science activity. 2. Dispose of the samples as directed by your teacher. 3. Wash your hands with soap and water after the science activity. 4. Follow all laboratory instructions as directed by the teacher. Part I. Counting by weighing: Relating mass to number This method of counting by weighing is useful for counting very small objects, such as small candies, beans, etc. Ask them to perform the activity on relating mass to number of entities. The activity asks them to determine the number of entities in a given sample by weighing a given amount of sample and knowing the average mass of one entity. The activity uses common materials. Kidney beans or other beans such as peanuts, squash seeds, or broad beans or patani may be used for Sample 1. Use smaller sized samples like mongo, peas, or dried sago for Sample 2. Have them answer the data tables and the questions. See the attached laboratory sheet. 71
Take note and emphasize that not all the kidney beans will have the same mass since they are not all identical. Some will be heavier and some will be lighter. What was done was to obtain the average mass of the kidney bean. However, for purposes of counting, what is needed is only the average mass.
Part II. Relating mass to moles Ask them to perform the activity on relating mass to moles. The samples are common household materials: Aluminium, iron, sodium chloride, and sucrose. Have them answer the data tables and the questions. See the attached laboratory sheet. Sample of Teacher’s Reference Table A1
11 g
B1
14 g
A2
22 g
B2
27 g
A3
16 g
B3
18 g
A4
25 g
B4
30 g
1. For the Aluminium sample, crumple a sheet of Aluminium foil into a loose ball and place in a small paper cup. 2. Be careful in using iron nails. 3. At the end of the activity, instruct them where to place the samples. Put separate containers for each sample. 4. The samples may be reused for another class.
ENRICHMENT (20 minutes) Post-laboratory discussion
1. Relate counting by weighing to finding the number of atoms in a weighed sample of material. 2. Give more exercises on calculating moles and molar masses of elements and compounds.
EVALUATION
Check their accomplished data tables and worksheet for correct use of units and significant figures and the logical solutions.
72
Teacher Tip Place the samples in small paper cups prior to the class. There are four samples in this activity, and make sure to prepare enough samples for the class. Each group must work on all four samples. Label the cups (e.g. Sample A1, A2, A3, A4 for Group A, etc.). You must also pre-determine the approximate amount of sample to put in each cup and put these in your notes. This will serve as a reference for the masses measured by the learners. However, they must measure the masses up to .01 g. The masses do not have to be identical. For example, the mass of Al in one group may have a different value than the mass of Al in the other group.
LABORATORY ACTIVITY: THE MOLE CONCEPT AND MOLAR MASS Introduction Atoms have very small masses. Macroscopic samples contain a very large number of atoms. The mole is used to deal with these extremely large numbers of atoms in macroscopic samples. The mole is defined as the amount of substance that contains as many entities as there are in exactly 12 grams of C-12. This is experimentally determined to be 6.022 x 1023 and is referred to as Avogadro’s number. The molar mass is the mass in grams of one mole of a substance. It is possible and practical to count very small objects by determining an average mass then weighing a given sample. You will be asked to determine the number of entities in a given sample of material through this technique. You will also determine the number of moles of different substances and the corresponding number of atoms present in the sample. Objectives 1. To determine the number of entities present in a given sample by weighing it and identifying the average mass of a single entity of the sample. 2. To determine the number of moles and the number of atoms present in given samples of materials. Materials
1. Balance – triple beam or electronic balance
5. Aluminium metal or foil
2. Paper cups
6. Iron (e.g. iron nails or iron filings)
3. Samples – e.g. kidney beans, mongo beans, rice, dried sago
7. Sodium chloride (NaCl)
4. Plastic spoons
8. Sucrose (table sugar, C12H22O11)
Safety Precautions 1. Never taste anything during a science activity. 2. Dispose of the samples as directed by your teacher. 3. Wash your hands with soap and water after the activity. 4. Follow all laboratory instructions as directed by your teacher.
73
Part I. Counting by weighing: Relating mass to number Procedure
Sample 1
Sample 2
1. Count 20 beans (kidney, peanuts, patani, or other samples as given by your teacher) and place them in a paper cup.
1. Count 20 mongo beans (rice, dried sago, peas, or any smaller bean samples given by your teacher) and place them in a paper cup.
2. Determine the mass of the 20 pieces of beans. Remember to subtract the mass of the container. If using an electronic balance, tare or set the balance to zero.
2. Determine the mass of the 20 pieces of beans. Remember to subtract the mass of the container. If using an electronic balance, tare or set the balance to zero.
3. Determine the mass of one bean by dividing the mass of the sample by 20.
3. Determine the mass of one bean by dividing the mass of the sample by 20. Data Table SAMPLE 1
Sample Mass of 20 pieces of sample plus container Mass of container Mass of 20 pieces of sample Mass of one piece of sample (Show calculation here) Answer the following questions: 1. How much will 750 pieces of kidney beans weigh? 2. Calculate the mass of 5,500 mongo beans. 3. 158 grams of mongo beans is approximately how many pieces?
74
SAMPLE 2
Part II. Relating mass to moles Procedure: Determine the masses of Samples 1 to 4. Record these in the data tables provided. Calculate the number of moles in each sample. Show all calculations and observe the correct use of units and significant figures.
Sample 1
Sample 2
Sample 3
Sample 4
ALUMINIUM
IRON
SODIUM CHLORIDE (Table salt, NaCl)
SUCROSE (Table sugar, C12H22O11)
1. Mass of sample + container, g 2. Mass of container, g 3. Mass of sample, g 4. Molar mass of sample, g/mol 5. Number of moles in sample 6. Number of atoms in sample 7. No. of atoms in 1.0 gram of sample
75
General Chemistry 1
60 MINS
Lesson 10: Percent Composition and Chemical Formulas Lesson Outline
Content Standard The learners demonstrate an understanding of percent composition and chemical formulas. Performance Standards The learners shall be able to design, using multimedia, demonstrations, or models, a representation or simulation of any of the following: 1. Atomic structure 2. Mass relationships in reactions Learning Competencies At the end of the lesson, the learners: 1. Calculate the percent composition of a compound from its formula (STEM_GC11PC-If-31); 2. Calculate the empirical formula from the percent composition of a compound (STEM_GC11PC-If-32); and 3. Calculate molecular formula given molar mass (STEM_GC11PC-If-33). Specific Learning Outcomes
At the end of the lesson, the learners shall be able to: 1. Interpret the information provided by the chemical formula;
Introduction
Communicating Learning Objectives
5
Motivation
College Projections
5
Instruction
Percent Composition and Chemical Formula
Enrichment
Determine the Sodium Percent in Snack Food
Evaluation
Short Quiz
35
15
Resources (1) Burdge, J.& Overby, J. (2012). Chemistry: Atoms first. New York: McGraw-Hill. (2) Chang, R. & Goldsby, K. (2016). Chemistry. (12th ed.). New York: McGraw-Hill. (3) Moore, J.W., Stanitski, C.L. & Jurs, P.C. (2012). Chemistry: The molecular science (4th ed.). Belmont, CA: Brooks Cole/Cengage Learning. (4) Zumdahl, S.S. &Zumdahl, S.A. (2012). Chemistry: An atoms first approach. Belmont, CA: Brooks/Cole Cengage Learning.
3. Explain the procedure used to determine the empirical formula of a compound given the percent composition; and
2. Explain the procedure used to determine the percent composition of a compound;
4. Utilize molar mass data to obtain the molecular formula from the empirical formula.
76
INTRODUCTION (5 minutes)
1. Introduce the learning objectives using any of the suggested protocol (Verbatim, Own Words, or Read-aloud) a. Interpret the information provided by the chemical formula b. Explain the procedure used to determine the percent composition of a compound c. Explain the procedure used to determine the empirical formula of a compound given the percent composition
Teacher Tip Give examples of empirical formulas and molecular formulas such as a. Benzene, with molecular formula C6H6 and empirical formula CH b. Acetylene, with molecular formula C2H2 and empirical formula CH c. Ammonia, with molecular formula NH3 and empirical formula NH3.
d. Utilize molar mass data to obtain the molecular formula from the empirical formula 2. Present the keyword for the concepts to be learned: a. Percent composition mass 3. Review empirical formulas and molecular formulas and give examples.
MOTIVATION (5 minutes)
Percentage of the class planning (as first choice) to go to the different strands in STEM 1. Get total number of learners in class. 2. Get the number of learners who would like to pursue engineering in college. Get the percentage. 3. Get the number of learners who would like to pursue the sciences. Get the percentage. 4. Get the number of learners who would like to pursue mathematics. Get the percentage.
INSTRUCTION (35 minutes) 1. Information from the chemical formula What information can be obtained from a chemical formula? For example, what information can you get from the formula of carbon dioxide, CO2?
= CARBON
= OXYGEN
77
Teacher Tip This will give an indication of the interest of the learners and their planned careers. This will also review the concept of percentage and its application in the real world. Then, mention that the lesson will be about how the concept of percentage is used in chemistry, especially in chemical compounds.
a. The compound is made up of two elements, namely carbon and oxygen. b. One molecule of CO2 is made up of one atom of carbon and two atoms of oxygen. c. One mole of CO2 molecules will have one mole of C atoms and two moles of O atoms. d. The ratio of the moles of C to the moles of O in CO2 is 1:2. e. CO2 is composed of 27.29% carbon and 72.71% oxygen. The chemical formula provides the percent composition of CO2.
Teacher Tip You should motivate the learners to provide the answers instead of just stating them. It is important for them to understand the chemical formula and all information that can be obtained from it.
2. Percent Composition by Mass The percent composition by mass is the percent by mass of each element in a compound. Mathematically,
with n = the number of atoms of the element For CO2,!
The answer indicates that CO2 is composed of 27.29% C atom and 72.71% O atom. Ask them to answer the following practice exercises: I.
Calculate the percent composition of NaCl.
II. The chemical formula of glucose is C6H12O6. Determine its percent composition. III. Which element comprising Mg(OH)2 has the highest percentage by mass? 78
Answer Key 1. 39.34% Na, 60.66% Cl 2. 39.99% C, 6.727% H, 53.28% O 3. O; the composition is 41.68% Mg, 54.89% O, and 3.46% H Teacher Tip You may want to connect the lesson to some real world examples. Magnesium hydroxide, Mg(OH)2, is used as a medication to treat symptoms brought about by too much stomach acid such as heartburn or indigestion.
3. Empirical Formula from Percent Composition The empirical formula of a compound can be calculated from the percent composition. Because percentage is given, it is convenient to assume 100.00 grams of the compound. Illustrate using the following examples: A. A compound is found to consist of 7.81% C and 92.19% Cl. What is the empirical formula of the compound? Assume 100.00 grams of the compound. The sample will therefore contain 7.81 g C and 92.19 g Cl. The grams are converted to moles to get the ratios of the moles of the elements in the compound:
The compound is C0.650Cl2.601. But chemical formulas are expressed in whole numbers. Empirical formulas are expressed as the lowest whole number ratio between the atoms. To convert to whole numbers, divide the number of moles by the smallest value (that is 0.650).
The empirical formula is C1Cl4 or CCl4.
79
B. A compound is found to consist of 43.64% P and 56.36% O. The molar mass for the compound is 283.88 g/mol. What is the empirical formula and molecular formula of the compound? Assume 100.00 grams of the compound. What is the mass of each element in 100.00 grams of compound?
What are the moles of each element in 100.00 grams of compound?
Divide the mole values by the smallest value to get
The compound is PO2.5. But the subscripts are still not whole numbers. Multiply the subscripts by a factor to get the smallest whole number. When multiplied by 2, the empirical formula is P2O5. What is the molecular formula? Compare the mass of the empirical formula to the molar mass: Mass of P2O5 = 141.94 g/mol Molar mass = 283.88 g/mol
80
Answer Key 1. 39.34% Na, 60.66% Cl 2. 39.99% C, 6.727% H, 53.28% O 3. O; the composition is 41.68% Mg, 54.89% O, and 3.46% H
Therefore, the molecular formula is (P2O5)2 or P4O10.
ENRICHMENT
Determine the % sodium in snack food
Look at the food labels of some snack food like potato chips, peanuts, popcorn, etc. Fill up the table below.
Teacher Tip You may want to connect the lesson to some real world examples. Magnesium hydroxide, Mg(OH)2, is used as a medication to treat symptoms brought about by too much stomach acid such as heartburn or indigestion.
1. Identify your chosen snack food and brand. 2. Get the amount in grams of one serving of the snack food. 3. Get the amount of sodium in mg in one serving of the snack food. 4. Obtain the % sodium by mass in one serving of snack food.
EVALUATION (15 minutes)
Answer the following questions. Place the answers in the space provided. Show calculations where applicable. Observe the use of significant figures for calculations and indicate the appropriate units. Learners can use the periodic table to answer the questions. Aspirin has the molecular formula C9H8O4. _____1. What is the % C in aspirin by mass in aspirin? _____2. What is the % O in aspirin by mass in aspirin? _____3. An oxide of chromium is made up of 5.20 g chromium and 5.60 g oxygen What is the empirical formula of the oxide? (Note: An oxide of nitrogen contains 63.1% oxygen and has a molar mass of 76.0 g/mol.) _____4. What is the empirical formula for this compound? _____5. What is the molecular formula of the compound?
81
Answer Key 1. 60.00% 2. 35.53% 3. Cr2O7 4. N2O3 5. N2O3
General Chemistry 1
120 MINS
Lesson 11: Chemical Reactions and Chemical Equations (Lecture) Content Standard The learners demonstrate an understanding of the use of chemical formulas to represent chemical reactions.
Lesson Outline Introduction
Review and Communicating Learning Objectives
Motivation
Evidences of Chemical Change
Instruction
Chemical Reactions and Chemical Equations
Enrichment
Inquiry
Performance Standards The learners shall be able to design, using multimedia, demonstrations, or models, a representation or simulation of any of the following: 1. Atomic structure 2. Mass relationships in reactions Learning Competencies At the end of the lesson, the learners: 1. Write equations for chemical reactions and balance the equations (STEM_GC11CR-If-g-34); 2. Interpret the meaning of a balanced chemical reaction in terms of the Law of Conservations of Mass (STEM_GC11CR-If-g-35); 3. Describe evidences that a chemical reaction has occurred (STEM_GC11CRIf-g-36); and 4. Perform exercises on writing and balancing chemical equations (STEM_GC11CR-If-g-37). Specific Learning Outcomes At the end of the lesson, the learners shall be able to: 1. Write and balance chemical equations; 2. Derive pertinent information from a balanced chemical equation; 3. Determine whether a chemical reaction has occurred or not; and 4. Classify chemical reactions. 82
30 3 85 2
Resources (1) Burdge, J & Overby, J. (2012). Chemistry: Atoms first. New York: McGraw-Hill. (2) Chang, R. & Goldsby, K. (2016). Chemistry. (12th ed.). New York: McGraw-Hill. (3) Chemical equations [Online lecture]. Retrieved from CK-12 website: https://www.ck12.org/physical-science/Chemical-Equations-inPhysical-Science/ (4) Moore, J.W., Stanitski, C.L.& Jurs, P.C. (2012). Chemistry: The molecular science (4th ed.). Belmont, CA: Brooks Cole/Cengage Learning. (5) Recognizing chemical reactions [Online lecture]. Retrieved from CK-12 website: https://www.ck12.org/physical-science/RecognizingChemical-Reactions-in-Physical-Science/ (6) Zumdahl, SS. & Zumdahl, S.A. (2012). Chemistry: An atoms first approach. Belmont, CA: Brooks/Cole Cengage Learning.
INTRODUCTION (30 minutes) Review
1. Discuss the enrichment assignment of last lesson (% sodium in snack food). 2. Discuss the quiz on percent composition given in the last lesson. Communicating Learning Objectives 1. Introduce the following learning objectives using any of the suggested protocol (Verbatim, Own Words, or Read-aloud): a. Write and balance chemical equations b. Derive pertinent information from a balanced chemical equation c. Determine whether a chemical reaction has occurred or not
Teacher Tip Ask the learners to discuss what they found to be the % sodium in their snack samples. Alternatively, these can be written on the board. Call on some of them to show their calculations on the board. Have the class comment on the calculations, including the proper use of significant figures.
Teacher Tip List these keywords on the board or through PowerPoint slides. Alternatively, you can write them on flip charts.
d. Classify chemical reactions 2. Present the keywords for the concepts to be discussed: a. Chemical equation b. Reactant c. Product d. Aqueous e. Decomposition reaction f.
Synthesis reaction
g. Single displacement reaction h. Double displacement reaction i.
Combustion reaction
j.
Hydrocarbon Note Ask them to recall the Law of Conservation of Mass and express their understanding of it.
3. Review the Law of Conservation of Mass
83
MOTIVATION (3 minutes)
Teacher Tip You may bring some actual samples of rusty iron nails, bleached and unbleached hair, or other materials to show to the class.
Ask them what the following have in common: a. Rusty iron nail b. Change in color of leaves
These are evidences of chemical change.
c. Bleached hair
INSTRUCTION (85 minutes)
Writing and Balancing a Chemical Equation In a chemical reaction, a substance (or substances) is converted to one or more new substances. Chemical reactions follow the law of conservation of mass. No atoms are created or destroyed; they are just rearranged. Chemists have a way of communicating chemical reactions. They represent chemical reactions through chemical equations. Consider the reaction of hydrogen gas (H2) with chlorine gas (Cl2) to yield hydrogen chloride. The reaction is illustrated by the figure below.
Teacher Tip Emphasize the Law of Conservation of Mass. There must be the same type of atoms on both sides of the arrow.
We can represent this reaction through a chemical equation. The reactants (starting substances) are placed on the left side. The products (substances produced) are placed on the right. An arrow points towards the direction of the reaction. The equation has to be balanced so that the same number and types of atoms appear on the left and right side of the equation. To balance, coefficients (numbers preceding the chemical formula) are used. For additional information, the physical states of the reactants and products (s, l, g, for solid, liquid, or gas, respectively)are indicated. 84
Hence, the balanced chemical equation is:
Teacher Tip Show learners where to put the coefficients.
H2(g) + Cl2(g) ! 2 HCl(g) Check if the equation is balanced: Reactants
Products
H (2)
H (2)
Cl (2)
Cl (2)
When a substance is placed in water, we indicate this with aq, meaning it is in an aqueous environment. For example, when KBr reacts with AgNO3 in an aqueous environment, KNO3 and solid AgBr are produced. This reaction is represented as KBr(aq) + AgNO3(aq) ! KNO3(aq) + AgBr(s) Show them the procedure of balancing equations through this example: Ethane (C2H6) reacts with oxygen gas (O2) to produce carbon dioxide and water. Write the balanced chemical equation for the reaction. 1. Identify reactants and products and write their correct formulas. Put reactants on the left side and products on the right. C2H6 + O2 ! CO2 + H2O
85
2. Balance the equation by changing the coefficients of the reactants or products. Do not change the subscripts or the chemical formula. C2H6 + 7/2 O2 ! 2CO2 + 3 H2O To use the smallest whole number coefficients, we multiply the equation by 2 to give: 2C2H6 + 7O2 ! 4CO2 + 6H2O 3. Check to make sure that the number of each type of atom is the same on each side of the equation. Reactants
Products
4
4
C
C
12 H
12 H
14 O
14 O
Ask them to answer these exercises: State if the illustrated equation below is balanced or not. If not, explain why it is not balanced. Illustrate by a drawing how you would balance the equation. 1.
2.
3. 86
Teacher Tip Show learners where to put the coefficients.
Balance the following equations
Teacher Tip Show them where to put the coefficients.
1. ____ C + ____ O2 ! ___ CO
Hint: Start with elements that appear only once on each side.
2. ____ Mg + ____ O2 ! ____ MgO 3. ____ H2O2 ! ____ H2O + ____ O2 4. ____ CH4 + O2 ! ____ CO2 + ____ H2O 5. ____ N2O5 ! _____ N2O4 + _____ O2 Interpretation of a Chemical Equation How can a balanced chemical equation be interpreted? See the example: H2
+
Cl2
!
2HCl
One molecule
+
One molecule
!
Two molecules
One mole
+
One mole
!
Two moles
2 (1.008 g) = 2.016 g
+
2 (35.45g) = 70.90 g
!
Ask them to answer 72.929 this exercise:
2 (1.008 g + 35.45 g) = 72.92 g 72.92 g
Interpret the balanced equation: 2C2H6 + 7O2 ! 4CO2 + 6H2O Show that the Law of Conservation of Mass is followed.
87
Answer Key 1. 2, 1, 2 2. 2, 1, 2 3. 2, 2, 1 4. 1, 2, 1, 2 5. 2, 2, 1
Types and Evidences that a Chemical Reaction has Occurred
Teacher Tip Ask them to give examples of evidences of chemical changes they have observed around them. Some examples are bleach turning hair yellow, milk going sour, or apple slices becoming brown.
Here are some evidences that a chemical reaction has occurred: a. Change in color b. Formation of a solid (a precipitate) c. Evolution of gas (bubble formation) d. Change in temperature (heat is released or absorbed) Most chemical reactions can be classified into five types: 1. Decomposition reaction – a reactant breaks down into two or more products
Chemical reactions can be classified in other ways such as acid-base reactions and oxidationreduction reactions. However, these concepts will be introduced in later chapters.
AB ! A + B Li2CO3 ! Li2O + CO2 2. Synthesis reaction – two or more reactants form a single product A + B ! AB! 2NO + O2 ! 2NO2! 3. Single displacement reaction – one element replaces another in a compound A + BC ! AC + B Cu(s) + 2AgNO3(aq) ! Cu(NO3)2(aq) + 2Ag(s) 4. Double displacement – two ionic compounds exchange ions AB + CD ! AD + CB 2KI(aq) + Pb(NO3)2(aq) ! 2KNO3(aq) + PbI2(s) 5. Combustion reaction – a hydrocarbon (a compound containing carbon and hydrogen) reacts with oxygen to form carbon dioxide and water. Hydrocarbon + O2 ! CO2 + H2O 2C2H6 + 7O2 ! 4CO2 + 6H2O
Teacher Tip H2O2 breaks down into H2O and O2 aided by light. This is a decomposition reaction. Write and balance the equation. See Practice Exercise 3 on balancing equations above.
ENRICHMENT (2 minutes)
1. Why do you need to store hydrogen peroxide away from light often in dark colored bottles? 2. Learners may watch the videos in the sites given in the Resources section above. These can be given as assignments. 88
Note The Evaluation will be through the exercises in the laboratory session hour.
General Chemistry 1
120 MINS
Lesson 12: Chemical Reactions and Chemical Equations (Laboratory) Content Standard The learners demonstrate an understanding of the use of chemical formulas to represent chemical reactions. Performance Standards The learners shall be able to design, using multimedia, demonstrations, or models, a representation or simulation of any of the following: 1. Atomic structure 2. Mass relationships in reactions Learning Competencies At the end of the lesson, the learners: 1. Write equations for chemical reactions and balance the equations (STEM_GC11CR-If-g-34); 2. Interpret the meaning of a balanced chemical reaction in terms of the Law of Conservations of Mass (STEM_GC11CR-If-g-35); 3. Describe evidences that a chemical reaction has occurred (STEM_GC11CRIf-g-36); and 4. Perform exercises on writing and balancing chemical equations (STEM_GC11CR-If-g-37). Specific Learning Outcomes At the end of the lesson, the learners shall be able to: 1. Write and balance chemical equations; 2. Derive pertinent information from a balanced chemical equation; 3. Determine whether a chemical reaction has occurred or not; and 4. Classify chemical reactions. 89
Lesson Outline Practice
Laboratory Exercises
120
Resources (1) Burdge, J & Overby, J. (2012). Chemistry: Atoms first. New York: McGraw-Hill. (2) Chang, R. & Goldsby, K. (2016). Chemistry. (12th ed.). New York: McGraw-Hill. (3) Chemical equations [Online lecture]. Retrieved from CK-12 website: https://www.ck12.org/physical-science/Chemical-Equations-inPhysical-Science/ (4) Moore, J.W., Stanitski, C.L.& Jurs, P.C. (2012). Chemistry: The molecular science (4th ed.). Belmont, CA: Brooks Cole/Cengage Learning. (5) Recognizing chemical reactions [Online lecture]. Retrieved from CK-12 website: https://www.ck12.org/physical-science/RecognizingChemical-Reactions-in-Physical-Science/ (6) Zumdahl, SS. & Zumdahl, S.A. (2012). Chemistry: An atoms first approach. Belmont, CA: Brooks/Cole Cengage Learning.
PRACTICE
Laboratory exercises (120 minutes) Give the following exercises in a separate time slot like the laboratory session. This provides practice for the learners. Allow them 60 minutes to answer the exercises. Then go over the exercises together. Ask them to show their answers on the board and explain their answers. Alternatively, part of the exercises can be taken as a quiz or as an assignment for evaluation.
Answer Key 1. B. 2. B. 3. D. 4. B. 5. D. 6. E. 7. C. 8. A. 9. A. 10. E.
11. A. 12. D. 13. A. 14. B. 15. A. 16. D. 17. C. 18. B. 19. E. 20. B.
LABORATORY ACTIVITY: CHEMICAL REACTIONS AND CHEMICAL EQUATIONS Directions: Choose the best answer. Encircle the letter corresponding to your answer.
1. Balanced chemical equations imply which of the following? A. Numbers of molecules are conserved in chemical change. B. Numbers of atoms are conserved in chemical change. C. Volume is conserved in chemical change. D. A and B E. B and C
3. The catalytic conversion of ammonia to nitric oxide is the first step in a three-step process, which ultimately results in nitric acid. Balance the equation for the reaction. a NH3(g) + b O2(g) ! c NO(g) + d H2O(g) A. B. C. D. E.
2. In balancing an equation, we change the __________ to make the number of atoms on each side of the equation balance. A. formulas of compounds in the reactants B. coefficients of reactants and products C. formulas of compounds in the products D. subscripts of compounds E. the reactants
90
a = 2, b = 1, c = 2, d = 1 a = 3, b = 2, c = 3, d = 3 a = 4, b = 3, c = 2, d = 6 a = 4, b = 5, c = 4, d = 6 a = 6, b = 15, c = 6, d = 9
4. In the reaction: a BaCl2 + b AgNO3 ! c Ba(NO3)2 + d AgCl
7. Which of the following equations is not balanced?
What is the coefficient, d, of silver chloride in the balanced equation? A. 1 B. 2 C. 3 D. 4 E. 5
A. 4Al + 3O2 ! 2Al2O3 B. C2H6 + O2 ! 2CO2 + 3H2O C. 2KClO3 ! 2KCl + O2 D. 4P4 + 5S8 ! 4P4S10 E. P4 + 5O2 ! P4O10 8. The first step in the Ostwald process for making nitric acid is the formation of NO as follows:
5. Balance the following equation with the smallest set of whole numbers. C4H10 + O2 ! CO2 + H2O
4NH3 + 5O2 ! 4NO + 6H2O
What is the coefficient for CO2 in the balanced equation? A. 1 B. 4 C. 6 D. 8 E. 12
According to the equation, 5 moles NH3 will react with ________ moles O2 to form _______ moles of NO. A. 5, 4 B. 4, 5 C. 25, 20 D. 5/4, 4/5
6. Balance the following equations:
E. 25/4, 5
_____ P4O10 + _______ H2O ! _______ H3PO4
9. Potassium metal and chlorine gas (Cl2) react in a combination reaction to produce potassium chloride. What is the correct balanced equation for this reaction?
What is the coefficient of H2O in the balanced equation? A. B. C. D. E.
1 2 4 5 6
A. 2 K(s) + Cl2(g) !! 2 KCl(s) B. K(s) + Cl2(g) ! KCl(s) C. K(s) + Cl(g) !! KCl(s) D. K2(s) + Cl2(g) ! 2 KCl(s) E. K(s) + Cl2(g) !! KCl2(s)
91
10. In the reaction given below, how many grams of water are consumed if 4.0 g hydrogen gas and 32.0 g oxygen gas are produced?
13. Balance the following equation: a NaNO3 ! b NaNO2 + c O2 What are the coefficients of the balanced equation for
2 H2O ! 2 H2 + O2 A. B. C. D. E.
a, b, and c?
2.0 g 4.0 g 18.0 g 20.0 g 36.0 g
A. B. C. D. E.
11. In the reaction given below, for every two molecules of hydrogen peroxide (H2O2) consumed, how many molecules of oxygen are produced?
14. Balance the following chemical reaction: a CO + b NO ! c CO2 + d N2 The coefficients a, b, c, and d for the balanced chemical equation are:
2H2O2 ! 2H2O + O2 A. B. C. D. E.
2, 2, 1 1, 1, 2 1, 2, 1 2, 3, 1 3, 1, 1
1 2 3 6 9
A. 2, 2, 2, 3 B. 2, 2, 2, 1 C. 1, 1, 1, 2 D. 2, 1, 2, 1 E. 1, 2, 2, 1
12. Balance the following reaction: 15. Classify the following reaction:
a Al2O3 ! b Al + c O2
2Na + Cl2 ! 2 NaCl
What is the sum of the coefficients of the reactant and products (a + b + c) in the balanced equation using the smallest set of whole numbers as coefficients? A. B. C. D. E.
A. Synthesis B. Decomposition
3 5 6 9 10
C. Combustion D. Single Displacement E. Double Displacement
92
16. Classify the following reaction:
19. Classify the following reaction:
Zn + 2HCl ! ZnCl2 + H2
NaCl(aq) + AgF(aq) ! NaF(aq) + AgCl(s)
A. Synthesis
A. Synthesis
B. Decomposition
B. Decomposition
C. Combustion
C. Combustion
D. Single Displacement
D. Single Displacement
E. Double Displacement
E. Double Displacement
17. Classify the following reaction: H2SO3 ! H2O + SO2
20. Classify the following reaction: CaCO3(s) ! CaO(s) + CO2(g)
A. Synthesis B. Decomposition
A. Synthesis
C. Combustion
B. Decomposition
D. Single Displacement
C. Combustion
E. Double Displacement
D. Single Displacement E. Double Displacement
18. Classify the following reaction: CH4 + 2O2 ! CO2 + 2H2O A. Synthesis B. Decomposition C. Combustion D. Single Displacement E. Double Displacement
93
General Chemistry 1
Lesson 13: Mass Relationships in Chemical Reactions Content Standard The learners demonstrate an understanding of quantitative relationship of reactants and products in a chemical reaction.
180 MINS
Lesson Outline
Introduction
Communicating Learning Objectives
5
Motivation
Let us Make Sandwiches
5
Performance Standard The learners shall be able to design, using multimedia, demonstrations, or models, a representation or simulation of any of the following: 1. Atomic structure 2. Mass relationships in reactions
Instruction, Delivery and Practice
I. Reactants and Products II. Limiting Reagents III. Theoretical Yield, Actual Yield, and Percent Yield
140
Enrichment
Animation Videos of Limiting Reagents
15
Learning Competencies At the end of the lesson, the learners: 1. Construct mole or mass ratios for a reaction in order to calculate the amount of reactant needed or amount of product formed in terms of moles or mass (STEM_GC11MR-Ig-h-38); 2. Calculate percent yield and theoretical yield of the reaction (STEM_GC11MR-Ig-h-39); 3. Explain the concept of limiting reagent in a chemical reaction; identify the excess reagent (STEM_GC11MR-Ig-h-40); and 4. Calculate reaction yield when a limiting reagent is present (STEM_GC11MR-Ig-h-41).
Evaluation
Check up Quiz
15
Specific Learning Outcomes At the end of the lesson, the learners shall be able to: 1. Identify mole ratios of reactants and products from balanced chemical equations; 2. Perform stoichiometric calculations related to chemical equations; 3. Define theoretical, actual, and percent yield of reactions; 4. Calculate theoretical and percent yield of a reaction; 5. Identify the limiting and excess reagent(s) of a reaction; and 6. Calculate reaction yield in the presence of a limiting reagent. 94
Materials
Periodic table, calculator Resources (1) Allan, Andy. Stoichiometry [PowerPoint presentation]. Retrieved from http://www.sciencegeek.net/APchemistry/FlashPPT/3_Stoichiometry/ index.html (2) Burdge, J. & Overby, J. (2012). Chemistry: Atoms first. New York: McGraw-Hill (3) Chang, R. &Goldsby, K. (2016). Chemistry. (12th ed.). New York: McGraw-Hill (4) Limiting reagent [Vector animation]. Retrieved from McGraw Hill Education web site: http://www.mhhe.com/physsci/chemistry/ essentialchemistry/flash/limitr15.swf (5) Limiting reactant [Vector animation]. Retrieved from North Carolina School of Medicine and Mathematics web site: http:// www.dlt.ncssm.edu/core/Chapter6-Stoichiometry/Chapter6Animations/LimitingReactant.html (6) Moore, J.W., Stanitski, C.L. & Jurs, P.C. (2012). Chemistry: The molecular science (4th ed.). Belmont, CA: Brooks Cole/Cengage Learning. (7) Reactants, products and leftovers [Simulation]. Retrieved from PhEt Interactive Simulations web site: http://phet.colorado.edu/sims/html/ reactants-products-and-leftovers/latest/reactants-products-andleftovers_en.html (8) Zumdahl, SS. &Zumdahl, S.A. (2012). Chemistry: An atoms first approach. Belmont, CA: Brooks/Cole Cengage Learning.
INTRODUCTION (5 minutes)
1. Introduce the learning objectives using any of the suggested protocol (Verbatim, Own Words, or Read-aloud): a. Identify mole ratios of reactants and products from balanced chemical equations b. Perform stoichiometric calculations related to chemical equations c. Define theoretical, actual, and percent yield of reactions d. Calculate theoretical and percent yield of a reaction e. Identify the limiting and excess reagent(s) of a reaction f. Calculate reaction yield in the presence of a limiting reagent
Teacher Tip List these keywords on the board or through PowerPoint slides. Alternatively, you can write them on flip charts.
2. Present the keywords for the concepts to be learned: a. Stoichiometry b. Limiting reagent c. Theoretical yield d. Actual yield e. Percent yield 3. Review the Law of Conservation of Mass
MOTIVATION (5 minutes)
Teacher Tip Ask the learners to show the equation:
Let us make sandwiches!
Some learners are going on a road trip and they are to bring some food to eat along the way. Karen was asked to bring hamburger sandwiches for which she will use two slices of bread and one hamburger patty to make one sandwich. Show the equation.
INSTRUCTION/DELIVERY/PRACTICE(40 minutes) Amounts of Reactants and Products
Stoichiometry is the study of the quantities of materials consumed and produced in chemical reactions. From the balanced chemical equation, we will be able to: a. Determine how much products will be produced from a specific amount of reactants b. Determine the amount of reactants needed to produce a specific amount of products
95
two slices of bread + one hamburger patty " one hamburger sandwich This analogy will be used for mass relationships in chemical equations.
1. Illustrate stoichiometry using the following examples: a. Let us make hamburger sandwiches again. The equation is
Teacher Tip Illustrate stoichiometry with real life applications.
two$slices$of$bread$+$one$hamburger$patty$$!$$one$hamburger$sandwich$ Suppose Karen has 14 hamburger patties, how many slices of bread will she need to consume all the patties? The ratio of slices of bread to hamburger patty is 2: 1.
2. Suppose that instead of plain burgers, Karen is to make double cheeseburgers. Show the equation so Karen can shop for enough ingredients. two$slices$of$bread$$+$one$hamburger$patty$$+$$two$slices$of$cheese$$$!$$$one$double$cheeseburger$ How many slices of cheese, hamburger patties, and slices of bread will Karen need to make 25 double cheeseburgers?
Karen will therefore have to buy 50 slices of cheese, 25 hamburger patties, and 50 slices of bread.
96
3. Ammonia, NH3, is a leading industrial chemical used in the production of agricultural fertilizers and synthetic fibers. It is produced by the reaction of nitrogen and hydrogen gases: 3 H2(g) + N2(g) ! 2 NH3(g) The balanced equation says that 3 moles H2 are stoichiometrically equivalent to 1 mole N2 and to 2 moles NH3. The ratio of moles H2 to moles NH3 is 3:2; the ratio of moles N2 to moles NH3 is 1:2. a. How many moles of NH3 will be produced if 10.4 moles H2 react completely with N2? (moles H2 ! moles NH3)
b. How many moles of N2 are needed to produce 42.4 moles NH3? (moles NH3 ! moles N2)
c. How many grams of NH3 will be produced from 25.7 moles N2 (moles N2 ! moles NH3 ! g NH3)
d. How many grams of NH3 will be produced if 122 g N2 reacts completely with H2? (g N2 ! moles N2 ! moles NH3 ! g NH3)
97
Teacher Tip Before doing any calculations involving chemical reactions, make sure that the chemical equation is balanced. Recall the interpretation of a balanced chemical equation. Teacher Tip Give these examples but ask them to supply the appropriate ratios or factors. Before calculating, ask them the steps they will be taking to get the answer.
4. Solid lithium hydroxide is used to remove carbon dioxide and is called a CO2 scrubber. This technique has been used for space vehicles. The reaction is: 2 LiOH(s) + CO2(g) ! Li2CO3 (s) + H2O (l) How many grams of CO2 can be absorbed by 785.0 g LiOH? What are the steps of the solution? a. Convert grams LiOH to moles LiOH. b. Get the moles of CO2 stoichiometrically equivalent to moles LiOH c. Convert moles CO2 to grams CO2. (grams LiOH ! moles LiOH ! moles CO2 ! grams CO2)
Ask them to answer the following practice exercises: 1. The combustion of carbon monoxide gas in oxygen gas is represented by the following balanced equation: 2 CO(g) + O2(g) ! 2CO2(g) How many moles of carbon dioxide gas will be produced from the complete combustion of 4.60 moles CO(g)? 2. Consider the reaction: 2 KClO3 ! 2 KCl + 3 O2 a. How many moles of KClO3 are required to produce 22.8 moles oxygen gas, O2? b. How many moles of KCl will be produced from the total decomposition of 18.8 moles KClO3?
98
Teacher Tip Assign the exercises to different groups. Ask them to show the calculations on the board. If there is no longer enough time, this could be given as an assignment. Alternatively, this can also be used as a quiz to check on their understanding of the concept. Answer Key 1. 4.60 moles CO2 2.a. 15.2 moles 2.b. 18.8 moles 3.a. 388 g Fe2O3 3.b. 6.43 x 103 g 4.a. 156.0 g MgO 4.b. 135.0 g MG
3. Given the reaction 4 Fe + 3 O2 ! 2 Fe2O3 a. How many grams of Fe2O3 will be formed from 4.86 moles Fe reacting with sufficient oxygen gas? b. How many grams of Fe are needed to react with sufficient oxygen to produce 28.8 moles Fe2O3? 4. Consider the reaction 2Mg + O2 ! 2MgO a. How many grams of MgO are produced from the complete reaction of 94.2 g Mg? b. How many grams of Mg are needed to produce 224 g of MgO in the complete reaction of Mg with oxygen gas? Limiting Reagents The reactant used up first in the chemical reaction is called the limiting reagent. Excess reagents are reactants present in quantities greater than what is needed by the reaction. Illustrate using the following examples: 1. Recall the example of the double cheeseburger. The equation is: two$slices$of$bread$$+$$one$hamburger$patty$$+$$two$slices$of$cheese$$!$$$one$double$cheeseburger$ When Karen went shopping, she was able to buy 50 slices of cheese, 20 hamburger patties, and 50 slices of bread. How many double cheeseburgers can she make? What is the limiting material or reagent? What are the excess reagents? To find the limiting reagent, determine which reagent will give the smallest amount of product.
99
Therefore, the limiting reagent is the hamburger patty. 40$slices$of$bread$ 10$slices$in$excess
+
20$patties +
40$slices$of$cheese$ 10$slices$in$excess
!
20$double$cheeseburger
Karen can only make 20 double cheeseburgers. The limiting reagent is the hamburger patty. There are ten slices of bread and ten cheese slices in excess. Karen cannot make more than 20 sandwiches because all the hamburger patties have been used up. 2. Consider again the reaction: 3H2(g) + N2(g) ! 2NH3(g) a. If 6.60 moles H2 are made to react with 4.42 moles N2, what is the limiting reagent? How many moles NH3 will be produced? What reagent is in excess and by how much? Determine which reagent will produce the smallest amount of product:
Therefore, the limiting reagent is H2. 100
The amount of limiting reagent present at the start of the reaction determines the theoretical yield. To determine the amount of NH3 produced, use the limiting reagent.
The excess reagent is N2. If you have 6.60 moles H2 then you will need
But you have 4.42 moles N2. Therefore, the excess amount of N2 is 4.42 moles – 2.20 moles = 2.22 moles N2. b. If 25.5 g H2 are made to react with 64.2 g N2, what is the limiting reagent? What is the theoretical yield in g of NH3 that will be produced? How do you determine the limiting reagent? i. Get the number of moles of each reactant. ii. Calculate the number of moles of product using each reagent. iii. The one that yields the smallest number of moles of product is the limiting reagent.
From 12.6 moles of H2, how many moles of NH3 do we expect to get?
101
Teacher Tip This example shows that even though the mass of N2 was greater than the mass of H2, the limiting reagent was still N2. This illustrates that the limiting reagent is not determined by which reactant is present in greater amount. It is only by considering the mole ratios and relationships in the balanced chemical reaction that the limiting reagent can be determined.
From 2.29 moles of N2, how many moles of NH3 do we expect to get?
The limiting reagent is N2. What amount of NH3 will be formed in this example? The amount of product that can be produced is determined by the limiting reagent. Once the limiting reagent is consumed, there is no further reaction. Hence, to calculate the amount of NH3 produced, we use 2.29 moles N2, the limiting reagent. Theoretical Yield, Actual Yield, and Percent Yield The theoretical yield is the maximum amount of product that would result if the limiting reagent is completely consumed. It is the amount of product predicted by stoichiometry (as shown in the above example). The actual yield is the quantity of the desired product actually formed.
If in the example given above, only 54.0 g NH3 were produced, then the actual yield is 54.0 g; the theoretical yield is 78.0 g and the % yield is:
102
Teacher Tip Explain why the theoretical yield is not obtained in actual work. Ask them for possible reasons.
Ask them to answer this practice exercise: 1. Silver metal reacts with sulfur to form silver sulfide according to the following reaction: 2Ag (s) + S(s) ! Ag2S (s)
Answer Key 1. Ag 2. 57.5 g 3. 2.57 g 4. 78.3 %
a. Identify the limiting reagent if 50.0 g Ag reacts with 10.0 g S. b. What is the theoretical yield in g of Ag2S produced from the reaction? c. What is the amount in g of the excess reactant expected to remain after the reaction? d. When the reaction occurred, the amount of Ag2S obtained was 45.0 g. What is the percent yield of the reaction?
ENRICHMENT (15 minutes)
Watch the animation videos of limiting reagent from the following sources: •
Limiting reagent [Vector animation]. Retrieved from McGraw Hill Education web site:
Teacher Tip If this cannot be shown in the classroom, learners can be asked to view the animation at home or in the library.
http://www.mhhe.com/physsci/chemistry/essentialchemistry/flash/limitr15.swf •
Limiting reactant [Vector animation]. Retrieved from North Carolina School of Medicine and Mathematics web site: http://www.dlt.ncssm.edu/core/Chapter6-Stoichiometry/Chapter6-Animations/ LimitingReactant.html
•
Reactants, products and leftovers [Simulation]. Retrieved from PhEt Interactive Simulations web site: http://phet.colorado.edu/sims/html/reactants-products-and-leftovers/latest/ reactants-products-and-leftovers_en.html
EVALUATION (15 minutes)
Encircle the letter of the best answer.
2. Given the reaction CH4 + 2O2 ! CO2 + 2H2O, what amount of O2 is needed to completely react with 14.0 moles CH4? A. 2.0 moles B. 28.0 moles C. 12.0 moles D. 6.0 moles E. 1.0 mole
1. Stoichiometry deals with A. Combustion reactions B. Rates of chemical reactions C. Heat evolved or absorbed during chemical reactions D. The study of amounts of materials consumed and products formed in chemical reactions E. Activation energy of chemical reactions 103
3. How much of SnF2 (stannous fluoride, active ingredient in toothpaste) in g can be prepared from the reaction of 10.0 g SnO with excess HF according to the following reaction? SnO + 2HF ! SnF2 + H2O A. 11.6 g B. 10.0 g C. 9.62 g D. 26.0 g E. 104.0 g
6. The reaction N2(g) + 2O2(g) ! N2O4(g) occurs in a closed container. If 8.0 moles N2(g) are made to react with 12.0 moles O2, the limiting reagent and the theoretical yield of N2O4 are: A. The limiting reagent is N2; the theoretical yield of N2O4 is 8.0 moles B. The limiting reagent is N2; the theoretical yield of N2O4 is 16.0 moles C. The limiting reagent is O2; the theoretical yield of N2O4 is 12.0 moles D. The limiting reagent is O2; the theoretical yield of N2O4 is 6.0 moles E. The limiting reagent is O2; the theoretical yield of N2O4 is 8.0 moles
4. What is a limiting reagent? A. The reactant that is used up last and prevents more product from being made B. The reactant that is never used up C. The reactant that is used up first and prevents more products from being made D. The reactant that is in excess and does not get used up in the reaction E. The reactant that is always in greater quantity
7. The reaction of 5.0 g hydrogen with 5.0 g carbon monoxide produced 4.5 g methanol. What is the percent yield for the reaction 2H2 + CO ! CH3OH? A. 11% B. 79% C. 96% D. 24% E. 63%
5. A mixture of 2.0 moles I2 and 4.0 moles Zn are reacted to completion in a closed container according to the following chemical equation: I2 + Zn ! ZnI2. What are the contents of the container after the reaction? A. Zn and ZnI2 B. I2 and ZnI2 C. Zn and I2 D. I2, Zn, and ZnI2 E. ZnI2
8. The reaction of 5.0 g fluorine with excess chlorine produced 5.6 g ClF3 in the reaction Cl2 + 3F2 ! 2ClF3. What was the percent yield of the reaction? A. 58% B. 69% C. 76% D. 86% E. 92%
104
General Chemistry 1
120 MINS
Lesson 14: Mass Relationships in Chemical Reactions (Laboratory) Content Standard The learners demonstrate an understanding of the quantitative relationship of reactants and products in a chemical reaction. Performance Standards The learners shall be able to design, using multimedia, demonstrations, or models, a representation or simulation of any of the following: 1. Atomic structure 2. Mass relationships in reactions Learning Competencies The learners determine the mass relationship in a chemical reaction (STEM_GC11MR-Ig-h-42). Specific Learning Outcomes At the end of the lesson the learner will be able to: 1. Prepare NaCl from the reaction of sodium bicarbonate and hydrochloric acid; 2. Determine the actual yield of the reaction; 3. Illustrate the mass relationship in a chemical reaction by calculating the theoretical yield of the reaction; and 4. Determine the percentage yield of the reaction.
105
Lesson Outline Motivation
Why is Baking Soda Added to Cakes and Cookies?
Introduction
Introduction to Laboratory Activity
17
Instruction, Delivery and Practice
Laboratory Activity
80
Enrichment
Post-laboratory Session
20
Evaluation
Data Sheet and Activity Sheet
3
Materials Evaporating dish, watch glass, balance (triple beam or electronic balance), Sodium bicarbonate, spatula or small plastic knife, dilute hydrochloric acid (3 moles), beaker or glass container for the acid, long dropper, Bunsen burner, wire gauze or mesh, iron stand, iron ring, wash bottle, and distilled water Resources (1) Burdge, J. & Overby, J. (2012). Chemistry: Atoms first. New York: McGraw-Hill. (2) Chang, R. & Goldsby, K. (2016). Chemistry. (12th ed.). New York: McGraw-Hill. (3) Moore, J.W., Stanitski, C.L. & Jurs, P.C. (2012). Chemistry: The molecular science (4th ed.). Belmont, CA: Brooks/Cole Cengage Learning. (4) Zumdahl, SS. &Zumdahl, S.A. (2012). Chemistry: An atoms first approach. Belmont, CA: Brooks/Cole Cengage Learning.
MOTIVATION (3 minutes)
Why do we add baking soda, NaHCO3, in baking cookies and cakes?
INTRODUCTION (17 minutes)
Introduce the laboratory experiment as indicated in the Laboratory Sheet.
INSTRUCTION: Laboratory activity (80 minutes) Safety Precautions: 1. Never taste anything during a science activity. 2. Wear appropriate laboratory attire; goggles and apron must be worn throughout the experiment. 3. Dispose of the materials as directed by your teacher. 4. Wash your hands with soap and water after the science activity.
Teacher Tip Baking soda is used to make cakes and cookies ‘rise’. When a weak acid such as lemon juice, vinegar, or buttermilk is added to baking soda, bubbles of carbon dioxide are produced. The release of gas is what causes the cake to ‘rise’. Teacher Tip 1. Prepare the classroom or laboratory, the materials, the laboratory sheets to be used. 2. Distribute the laboratory sheets at the start of the lesson. 3. After the introduction and motivation, explain the procedures of the activity. 4. Explain the safety precautions. 5. This activity can be performed individually or in groups. Teacher Tip It is important to discuss the safety precautions thoroughly before starting the experiment.
5. Follow all laboratory instructions as directed by your teacher. Procedure: 1. Discuss the procedure of the experiment. Demonstrate the setup to be used. 2. During the activity, they should record the data in the activity sheets. 3. After the activity, make sure that the learners clean the materials and equipment and properly dispose of the product.
106
Note the precautions in handling the acid and in lighting the burner. The product, NaCl, can be disposed in the sink during washing.
ENRICHMENT (20 minutes) POST-LABORATORY SESSION Give them enough time to accomplish the activity sheet of the experiment. EVALUATION EXCEEDS EXPECTATIONS The learner performed the experiment using proper laboratory techniques while observing safety precautions; and was able to answer at least 75% of the calculations and discussions in the activity sheet.
MEETS EXPECTATIONS
NEEDS IMPROVEMENT
NOT VISIBLE
The learner needed to improve his/ her use of laboratory techniques; but was able to observe safety precautions; and was able to answer at least 60% of the calculations and discussions in the activity sheet.
The learner needed to improve his/ her use of laboratory techniques and observance of safety measures; and was able to answer at least 50% of the calculations and discussions in the activity sheet.
The learner did not observe proper safety procedures for the experiment; did not use proper laboratory techniques; and was able to answer less than 25% of the items in the activity sheet.
LABORATORY EXPERIMENT MASS RELATIONSHIPS IN CHEMICAL REACTIONS Introduction
A reaction is said to have been completed if one of the reactants is completely consumed by the reaction. In this experiment, sodium bicarbonate (baking soda) is made to react with hydrochloric acid to produce sodium chloride according to the reaction: NaHCO3 +HCl → NaCl + H2O + CO2(g) You will use an accurately measured amount of NaHCO3 and add enough HCl until the bicarbonate is completely used up. You will isolate the product, NaCl, from the other products and determine its mass. This is the actual yield of the reaction. The theoretical yield can be calculated by using the mass relationships in the balanced chemical equation above. The percentage yield can be determined from the ratio of the actual yield to the theoretical yield. Objectives 1. To perform a chemical reaction and measure the actual yield of sodium chloride from the chemical reaction. 2. To determine the percent yield of the reaction.
107
Materials
a. Evaporating dish b. c. d. e. f. g.
h. i. j. k. l. m. n.
Long dropper Watch glass Bunsen burner Balance (triple beam or electronic balance) Wire gauze or mesh Sodium bicarbonate
Iron stand Spatula or small plastic knife Iron ring Dilute hydrochloric acid (3 moles) Wash bottle Beaker or glass container for the acid Distilled water (use commercially available distilled water)
Safety Precautions 1. Never taste anything during a science activity. 2. 3. 4. 5.
Wear appropriate laboratory attire; goggles and apron must be worn throughout the experiment. Dispose of the materials as directed by your teacher. Wash your hands with soap and water after the science activity. Follow all laboratory instructions as directed by your teacher.
Procedure: 1. 2. 3. 4.
Clean and dry an evaporating dish and a watch glass. The watch glass will be used as the cover of the evaporating dish. Weigh the combination of the evaporating dish and the watch glass to the nearest 0.01 g. Record the mass in the data table.
Put about 2.00 grams of pure NaHCO3 into the dish. Weigh the dish, the contents, and the cover watch glass to the nearest 0.01 g. Set up the Bunsen burner, ring, and wire mesh, and set the evaporating dish on the wire mesh. To cover the dish, place the curved side down and the glass slightly off center so that the lip of the dish is uncovered. Do not light the burner yet. 5. Add dilute hydrochloric acid drop wise down the lip of the dish to the bicarbonate sample in the dish. Continue adding the acid dropwise until no more reaction takes place when a drop of acid is added. Gently swirl the contents of the dish so that all of the solid gets in contact with the liquid. Do not add excess acid. Caution: HCl acid is caustic and corrosive. Avoid contact with skin and eyes. Avoid breathing the vapors. Wear safety goggles and apron. Wipe away all spills. If any acid spills on you, immediately flush the area with water and notify your teacher.
108
6. Carefully rinse the bottom of the watch glass with distilled water, a few drops at a time, and collect all the washings in the evaporating dish. 7. Gently heat the evaporating dish, contents, and cover with a low flame until the salt is completely dry. Move the burner back and forth to avoid spattering. If the contents of the dish spatter, reduce the flame. 8. Turn off the flame. Allow the dish to cool to room temperature. Weigh the dish, contents, and cover watch glass to the nearest 0.01 g. Caution: Before you light the burner, make sure that long hair and loose clothing have been confined. Remember to allow all apparatus to cool before you handle it again. 9. Repeat Steps 7 and 8 to be sure that constant weight has been obtained. Two consecutive mass readings should agree within 0.02 g. Calculations: 1. Calculate the theoretical yield of NaCl that should have been obtained from the reaction. Show your calculations. 2. Determine the percentage yield. Discussion: 1. What is the cause of the effervescence that you observed during the reaction? 2. How can you conclude that the reaction has gone to completion? 3. How do you know if the product was completely dry? 4. List possible sources of error which you think affected the yield of your reaction. Did your error cause your result to be higher or lower than the theoretical yield? Additional Exercises - Practice: 1. Suppose you started with 6.4 g NaHCO3 and added sufficient HCl for a complete reaction, how much NaCl in g would you expect to produce? Show your calculations. 2. If you wanted to produce 3.78 g NaCl, how much NaHCO3 in g would you start with, assuming no loss of product occurs? REPORT SHEET: MASS RELATIONSHIPS - LABORATORY DATA TABLE Mass of evaporating dish, watch glass, and NaHCO3 Mass of evaporating dish and watch glass Mass of NaHCO3 Mass of evaporating dish, cover, and NaCl
Trial 1 ________________ Trial 2 ________________ Trial 3 _________________(if needed)
Mass of NaCl obtained (experimental)
109
General Chemistry 1
240 MINS
Lesson 15: Gases (Lecture) Content Standard The learners demonstrate an understanding of:
Lesson Outline
1. The mathematical relationship between the pressure, volume, and temperature of a gas;
Introduction
Presentation of Learning Objectives and Important Keywords
5
2. The partial pressures of a gas; 3. The quantitative relationships of the reactants and products in a gaseous reaction; and 4. Behavior and properties of gases at the molecular level.
Motivation
Relate the Presence of Gases in Daily Life
5
Instruction, Delivery and Practice
I. II. III. IV.
Enrichment
Problem Solving
Evaluation
Check-up Quiz
Performance Standards The learners shall be able to: 1. Determine the volume or pressure of gas at different conditions; 2. Determine the pressure of the components or of the whole gas mixture; 3. Calculate the amount of products or reactants involved in a gaseous reaction; and 4. Explain the properties and behavior of a gas in terms of its molecular composition. Learning Competency At the end of the lesson, the learners:
1. Define pressure and give the common units of pressure (STEM_GC11G-Ih-i-43); 2. Express the gas laws in equation form (STEM_GC11G-Ih-i-44); 3. Use the gas laws to determine pressure, volume, or temperature of a gas under certain conditions of change (STEM_GC11G-Ihi-45); 4. Use the Ideal Gas Equation to calculate pressure, volume, temperature, or number of moles of a gas (STEM_GC11G-Ihi-46);
Gas Laws Gas Mixtures Reaction Stoichiometry Involving Gases Kinetic Molecular Theory of Gases
200
30
Materials Projector, computer, flip charts Resources (1) Brown, T. L., Bursten, B. E., LeMay Jr., H. E., Murphy, C., & Woodward, P. (2014). Chemistry: The central science. (13th ed.). Upper Saddle River, NJ: Prentice Hall. (2) Burdge, J. (2013).Chemistry. (3rd ed). New York: McGraw-Hill. (3) Chang, R. & Goldsby, K. (2016). Chemistry. (12th ed.). New York: McGraw-Hill.
6. Apply the principles of stoichiometry to determine the amounts (volume, number of moles, or mass) of gaseous reactants and products (STEM_GC11GS-Ii-j-48); 7. Explain the gas laws in terms of the kinetic molecular theory of gases (STEM_GC11KMT-Ij-49); and 8. Relate the rate of gas effusion with molar mass (STEM_GC11KMT-Ij-50).
5. Use Dalton’s Law of Partial Pressures to relate mole fraction and partial pressure of gases in a mixture (STEM_GC11DL-Ii-47); 110
INTRODUCTION (5 minutes)
1. Introduce the learning objectives using the suggested protocol (Read-aloud): a. I will be able to calculate the pressure or volume of a gas at different conditions
Teacher Tip Display the objectives prominently on the board, so that the learners can track the progress of their learning. List these keywords on the board.
b. I will be able to determine the pressure of a gas mixture or of its components c. I will be able to determine the amount of products or reactants involved in a gas phase reaction d. I will be able to discuss the properties and behavior of gases in terms of its molecular composition 2. Present the keywords for the concepts to be learned: a. Boyle’s Law b. Charles’s Law c. Avogadro’s Law d. Ideal Gas Equation e. Partial pressure f.
Dalton’s Law
g. Gas reaction stoichiometry h. Kinetic Molecular Theory
MOTIVATION (5 minutes)
Point out the abundance of gases in their surroundings, such as in the environment, at home, and in other places. Ask them where gases are encountered or used in everyday life. Some expected responses are: a. In the air, which supplies us with the gases we breathe b. In the kitchen, wherein a gas (liquid petroleum gas) is used for heating or cooking c. In the hospital, wherein gases are used to aid the breathing of patients d. In the automobile, wherein gases are burned in order to make the cars move e. In carbonated drinks, wherein a gas (carbon dioxide) makes the drinks refreshing 111
Teacher Tip The lesson is essentially a review of the basic concepts presented and used in junior high school.
INSTRUCTION / DELIVERY / PRACTICE (200 minutes)
Note: The delivery will be done in four 50-minute sessions. It is best to summarize the concepts learned at the end of each session. PART ONE Ask the learners to recall the definition of pressure – the amount of force exerted per unit area. Let them give the unit for pressure. Show them a balloon and ask them to point out the role of the pressure of the gas inside the balloon. Ask them to imagine the tire of a vehicle and the need to pump air into the tire up to a given pressure. a. What will happen if the pressure is much lower than what it should be? b. What will happen if the pressure is much greater than what it should be? Ask them if they know how the air pressure of the tire is measured and expressed. Point out the various units used for pressure: a. The old air pump in the gasoline stations used the unit pounds per square inch (psi), which is widely used especially in the United States, but usage of which is supposedly discouraged. b. Later on, the new air pumps used kilopascal (KPa) (or newton per square meter, N/m2), which is the SI unit for pressure. c. In chemistry, a widely used unit for pressure is the atmosphere (atm), but the International Union of Pure and Applied Chemistry discourages its usage. However, it takes some time for usage of this unit to be discontinued. d. Another old and popular unit for pressure is the Torr (or mmHg), yet the International Union of Pure and Applied Chemistry is also discouraging the usage of this unit.
112
Teacher Tip Ask them how the balloon would look like if there is no gas inside the balloon. It would be helpful if a picture of an automobile with tires is shown to them.An alternative example is an air mattress. Showing a picture of a gasoline station air pump may also be beneficial to the learners.
Write on the board the relationship between the different units:
Teacher Tip Ask them the origin of the unit Torr and the experiment of Torricelli. If they are not familiar with this, assign them to read about this from the internet and to write a report on what they read.
1 atm = 760 Torr (mm Hg) 1 atm = 101.3 kPa Point out to that aside from pressure, the other parameters (or variables) used to describe gases are volume and temperature. The common unit for volume is the liter (L), but the SI unit for volume is m3. The equivalence of the liter in SI units is simple: 1 L = 1000 m3 1 L = 1 dm3 1 mL = 1 cm3 The common unit for temperature is degree Celsius (oC), but the SI unit is Kelvin (K). The relationship between the units is K = oC + 273 Point out that the relationship between these three parameter are expressed by the Gas Laws: 1. Boyle’s Law 2. Charles’s Law 3. Avogadro’s Law Ask them to state Boyle’s Law and emphasize on expressing the law correctly: ‘The volume of a given amount of gas is inversely proportional to its pressure at constant temperature.’ •
Highlight that Boyle’s Law is valid only if the amount of the gas and the temperature is constant.
•
Write the mathematical expression for Boyle’s Law:
In terms of a proportion: V α 1/ P (at constant amount and temperature) In terms of an equation: PV = k
V = k/P
(at constant amount and temperature)
or
P1 V1 = P2 V2 113
Teacher Tip They would be familiar with Boyle’s Law and Charles’s Law from their Chemistry course in junior high school. A review and an enrichment could be done for these topics. It would be helpful to demonstrate Boyle’s Law through a 60-mL plastic syringe sealed at its inlet. Show what happens to the volume of the trapped gas once pressure of the gas is increased by pushing the plunger inwards.
Draw the graph relating pressure and volume. Point out that the plot is called an isotherm, since the relationship is exhibited only at constant temperature.
Teacher Tip They should be familiar with the expressions, and they could be asked to write them on the board. They might be familiar with the plot from junior high school, and they could be asked to sketch it on the board.
They might have learned how to solve this type of problem from the Chemistry course in junior high school. Let them recall how to solve the problem. Answer key: 1. 3.75 L 2. 3.0 atm
A graph showing the relationship between volume and pressure, as stated by Boyle’s Law Ask them to solve the following problems: 1. A gas sample occupies a volume of 2.5 L at a pressure of 1.5 atm. What would be the volume of the gas if its pressure is reduced to 1 atm at the same temperature? 2. The gas inside a balloon has a volume of 15.0 L at a pressure of 2.0 atm. Calculate the pressure of the gas if its volume is compressed to 10.0 L at the same temperature. Ask them to state Charles’s Law and emphasize on expressing the law correctly: ‘The volume of a given amount of gas is directly proportional to its absolute temperature at constant pressure.’ 1. Highlight that Charles’s Law is valid only if the amount of the gas and the pressure is constant. Also, point out that the temperature should be expressed in the unit Kelvin (K). 114
Teacher Tip They should be familiar with the expressions, and they could be asked to write them on the board.
Teacher Tip They might be familiar with the plot from junior high school, and they could be asked to sketch it on the board.
Write the mathematical expression for Charles’s Law: In terms of a proportion: V α T (at constant amount and pressure) In terms of an equation: V/T=k
V = k T (at constant amount and pressure)
or
V1 / T1 = V2 / T2 Draw the graph relating volume and temperature. Point out that the plot is called an isobar, since the relationship is exhibited only at constant pressure.
A graph showing the relationship between volume and pressure, as stated by Charles’ Law Ask them to solve the following problems: 1. At 30oC, the volume of a sample of air was 5.8 L. What would be the volume of the air sample if it is heated to 60oC at the same pressure? 2. A given amount of oxygen gas has a volume of 25.0 L at a temperature of and a pressure of 1.0 atm. At what temperature would this gas occupy a volume of 22.0 L at a pressure of 1.0 atm? 37oC
115
They might have learned how to solve this type of problem from the Chemistry course in junior high school. Let them recall how to solve the problem. Answer key: 1. 6.37 L 2. 273 K
State Avogadro’s Law: ‘The volume of a gas at a given temperature pressure is directly proportional to the number of moles contained in the volume. •
Mention that this law is based on Avogadro’s hypothesis that ‘the same volume of two gases at the same temperature and pressure contain the same number of molecules’.
•
Let them recall that the SI unit mole is related to the number of molecules in a substance.
•
Point out that experiments have shown that the volume of 1.0 mole of a gas at 0oC and 1 atm is 22.4 L.
•
Write the mathematical expression for Avogadro’s Law:
In terms of a proportion: V α n In terms of an equation: V/n = k
V = kn
(at constant temperature and pressure) (at constant temperature and pressure)
or
V1 / n1 = V2 / n2 Answer Key 1. 168 L 2. 0.446 mol
Ask them to solve the following problems: 1. 1.0 mole of a gas occupies a volume of 22.4 L gas at 0oC and 1 atm. What would be the volume of 7.5 mol of the gas at the same temperature and pressure? 2. The volume of a gas sample at 0oC and 1.0 atm is 10.0 L. How many moles of gas are contained in the sample? The three gas laws can be combined into a single equation known as the Ideal Gas Equation: PV = nRT
116
This equation can be rearranged into an equation known as the combined gas law, which holds true for a given amount of gas: PV = nR = k T P 1 V1
P2V2
=
T1
T2
The combined gas law reduces to Boyle’s Law, if temperature is kept constant (i.e. T1#=#T2): P1 V1 = P2V2 It also reduces to Charles’s Law, if pressure is kept constant (i.e. P1 = P2): V1
=
T1
V2 T2
It will also show that pressure of a gas is directly proportional to its absolute temperature, if the volume is kept constant (i.e. V1 = V2): P1 T1
=
P2 T2
The value of R can be calculated from the molar volume at 0oC and 1 atm (V#=#22.4#L). R
=
PV nT
=
(1 atm) (22.4 L) (1 mol) (273 K)
=
0.0821
atm L mol K
Point out the importance of R, which is known as the gas constant, and that it is one of the few universal constants, i.e. its value is the same anywhere and anytime. 117
Ask them to solve the following problems: 1. A gas sample occupies a volume of 12.0 L at 50oC and 700 Torr. How many moles of gas are contained in the sample? 2. Calculate the volume that will be occupied by 20.0 g carbon dioxide at 25oC and 1.25 atm. 3. What would be the pressure of 6.40 g oxygen gas in a vessel with a volume of 4.5 L at 20oC? The ideal gas equation can be transformed into an expression involving density. The number of moles n can be expressed in terms of mass and molar mass (or weight and molecular weight, respectively): n = w/M Introducing this into the ideal gas equation gives: P V = (w / M) R T which can be rearranged into P M = (w / V) R T The term w / V is recognized as equal to density, d, so that the equation becomes: PM = dRT Note that if the value of R as 0.0821 (atm L) / (mol K) is used, the unit for density in the equation should be g#/#L. 118
Answer Key 1. 0.0347 mol 2. 8.90 L 3. 1.07 atm
For a given gas (i.e. M = constant ) at a given pressure, the equation can be reduced into dT = K
or
d1 T1 = d2 T2 This equation shows that the density of a gas is inversely proportional to its temperature. This means that hot air has a lower density than cold air. The relationship between density and temperature can explain: a. The principle of the hot air balloon b. The principle behind passive cooling in building design PART TWO Highlight that many of the gases encountered in the surrounding are mixtures. Point out that the Ideal Gas Equation can also be applied to not only to pure gases, but also to mixtures of gases. Present a system composed of three gases contained in a vessel of volume V and kept at a temperature T. The number of moles of each gas is n1, n2 and n3, for gases 1, 2 and 3, respectively, so that the total number of moles of gases is
n total = n1 + n2 + n3 The pressure of the mixture is given by the Ideal Gas Equation:
P mixture V = n total R T
119
Teacher Tip Assign them to read on ‘passive cooling’ from internet resources. This reading will make them realize the relevance of the gas laws in building design.
Expressing ntotal in terms of the number of moles of each gas and solving for Pmixture will result in
P mixture
=
( n 1 + n 2 + n 3 ) RT
Teacher Tip The learners could be asked to write the resulting expression for each step of the derivation.
V
If the right-hand side of the equation is expanded, the expression becomes
P mixture
n 1 RT =
V
+
n 2 RT V
+
n 3 RT V
Let them realize that the term ni R T / V is equal to P and see that the previous equation can be written as
P mixture = P1 + P2 +P3 The pressures P1, P2, and P3, called partial pressure of each gas, corresponds to the pressure that the gas will exert in a volume equal to that of the mixture. According to this expression, the total pressure of a gas mixture is equal to the sum of the partial pressure of each gas. This is known as Dalton’s Law of Partial Pressure. Ask them to solve the following problems: 1. In a gas mixture composed of N2, Ne, and He, the partial pressure of N2 is 0.50 atm, that of Ne is 1.1 atm, and that of He is 0.80 atm. What is the total pressure of the mixture? 2. A sample of oxygen gas, which is saturated with water vapor, is kept in a 10-L vessel at 30oC and has a pressure of 758 Torr. If the pressure of the water vapor at this temperature is 31.8 Torr, what would be the pressure of the dry oxygen?
120
Answer Key 1. 2.4 atm 2. 726.2 Torr
The application of Dalton’s Law can also yield information about the composition of the mixture, in terms of the mole fraction of each component. Let them write on the board the expression for the pressure of gas 1 and that of the mixture: P1
=
n 1 RT
P mixture
V
=
Teacher Tip They could be asked to write the resulting expression for each step of the derivation.
n total RT V
Dividing P1 by Pmixture gives the following expression: P1 P mixture
=
n1 n total
=
x1
The term at the right-hand of the equation is actually a fraction, i.e. part divided the whole, and is known as mole fraction X1. Rearranging the expression leads to an important relationship:
P1 = Pmixture X1 Ask them to solve the following problems: 1. In a gas mixture composed of N2, Ne, and He, the partial pressure of N2 is 0.50 atm, that of Ne is 1.1 atm, and that of He is 0.80 atm. Calculate the mole fraction of each gas. 2. A gas mixture contains 2.5 mol N2 and 9.7 mol CO2, and has a pressure of 2.3 atm. What is the partial pressure of each gas? PART THREE As an introduction to this section, make the learners realize that reactions involving gases are common, such as the burning of fuel and the digestion of sugars: 2 C4H10 (g) + 13 O2 (g) g 8 CO2 (g) + 10 H2O (l) C6H12O6 (aq) + 6 O2 (g) g 3 CO2 (g) + 6 H2O (l) 121
Teacher Tip 1. N2: X = 0.21; Ne: X = 0.46; He: X = 0.33 2.
N2: P = 0.47 atm; CO2: P = 1.83 atm
Let them recall the basic principle of reaction stoichiometry, which is expressed by the balanced chemical equation.
Note Reaction stoichiometry – the relationship between the moles of reactants and products in a reaction
Highlight the following principles learned in the previous discussion:
Answer Key 1. They can be guided in solving the problem by asking them to solve first the number of moles of NaN3: mol NaN3 = 2
1. The Ideal Gas Equation enables the calculation of the number of moles of a gas from its pressure, volume, and temperature: n
=
PV
Then, ask them to solve the number of moles of N2 produced, using the balanced equation: mol N2 = 3
RT
2. The volume of a gas at standard temperature (0oC) and standard pressure (1 atm) can provide information about the number of moles of the gas, through the known molar volume under the standard condition: V STP n
Finally, ask them to recall the molar volume at STP (22.4 L) and use it to solve for the final answer: VN2 = 67.2 L 2.
= 22.4
Point out that these principles are useful in calculating the amount of gases involved in a reaction. Ask them to solve the following problems: 1. The airbag is a safety device used in cars to cushion the passenger during a crash. It involves the following chemical reaction which is triggered by an impact: 2 NaN3 (s) g 2 Na (s) + 3 N2 (g) Calculate the volume of N2 gas (measured at STP) that can be produced from 130.0 g of NaN3 (molar mass = 65).
122
The mole of C2H2 is first calculated through the molar volume (22.4 L) at STP: mol C2H2 = 0.446 From this, the mole of CaC2 is calculated: mol CaC2 = 0.446 The weight of CaC2 is calculated from the number of moles: g CaC2 = 28.6 g
2. Acetylene is formed by the reaction of water with calcium carbide, according to the following equation: CaC2 (s) + 2 H2O(l) g Ca(OH)2 (aq) + C2H2 (g)
Answer Key 3. The mole of C2H2 is first calculated through the molar volume (22.4 L) at STP: mol of C2H2 = 2.23 From this, the mole of CO2 is calculated: mol CO2 = 4.46
How many grams of CaC2 would be needed to produce 10.0 L (measured at STP) acetylene? 3. The reaction involved in the explosive combustion of acetylene is:
The volume (measured at STP) of CO2 is calculated from the number of moles: V CO2 = 100.0 L This answer can also be obtained through the ratio of the mol CO2 to mol C2H2.
2 C2H2 (g) + 5 O2(g) g 4 CO2 (g) + 2 H2 (g) How many L of CO2 gas (measured at STP) will be formed during the combustion of 50.0 L C2H2 gas (measured at STP)?
Teacher Tip They can be given a set of problems involving reaction stoichiometry as homework.
PART FOUR Emphasize to the learners that the gas laws summarize the general behavior of gases. Through these laws, the behavior of gases can be predicted. However, no explanation is given for this behavior. Point out that the explanation is provided by the Kinetic Molecular Theory. The theory assumes a model which can be used to explain why gases behave the way they do. The model is described through a set of postulates: 1. Gases are made up of very small molecules,which are separated by a very great distance between them. The dimension of the molecules is very much smaller than the distance between them. 2. Because of the very great distance between them, the force of attraction between the molecules is negligible. The molecules are independent of each other. 3. The molecules are in constant motion, moving in randomly in all directions. 4. Due to the great number of molecules and their random motion, it is unavoidable that the molecules will collide with each other and with the walls of the container. 123
Teacher Tip To highlight each postulate, write the keyword for each statement on the board.
5. During these collisions, there is no change in the momentum of the molecules. 6. The average kinetic energy of the molecules is determined only by the absolute temperature of the gas. Ask them to draw a representation of the model of the Kinetic Molecular Theory. The model should be: Teacher Tip The model would be familiar to them from the Science course in junior high school. What might not have been emphasized then is the motion of the molecules.
Ask them to apply this model to explain some properties of gases: a. Why can gases be compressed? Make them see that because of the great distance between them, gases can be forced to be close to each other by compressing it. b. Why does the volume of a gas decrease as the pressure is increased at constant temperature? This question asks for an explanation for the behavior described by Boyle’s Law. The answer would be similar to the previous question on the compressibility of gases. As the molecules become closer to each other, the volume of the gas becomes smaller. c. Why do gases exert pressure? Remind them that pressure is actually a force acting on a unit area. Help them realize that the collision of the molecules with the walls of the container produces a force acting on the wall. 124
d. Why does the volume of a gas increase as it is heated at constant pressure? This question asks for an explanation for the behavior described by Charles’s Law. Help them recognize that Postulate 5 expresses the effect of temperature on gases. According to this postulate, if the temperature is increased, the kinetic energy of the molecule increases. The increased kinetic energy makes the molecules to move faster and farther apart from each other, leading to a greater volume. Point out that a mathematical treatment of the Kinetic Molecular Theory would lead to an equation for the root-mean-square velocity of the molecule:
This equation clearly shows that as the temperature increases, the velocity of the molecule increases.The gas molecules move faster at a higher temperature. It also shows that as the molar mass M of the molecule increases, the velocity of the molecule decreases. If the velocity of two molecules of molar mass M1 and M2 are compared, the result is 2 1
The velocity of the molecules determines the rate of diffusion of the gases. The relationship between diffusion rate and molar mass has been verified by experiments, and is known as Graham’s Law of Diffusion. Ask them to imagine that two bottles are placed at opposite ends of the room at equal distance from them. One bottle contains ammonia gas, NH3 (M = 17) which has a pungent odor, and the other contains hydrogen sulfide, H2S (M = 34) which has an odor like that of a rotten egg. Which odor will they sense first?
125
Teacher Tip The root-mean-square velocity is the squareroot of the mean of the square of the velocities of the molecules:
ENRICHMENT
1. As mentioned in the Delivery, at the end of each session, summarize the concepts discussed in the session. 2. Conduct a session on problem solving to provide more exercises on the application of the gas equations. 3. Conduct a laboratory activity on Graham’s Law of Diffusion.
EVALUATION (30 minutes) Check-up quiz
1. Under which of the following volumes will 1.00 mol of an ideal gas exhibit the greatest pressure at 300 K? A. 0.01 L
C. 1.00 L
B. 0.10 L
D. 10.0 L
2. How will the volume of 0.50 mol of a gas behave if the temperature is raised from 30oC to 60oC at constant pressure? A. The volume will increase.
C. The volume will be doubled.
B. The volume will decrease.
D. The volume will be halved.
3. Which among the following systems will have the greatest volume at STP? A. 1.00 g N2 gas (M = 28 g/mol) B. 1.00 g NH3 gas (M = 17 g/mol) C. 1.00 g CO2 gas (M = 44 g/mol) D. 1.00 g He gas (M = 4 g/mol) 4. How will the density of a gas vary if its temperature is increased from 25oC to 50oC at constant pressure? A. The density of the gas will not change! ! B. The density of the gas will increase C. The density of the gas will decrease D. The density of the gas will double
126
5. Which of the following volumes of oxygen will contain the greatest number of molecules at 300K and 1 atm pressure? A. 0.01 L
C. 1.00 L
B. 0.10 L
D. 10.0 L!
6.!In which of the following gas mixtures of N2(g) and He(g) is the partial pressure of He(g) the greatest? A. 2 moles N2(g) and 3 mole He(g) B. 3 moles N2(g) and 1 mole He(g) C. 4 moles N2(g) and 2 mole He(g) D. 5 moles N2(g) and 5 mole He(g) 7. Hydrogen, H2(g),reacts with oxygen, O2(g), to form water H2O(l): !
2 H2(g) + O2(g) ##!###2 H2O(g)# How many liters of oxygen gas, measured at STP, will be needed to react completely with 10.0 L hydrogen gas, also measured at STP? A. 5.00 L
C. 20.0 L
B. 10.0 L
D. 100.0 L
8. Which of the following postulates of the Kinetic Molecular Theory for gases can explain why gases exhibit pressure? A. The molecules are in constant random motion! B. The molecules collide with the walls of the vessel! C. The distance between the molecules is great! D. The molecular kinetic energy depends on temperature! 9. How will the velocity of a gas molecule vary if its molecular weight is increased from 32 g mol-1 to 64 g mol-1? A. The velocity will increase
!
B. The velocity will decrease C. The velocity will double D. The velocity will remain the same
127
General Chemistry 1
120 MINS
Lesson 16: Gases (Laboratory)
Lesson Outline
Content Standard The learners demonstrate the ability to explain experimental observations using the laws and theories learned in the lecture course.
Introduction
Communicating Learning Objectives
5
Motivation
Inquiry
5
Instruction
Pre-activity
60
Enrichment
Problem solving
25
Learning Competencies
Evaluation
Post-laboratory
25
At the end of the lesson, the learners: 1. Observe and measure the difference in the diffusion rate of two gases. (STEM_GC11KMT-Ij-51);
Materials Laboratory glassware
Performance Standards The learners shall be able to compare the rates of diffusion of two gases and explain the observed behavior.
Resources (1) Laboratory experiments found in the internet, such as: Fasano, Janet. Graham’s Law lab [PDF document]. Retrieved from Needham Public Schools: http://fcw.needham.k12.ma.us/~Janet/FOV1-00108AC5/Graham's%20Law %20Lab.pdf
128
INTRODUCTION (5 minutes)
1. State the objective of the experiment that the learner will be performing. 2. Ask them to recall the diffusion property of gases and explain in terms of the kinetic molecular theory. 3. Point out safety measures to be observed.
Teacher Tip A laboratory experiment sheet has to be prepared and distributed to the learners. The experiment found in the internet could be revised or simplified to suit the available facilities in the laboratory.
Point out why we can smell the odor of a fruit (such as durian) or a flower (such as sampaguita) from a distance.
Teacher Tip An alternative experiential approach to the motivation can be done using an open bottle of perfume in front of the class.
INSTRUCTION (60 minutes)
Remind them to observe safety precautions during the experiment.
MOTIVATION (5 minutes)
1. Provide each group with the necessary materials. 2. Ask them to follow the procedure in the experiment sheet. Sample Problems a. Gas X has a molar mass of 72 g/mol and Gas Y has a molar mass of 4 g/mol. How much faster or slower does Gas Y effuse from a small opening than Gas X if they are at the same temperature? b. If the density of hydrogen is 0.090 g/L and its rate of diffusion is 5.93 times that of chlorine, what is the density of chlorine?
ENRICHMENT (25 minutes)
Assign them to solve some problems involving Graham’s Law.
Teacher Tip Provide them with the worksheet that they have to fill up. It could include some more questions.
EVALUATION (25 minutes)
Ask them to submit a report on the experiment.
129
DIFFUSION OF GASES Introduction
One of the properties of gases is its ability to diffuse easily. This property can be explained by the motion of the gas molecules and the absence of intermolecular forces of attraction. As a result of this property, a gas spreads easily in the air and fills up all available space. In this experiment, the diffusion of two gases will be investigated and their relative rates of diffusion will be measured. The gases will be confined in a glass tube and will be introduced at the opposite ends of the tube.The mixing of the two gases will be indicated by the formation of a white solid in the tube.
Materials
a. Concentrated hydrochloric acid, HCl b. Ammonia solution, NH3 c. Glass tube d. Cotton buds (Q-tips)
Procedure
1. Set the glass tube against a black background and place markings on both ends to indicate where the cotton tips will be introduced. 2. Place two drops of concentrated HCl in one cotton bud, and two drops of NH3 solution in the second cotton bud. Caution: These solutions can irritate your skin. Use gloves, if possible. 3. Simultaneously insert the cotton buds in the opposite ends of the glass tube. 4. Note which part of the tube a white ring will form. Mark this part and measure its distance from the HCl end and from the NH3 end. 5. Repeat Steps 1 to 4 to provide a duplicate measurement. This will be used to check the repeatability of the results. 6. Dispose the cotton buds in the designated waste container.
130
Treatment of results
1. Record the distance of the white ring formed in the tube from the ends where the two gases were introduced. Calculate the ratio of these two distances. This ratio is equal to the ratio of the rates of diffusion of the two gases. 2. Obtain the molar mass of HCl and NH3, and calculate the ratio of the diffusion rates of the two gases using Graham’s Law of Diffusion. 3. Compare the observed and predicted ratio of the diffusion rates.
TRIAL 1
TRIAL 2
TRIAL 1
Distance from the HCl end (dHCl)
Molar mass of HCl (MHCl)
Distance from the NH3 end (dNH3)
Molar mass of NH3 (MNH3)
Ratio of distances (dHCl) / dNH3)
Calculated ratio of diffusion rates (dHCl) / dNH3)
TRIAL 2
EVALUATION EXCEEDS EXPECTATIONS
MEETS EXPECTATIONS
NEEDS IMPROVEMENT
NOT VISIBLE
The learner:
The learner:
The learner:
The learner:
i.
i.
i.
i.
performed the experiment correctly;
performed the experiment correctly;
performed the experiment correctly;
ii. described the results correctly; and
ii. described the results correctly; and
ii. described the results correctly; but
iii. discussed the results of the experiment very well.
iii. discussed the results of the experiment well.
iii. did not discuss the results of the experiment.
131
did not do the assigned task.
General Chemistry 1
Lesson 17: Electromagnetic Waves, Planck’s Quantum Theory, and Photoelectric Effect
120 MINS
Content Standard The learners demonstrate an understanding of the quantum mechanical description of the atom and its electronic structure. Performance Standards The learners can describe the dual nature of an electron. Learning Competencies At the end of the lesson, the learners: 1. Describe the quantum mechanical model of the atom (STEM_GC11CB-IIa-b-52) a. Identify the inadequacies of the classical physics in explaining experimental results that brought about the quantum theory b. Discuss Planck’s quantum theory c. Describe the particle-wave duality of light in relation to the photoelectric effect d. Recognize the contribution of scientists to the development of the Quantum Mechanical Model of the Atom Specific Learning Competencies At the end of the lesson, the learners will be able to: 1. Describe the characteristics of a wave; 2. Relate the order of the regions of the electromagnetic spectrum in terms of their wavelength and frequency; 3. State Planck’s equation; 4. Solve problems related to electromagnetic radiation, its energy, wavelength, and frequency; 5. Describe the particle-wave duality of light; and 6. Recognize technological applications of the photoelectric effect 132
Lesson Outline Introduction
Presentation of Learning Objectives and Keywords
10
Motivation
Thermal Imaging Infrared Photography
10
Instruction, Delivery and Practice
I. II. III. IV.
Enrichment
Revisit Introductory Questions
Evaluation
Take Home Activity
The Characteristics of a Wave Planck’s Quantum Theory The Photoelectric Effect The Particle-Wave Duality of Light
90
10
Materials Calculator Resources (1) Chang, Raymond and Goldsby, Kenneth A. (2016). Chemistry (12th ed). New York: McGraw-Hill. (2) Petrucci, Herring, Madura, and Bissonnette (2011). General Chemistry and Modern Applications, 10th Ed. Pearson Canada, Inc. (3) Zumdahl, S.S. and Zumdahl, S.A (2013).Chemistry, 8th ed. Cengage Learning (4) Infrared Image Gallery: http://coolcosmos.ipac.caltech.edu/ image_galleries/ (5) PAGASA weather satellite maps: http://meteopilipinas.gov.ph/
INTRODUCTION (10 minutes)
1. Introduce the following learning objectives using any of the suggested protocol At the end of Part I, I will be able to: a. Describe the characteristics of a wave b. Relate the order of the regions of the electromagnetic spectrum in terms of their wavelength and frequency c. State Planck’s equation. d. Solve problems related to electromagnetic radiation, its energy, wavelength, and frequency. e. Describe the particle-wave duality of light f.
Recognize technological applications of the photoelectric effect
2. Present the keywords for the concepts to be learned:
a. wave h. Hertz b. frequency
i.
speed of light
c. wavelength
j.
blackbody radiation
d. amplitude
k. quantum theory
e. crest
l.
f.
trough
photoelectric effect
m. particle-wave duality of light
g. period 3. Certain experimental results observed at the beginning of the 20th century could not be explained by classical physics. These included the blackbody radiation, photoelectric effect, and the emission spectrum of hydrogen. The new age of physics began when the German physicist, Max Planck proposed his quantum theory of energy. The lesson will discuss the emergence of this new theory and the contributions of Max Planck and Albert Einstein to the development of the quantum theory.
133
4. Post on the board the following essential questions that will be answered after the discussion a. What is a wave? b. What is a particle? c. Is the electron a wave or a particle?
MOTIVATION (5 minutes)
Engage the students in a discussion and ask them the following questions: 1. What is a thermal imaging infrared photography? Thermal imaging infrared photography detects infrared light and converts this to an electronic signal that is processed to produce a thermal image. 2. Give some uses of thermal imaging photography. Examples of uses: Military operations Construction – check efficiency of insulation and detect where there are heat leaks; check electrical wirings in houses to see where there are overheating joints Fire fighters use this to locate hotspots in a building or locate people who are trapped. PAGASA weather maps showing warm and cool areas of the ocean
INSTRUCTION (90 minutes) I.
THE CHARACTERISTICS OF A WAVE
It was in the early 1900’s that a new way of looking at energy and matter began. It stemmed from Max Planck’s idea about blackbody radiation and culminated in Schrodinger’s wave equation known also as the wave function, ψ (psi), which described the hydrogen atom. 1. First let us define a particle and a wave. a. A particle is an object which has distinct chemical or physical properties such as volume or mass. b. A wave is a disturbance that travels from one location to another location. The highest peak of the wave is called the crest and the lowest point is named as the trough. c. The wave has distinct characteristics that include amplitude, wavelength and frequency. 134
Teacher Tip Check the Infrared Image Gallery site at http:// coolcosmos.ipac.caltech.edu/image_galleries/. If possible, print a few infrared pictures beforehand and pass these to the students during class. Alternatively, ask the students to visit the website from their homes, their mobile phones, or from the school library. Ask the students to go to the PAGASA website at http://meteopilipinas.gov.ph/ to see the IR weather satellite map.
The amplitude is defined as one-half the distance from crest to trough. The wavelength (symbolized by the Greek letter lambda, λ) is the distance from crest to crest or from trough to trough. d. Frequency (symbolized by the Greek letter nu, ν) is defined as the number of waves passing a fixed point in a specified period of time. Frequency has units of waves per second or cycles per second. Another unit for frequency is the Hertz (abbreviated Hz) where 1 Hz is equivalent to 1 cycle per second.
e. The period of a wave is the time for a particle on a medium to make one complete vibrational cycle.
135
Teacher Tip Remind the learners that i. A wave is characterized by its wavelength, frequency, and amplitude. ii. The wavelength, λ , has units of length (i.e. m, mm, nm, etc.) iii. Frequency, v", has the units of Hz (1 Hz = 1cycle per second). Sometimes it is also in terms of (1/time) for example sec-1.
The figures above show two waves travelling between two points at a constant speed. Note that the wave with longer wavelength has the lower frequency, and the wave with the shorter wavelength has higher frequency. Thus, wavelength, λ, and frequency, v, are indirectly related to one another. The wavelength of the wave multiplied by the frequency of the wave corresponds to the speed, µ, of the wave. In an equation form, λν = µ Waves can be classified as mechanical or electromagnetic waves. A mechanical wave requires a medium for it to travel, i.e. the sound wave, water wave, etc. An electromagnetic wave is a wave that is capable of transmitting its energy through an empty space or vacuum. Light is considered to be a electromagnetic wave. In electromagnetic radiation, the frequency of the wave when multiplied by its wavelength corresponds to the speed of light, c, as shown in the equation λν = c where c = 3.00 x 108 m/s. The figure on the right shows the various types of electromagnetic radiation, which differ from one another in terms of wavelength and frequency. The shortest waves which have the highest frequency, like the gamma rays, result from the changes within the nucleus of the atom. The visible light waves, with wavelength that range from about 400-700 nm, are produced by the motions of electrons within the atoms and molecules. The longest waves are those emitted by the antennas of broadcasting stations. ELECTROMAGNETIC SPECTRUM. Retrieved (https://upload.wikimedia.org/wikipedia/commons/ thumb/f/f1/EM_spectrum.svg/2000px-EM_spectrum.svg.png) (07/02/2016 12:56 PM)
136
Give the learners some practice exercises: 1. If the wavelength is decreased to half its original length, what happens to the frequency? Answer: The frequency is doubled. 2. A yellow light emitted by a sodium vapor lamp has a wavelength of 589 nm. What is the frequency of the yellow light? v## =
c# λ
=
3.00#x#10#8#m/s# 589#nm
x
10#9#nm# 1#m
!
5.09"x"10"14"s"+1""or##5.09"x"10"14"Hz
3. A radio station broadcasts at a frequency of 590 KHz. What is the wavelength of the radio waves?
λ## =
c# v
3.00#x#10#8#m/s# =
590#kHz
x
1#kHz# 1,000#Hz
1#Hz#
x
=##508#m##="""5.1"x"10"2"m
1#/#s
4. A particular electromagnetic radiation was found to have a frequency of 8.11 x 1014Hz. What is the wavelength of this radiation in nm? To what region of the electromagnetic spectrum would you assign it? II. PLANCK’S QUANTUM THEORY When an object is heated, the electrons on the surface are thermally agitated and begin to emit radiation. Physics around the 1900s was concerned with the spectrum of the light emitted by heated bodies, particularly by black bodies. A blackbody is a material that absorbs all radiation that falls on it and is therefore a perfect absorber. When such a blackbody is heated, it was expected to emit at every wavelength of light that it is able to absorb. Classical physics predicted that the maximum wavelength emitted by the blackbody would be infinite. However, results proved otherwise and classical physics could not explain the resulting spectrum of blackbody radiation. Experimental results showed that while blackbodies emitted radiation at various wavelengths, they showed a maximum wavelength (not infinite) that shifted toward lower wavelengths as the temperature increased. Planck made a radical proposal to explain the experimental results of the blackbody radiation. He proposed that the atoms on the surface of the heated solid could absorb energy only in discrete quantities or quanta, and not continuously as assumed by classical physics. The energy absorbed 137
Note Give the practice exercises as a seatwork. Ask some students to show their solutions on the board and explain their answer.
or released by any oscillator are in integer multiples, or quanta, of hν. This became known as Planck’s equation.
E"="hν" Energy, E, is equal to frequency, ν, multiplied by Planck’s constant, h, with a value of 6.626 x 10-34 J s. According to quantum theory proposed by Planck, the amount of energy emitted or absorbed by a body can have values of hν, 2 hν, 3 hν, 10 hν, but never 4.8 hν or 0.25 hν.
Teacher Tip Make sure the learner understands the meaning of integral multiples and quanta. There is no need for the learner to memorize Planck’s constant, h. The value should be given to the student during exams and quizzes.
And because v=c/λ, the equation can also be expressed as
E"="hν"="h""
c# λ
Ask the learners the following questions: a. Which is a quantized way of getting from the 1st to the 2nd floor of a building – using the stairs or using a ramp? b. Give some examples from daily life that shows quantization. III. EINSTEIN’S EXPLANATION OF THE PHOTOELECTRIC EFFECT Another stumbling block for classical physics was the photoelectric effect experiment. According to classical physics, when light hits a metal surface, the electrons in the metal should slowly absorb energy from the light until they have enough energy to be emitted to produce a current. It also predicted that as the intensity of the incident light increases, the kinetic energy of the emitted electrons should increase. However, the experiment did not support these predictions but provided the following observations:
Answer Key a. Using the stairs is quantized. One can take 1 step at a time or 2 steps at a time in going up to the second floor but never 1.6 steps at a time.! b. Some answers could be: The smallest denomination given by ATM machines; chairs around the table, etc.
a. When light is made to hit a metal surface, there is a threshold frequency below which no electrons can be ejected regardless of the intensity of the incident light. b. Above the threshold frequency, the number of ejected electrons was proportional to the intensity (or brightness) of the incident light but their energies were not. c. Above the threshold frequency, the kinetic energy of the emitted electrons increased linearly with the frequency of the incident light. 138
Source: Retrieved (https://upload.wikimedia.org/ wikipedia/commons/f/f5/Photoelectric_effect.svg) 11/02/16, 08:12 am
These results could not be explained by the wave theory of light. Waves can have any amount of energy - big waves have a lot of energy, small waves have very little. And if light is a wave, then the brightness of the light affects the amount of energy - the brighter the light, the bigger the wave, the more energy it has. IV. THE PARTICLE-WAVE DUALITY OF LIGHT Einstein proposed that the only way to explain the photoelectric effect was to say that instead of being a wave, as was generally accepted, light was actually made up of lots of small packets of energy called photons that behaved like particles. Each photon has energy given by the equation E#=#hν# Where is the frequency of the light and h is Planck’s constant: h = 6.626 x 10-34 J s. In explaining the results of the photoelectric effect experiment, the energy, hν, of the incident light is used to remove the electron from the surface of the metal. If the electron is tightly held by the metal and the energy of the incident light (corresponding to the threshold frequency) is not sufficient, no electron will be ejected. If the energy of the incident light is sufficient, it will use the energy to eject the electron; the rest will be given off as the kinetic energy of the electron. In equation form, this is given by hν#=#W#+#K.E.# where W is the work function (the energy needed to eject the electron) and K.E. is the kinetic energy. K.E.="hν"+"W" This explains the observation that the kinetic energy of the emitted electrons varied linearly with the frequency of the incident light. 139
Light is a wave as shown by different experiments like the diffraction of light by a prism to yield the visible spectrum. However, the photoelectric effect experiment showed that light also behaves like a particle. Thus light has both wavelike and particlelike properties. This concept is called the particle-wave duality of light.
Note The apparent mass of a photon of light with wavelength, λ, can be expressed from the relationship of Einstein’s famous energy equation from the theory of relativity:
E"="mc"2" And the energy equation by Planck:
Ephoton"="hv"=""
hc" λ
And is given by
E" λc/"λ" h" m"=""""""""""""""=""""""""""""""""""""""=" 2 2 c c λc Note that the apparent mass of a photon depends on its wavelength. However, a photon does not have a mass in a classical sense.
Give the learners practice exercises: 1. The work function or the energy needed to eject an electron in cesium metal is 3.42 x 10-19 J. If an incident light of frequency 1.00 x 1015 s-1 is used to irradiate the metal, will electrons be ejected? Show your calculations. The energy of the incident light can be calculated by E""=""hv""=""(6.626"x"10"+34"Js)"(1.00"x"10"15"s"+1")""=""6.626"x"10+19"J" This energy is greater than the work function of cesium metal. Therefore, electrons will be ejected from the metal. 2. What will be the kinetic energy of the ejected electron? KE""=""hv"+"W"=""(6.626"x"10"+19"J"+"3.42"x"10"+19"J""=""3.21"x"10+19"J" 140
Teacher Tip Discuss some problems in class. The rest may be given as a seatwork. The problems may be assigned individually or to groups. Then ask the learners to show their work on the board for discussion with the entire class. This may also be used for evaluation.
3. The blue color in fireworks is due to copper (I) chloride , (CuCl), is heated at a temperature of 1200 oC. What is the energy emitted at 4.50 x 102 nm by CuCl? Solution:))The)quantum)of)energy)can)be)calculated)from)Planck’s)equation)!E"="hv." a. The frequency can be calculated from the equation, c = λv ; rearranging the terms, we have: v"" =
c" λ
=
3.00"x"10"8"m"/"s" 4.50"x"10"+7"m
=""6.67"x"10"14"s"+1
b. Then, solve for the energy using Planck’s equation.
!"E"="hv"="(6.626"x"10"+34"Js")"("6.67"x"10"14"s"+1")"="4.41"x"10"+19"J"" This means that, CuCl emitting a blue light at 450 nm can lose energy only in the increments of 4.41 x 10-19 J, the size of the quantum in this case. 4. There are three types of UV radiation classified by wavelength: UVA (320 – 400 nm), UVB (290 – 320 nm), and UVC (180 – 280 nm). Which type of UV has the lowest energy? 5. A photon of ultraviolet (UV) light possesses enough energy to mutate a strand of human DNA. What is the energy of a single UV photon having a wavelength of 5.00 nm? 6. Compare the energy (in joules) of (a) photon with a wavelength of 5.00 x 104 nm and (b) photon with a wavelength of 5.00 x 10-2 nm. At what regions in the spectrum do the samples come from? Relate the relationship of the wavelength of a radiation to the energy. 7. Chlorophyll absorbs light energies of 3.06 x 10-19J/photon and 4.41 x 10-19J/photon. To what color and frequency do these absorptions correspond? 141
8. The protective action of ozone in the atmosphere comes through ozone’s absorption of UV radiation in the 230 to 290 nm wavelength range. What is the energy, in kJ/mol, associated with radiation in this wavelength range? 9. The work function of potassium metal is 3.68 x 10-19 J. Which of the following will cause electrons to be ejected from the surface of potassium metal? a. Red light ( λ = 7.00 x 10 -7 m) b. Green light ( λ = 5.51 x 10 -7 m) c. Violet light ( λ = 4.00 x 10 -7 m)
ENRICHMENT (10 minutes)
1. Go back to the essential questions presented during the introduction and ask the students to give their answers to check their understanding of the lesson. a. What is a wave? b. What is a particle? c. Is the electron a wave or a particle? 2. Return also to the motivation questions on thermal imaging infrared photography. Relate these to the lesson. a. Which has longer wavelength, IR or visible radiation? b. Which will have more energy, IR or visible radiation? 3. Discuss the modern uses of the photoelectric effect. After a century, Einstein's work on the photoelectric effect gave way to new and very useful technologies a. Photocell found in automatic door openers; b. Ruby lasers, red light emitters used to read bar codes and night vision devices c. Medical and dental devices d. Other image processing technologies. 142
Answer Key 4. UVA 6.a. E = 3.98 x 10 -21 J 6.b. E = 3.98 x 10 -15 J from x-ray region 6.c. Wavelength is inversely proportional to energy.
TAKE HOME ACTIVITY: Scientists on Parade Explain to the students the take home activity. This may be assigned to individual students or to groups. Ask the students to prepare a poster which illustrates or describes the role of the scientists listed below in the development of the quantum mechanical description of the atom. If an LCD projector and laptop are available, the students may present a 7-slide power point presentation. Ask the students also to reflect on the lives of these scientists and relate what they liked in the life story of the scientists and what attributes did they possess that are worthy to emulate.
Make a profile of each scientist and explain their contributions to the behavior of the electron. Samples 1. Max Planck 2. Albert Einstein 3. Niels Bohr
Give the students about 1 to 2 weeks to work on the project. They are to present their work in class at the end of the lesson on the quantum mechanical description of the atom.
4. Louis de Broglie 5. Werner Heisenberg 6. Erwin Schrödinger
EVALUATION CRITERIA
EXCEEDS EXPECTATIONS
MEETS EXPECTATIONS
NEEDS IMPROVEMENT
NOT VISIBLE
Information
3-4 unique details or examples provided; content is complete; all information clearly relates to topic
2-3 details are provided information clearly relates to topic; diagrams (if present) relate to topic and add to clarity
1-2 details are provided; some information provided is not closely related to topic
Incomplete information; irrelevant ideas or examples included
Organization
Clear organizational method chosen suits work; content flows in a clear pattern; reader is able to concentrate on the information
Information generally organized; the content flows nicely; the reader has no difficulty following the information despite a minor error or two
2-3 gaps or out of sequence information cause viewer or reader to re-read numerous times for clarity
Information is presented out of sequence
Presentation
An occasional grammar or spelling error may result from risk-taking; materials are organized and glued down; presenting a polished, pleasing result
3 or fewer minor grammar or spelling errors; mostly aesthetically pleasing, some messy parts; evident care of project
Many minor grammar or spelling errors; messy; inconsistent care in attaching materials or drawing or writing care of the project is inconsistent
Major and minor errors in grammar and spelling; information attached without attention to pleasing effect; care of project not evident (edges rolled, etc.)
Creativity
Images or layout show original ideas; reader is surprised, interested and pleased
Images or layouts use a common pattern which get intended audience interested
Images or layout reproduce common patterns, and give viewer/reader what he/she is accustomed to
Images or layout copied from others or standard or so sketchy that intent can’t be determined
143
General Chemistry 1
Lesson 18: Emission Spectrum of Hydrogen, and Dual Nature of Matter Content Standard The learners demonstrate an understanding of the quantum mechanical description of the atom and its electronic structure.
Lesson Outline
Introduction
Communicating Learning Objectives
Motivation
Recall
Learning Competencies At the end of the lesson, the learners: 1. Describe the quantum mechanical model of the atom (STEM_GC11CBIIa-b-52) a. Identify the inadequacies of the classical physics in explaining the emission spectrum of the hydrogen atom b. Discuss the use of quantum theory in explaining the emission spectrum of the hydrogen atom c. Describe the Bohr model of the atom and the inadequacies of the Bohr model d. Explain the wave-particle duality of matter.
Instruction
I.
Enrichment
Revisit Essential Question
Evaluation
Quiz
Specific Learning Competencies At the end of the lesson, the learners will be able to: 1. Explain the emission spectrum of hydrogen using the Bohr model of the hydrogen atom; 2. Calculate the energy, wavelength, and frequencies involved in the electron transitions in the hydrogen atom; 3. Relate the emission spectra to common occurrences like fireworks and neon lights; 4. Describe the Bohr model of the atom and the inadequacies of the Bohr model; 5. Explain the wave-particle duality of matter and 6. Perform calculations to determine wavelengths associated with moving bodies.
Material Calculator
144
120 MINS
The Emission Spectrum and the Bohr Theory II. Limitations of the Bohr Model III. The Dual Nature of the Electron; De Broglie Equation IV. Calculating the De Broglie Wavelength V. Experimental Evidence of De Broglie Wavelength
12 3 85
5 15
Resources (a) Chang, Raymond and Goldsby, Kenneth A. (2016). Chemistry (12th ed). New York: McGraw-Hill. (b) Petrucci, Herring, Madura, and Bissonnette (2011). General Chemistry and Modern Applications, 10th Ed. Pearson Canada, Inc. (c) Zumdahl, S.S. and Zumdahl, S.A (2013).Chemistry, 8th ed. Cengage Learning (d) Roque, et al. laboratory Manual in General Chemistry (2008). Philippine Normal University.
INTRODUCTION (12 minutes)
1. Introduce the following learning objectives using any of the suggested protocol (Verbatim, Own Words, or Read-aloud): At the end of Part II, I will be able to: a. Explain the emission spectrum of hydrogen using the Bohr model of the hydrogen atom b. Calculate the energy, wavelength, and frequencies involved in the electron transitions in the hydrogen atom. c. Relate the emission spectra to common occurrences like fireworks and neon lights. d. Describe the Bohr model of the atom and the inadequacies of the Bohr model e. Explain the wave-particle duality of matter f.
Perform calculations to determine wavelengths associated with moving bodies
2. Present the keywords for the concepts to be learned: a. Emission spectrum b. Rydberg’s constant c. Ground state d. Ground energy level e. Excited state f.
Excited energy level
g. Travelling wave h. Standing wave i.
De Broglie Equation
3. Post on the board the following essential questions that will be answered after the discussion Why do elements emit different colors when heated? What is the wave-particle duality of matter? 4. Review Rutherford’s nuclear model of the atom
145
MOTIVATION (3 minutes)
1. What causes the colors in fireworks displays? Ask the students to recall fireworks displays and ask them what they think give the colors in the fireworks? 2. What causes the colors in neon lights?
INSTRUCTION (85 minutes)
I. THE EMISSION SPECTRUM AND THE BOHR THEORY OF THE HYDROGEN ATOM When elements are energized by heat or other means, they give off a characteristic or distinctive spectrum, called an emission spectrum, which can be used to differentiate one element from another. While scientists recognized the usefulness of emission spectra in identifying elements, the origins of these spectra were unknown. From Rutherford’s theory, the atom was described to be mostly empty space having a very tiny but dense nucleus that contained the protons. The electrons whirled around the nucleus in circular orbits at high velocities. Classical mechanics and electromagnetic theory explained that any charged particle moving on a curved path would emit electromagnetic radiation. This implies that electrons would lose energy and spiral into the nucleus. Why this is not observed had to be explained.
Teacher Tip Here is another occurrence that classical mechanics is unable to explain.
In 1913, Niels Bohr proposed his model of the hydrogen atom to explain how electrons could stay in stable orbits around the nucleus. This model is no longer considered to be correct in all its details. However, it could explain the phenomenon of emission spectra. For his model of the hydrogen atom, Bohr made the following postulates:
Note Remember that the Bohr model is no longer considered correct. However, some of its features are still useful. One of this is the explanation of the emission spectrum. The limitations of the Bohr model will be pointed out in a later section.
a. Electrons go around the nucleus in circular orbits. However, not all circular orbits are allowed. The electron is allowed to occupy only specific orbits with specific energies. Therefore, the energies of the electron are quantized. b. If the electron stays in the allowed orbit, its energy is stable. It will not emit radiation and it will not spiral into the nucleus. c. If an electron jumps from one orbit to another, it will absorb or emit energy in quanta equal to #E"="hv
146
According to Bohr, the energy of the electron in the H atom is given by:
n%=%3 n%=%2 n%=%1
The negative sign is an arbitrary convention. A free electron is arbitrarily considered to have an energy of zero. A negative energy means that the energy of the electron is lower than the energy of a free electron. RH is the Rydberg constant for hydrogen equal to 2.18 x 10-18J. The number n is an integer equal to n = 1, 2, 3,…
Exercises
Teacher Tip It is important for the learner to understand the negative values for the energy. As the value gets more negative, the energy gets lower. As the value of the energy gets less negative, the energy gets higher. Comparing the energies for the first 3 energy levels, we see that E3 has the highest energy. It is less negative. Therefore, as n increases, energy increases. E1 = 2.18 x 10-18 J E2 = 0.545 x 10-18 J E3 = 0.242 x 10-18 J
1. What is the energy of the electron when it is in the first orbit, n=1?
E1 refers to the energy when the electron is in n=1. 2. What is the energy of the electron in orbit n = 2?
Do not make the learner memorise the value of RH. The value should be given to the learner.
147
3. What is the energy of the electron in orbit n = 3?
4. Plot the energies of the electron in n=1, n=2, n=3.
5. In which orbit will the electron have the highest energy, n=1, n=2, or n=3? Answer: n=3 6. As the value of n increases, what happens to the energy value of the electron? Answer: As n increases, energy increases.
148
E1 is the lowest energy and, therefore, the most stable state. It is called the ground state or the ground level. E2, E3, E4, etc. have higher energies and are less stable than E1. They are called excited states or excited levels. Note also that as the electron gets closer to the nucleus, it becomes more stable. When energy is absorbed by the atom, the electron gets excited and jumps from a lower orbit to a higher orbit. When electrons go from a higher energy level to a lower energy level, it emits radiation. According to Bohr, if an electron jumps from one orbit to another, it will absorb or emit energy in quanta equal to: ∆Ε = hν = h
Teacher Tip Note that this is the 3rd time that quantization of energy is used to explain an experimental result. The first was the blackbody radiation; the second was the photoelectric effect; and now the third is the emission spectrum of the hydrogen atom.
c ν
The Bohr model can explain the experimental emission spectrum of hydrogen which includes a wide range of wavelengths from the infrared to the UV region. These are summarized in the table below: SERIES
n final
n initial
Spectrum Region
Lyman
1
2, 3, 4
Ultraviolet
Balmer
2
3, 4, 5
Visible and ultraviolet
Paschen
3
4, 5, 6
Infrared
Brackett
4
5, 6, 7
Infrared 149
Teacher Tip Do not ask the learners to memorize the H atom emission spectrum series. This is only for illustration purposes only. We want them to understand the concept and not memorize.
Exercises 1. The electron in the hydrogen atom undergoes a transition from n=3 to n=2. a. Is energy absorbed or emitted? Answer: E is emitted because the electron goes from a higher energy level to a lower energy level. b. What is the energy involved in the transition?
c. What is the wavelength (in nm) corresponding to this transition?
d. What region of the electromagnetic spectrum will this be? Answer: This will be in the visible region. 2. Which transition of the electron in the hydrogen atom will involve the highest frequency? a. n = 5 to n = 3 b. n = 4 to n = 3 c. n = 5 to n = 2
150
Similarly, when substances like metal ions are subjected to heat, they absorb energy. The electrons jump from their ground state to an excited state. Once the electrons return from the excited state to the ground state, light is given off. The light emitted corresponds to the energy released. Example: Potassium emitted a pinkish purple color (approx. 400 nm) while lithium emitted a dark red hue (approx. 700 nm) when heated. a. What caused the color change during heating? b. Which element required the greater energy absorption for the electrons to be in the excited state and produce the observed emissions? Why? c. What is the relationship between wavelength and energy? II. THE LIMITATIONS OF THE BOHR MODEL OF THE ATOM Although the Bohr model could explain the emission spectrum of hydrogen and was an important step in the development of atomic theory, it has several limitations: a. It cannot explain the spectrum of atoms with more than one electron. b. It cannot explain the relative intensities of spectral lines (why are some lines more intense than others) c. It cannot explain why some lines are slit into several components in the presence of a magnetic field (called the Zeeman effect)
Teacher Tip In contrast to standing waves, travelling waves: are waves that travel in one-dimensional direction. Concrete examples of travelling waves can be seen from • Skipping rope held on one end, moved up and down, forming waves, from one end to the other • Ocean waves: the wind produces waves on the surface of water producing crests and troughs that travel great distances.
d. According to the Bohr model, when electrons go around the nucleus in certain orbits, its energy remains constant. But moving electrons would lose energy by emitting electromagnetic waves and the electron is expected to spiral into the nucleus. e. It violates the Heisenberg’s Uncertainty Principle. The Bohr model considers electrons to have a known radius and orbit which is impossible according to Heisenberg. This will be explained later in the next lesson.
For an allowed orbit, the circumference of the orbit must be equal to an integral number of wavelengths (a); otherwise the wave will cancel itself (b).
151
III. THE DUAL NATURE OF THE ELECTRON; DE BROGLIE’S EQUATION In 1924, Louis de Broglie made a bold proposition based on Planck’s and Einstein’s concepts. De Broglie reasoned that if light could have particle-like properties, then particles like electrons could also have wavelike properties. Why are only certain orbits allowed in the Bohr model? Following De Broglie’s idea, if the electron going around the nucleus in a circular orbit behaves as a wave, then it should behave as a standing wave as shown in Figure 1. In a standing wave, there are fixed points, or nodes, where the amplitude is zero. The length of the wave must fit the circumference of orbit (see Figure 2). Otherwise the wave would cancel itself.
The derivation of the De Broglie equation is only for clarification for the teacher and need not be included in the lesson. It is important that the student understand the concept more than knowing the derivation.
E%=%mc2% E%=%hv% mc2%=%hv% How did de Broglie arrive at his hypothesis? He combined the energy relationship of Einstein’s relativistic equation and Planck’s energy of a photon.
p%=%mc% The momentum, p, of a photon is the product of the relativistic mass of the photon, m, and the speed of light, c, or
p%=% (b) (a)
Standing Waves
(c)
For an allowed orbit, the circumference of the orbit must equal to integral multiple of wavelengths (b). Otherwise, the wave will cancel itself (c).
Mathematically, this means that the circumference of the allowed orbit (2r) must be equal to an integral multiple of the wavelength.
where n = 1, 2, 3…. Because n is an integer, the radius, r, can only have certain values corresponding to n. Therefore, only certain orbits with allowed r values are permissible.
152
hv% c
Substituting the relationship c = λv, the momentum expression becomes
p"="
h% λ
For a material particle, such as the electron, de Broglie substituted for the momentum its equivalent, the product of the mass of the particle, m, and its velocity, u. Therefore,
λ%de%Broglie%=
h% mu
How are the particle and wave properties related according to De Broglie? This is given by the De Broglie equation:
Where h is Planck’s constant, m is the mass of the particle, and u is the velocity. Therefore, a particle in motion can be treated as a wave and a wave can exhibit properties of a particle. An electron, for instance, has both particle and wavelike properties. This is referred to as the dual nature of matter. IV. CALCULATING THE DE BROGLIE WAVELENGTH 1. What is the relationship between the De Broglie wavelength and the mass of the moving particle? What happens to the wavelength as the mass increases? 2. Without doing any calculations, compare the wavelength associated with a moving airplane and an electron moving at the same speed. Which will have the smaller De Broglie wavelength? 3. How will the wavelength vary if the velocity of the particles increases? 4. Calculate the wavelength of the following “particles”: a. A 6.00 x 10-2 kg tennis ball travelling at 68 m/s. λ λ b. An electron moving at the same speed (mass of electron is 9.1094 x 10-31 kg) λ λ
153
Answer Key 1. λ and mass are inversely proportional. As the mass of the particle increases, λ becomes smaller. 2. The moving airplane will have the smaller wavelength.
The wavelength of the tennis ball is exceedingly small considering that the size of the atom is in on the order of 1 x 10-10 m. This makes it difficult for a tennis ball to be detected by any existing measuring device. Meanwhile, the wavelength of the electron is in the infrared region. This shows that only small particles like the electrons and other submicroscopic particles have measurable wavelengths. 5. What must be the velocity, in m/s, of a beam of electrons if they are to display a de Broglie wavelength of 1µm? 6. 2. What is the de Broglie wavelength, in nm, of a 2.4 g bird flying at 1.20 x 102 mph? (1 mile = 1.61 km) 7. What is the wavelength, in nm, associated with 1000 kg automobile travelling at a speed of 25 m/s. Comment on the experimental measurement of the wavelength associated with the moving automobile. V. EXPERIMENTAL EVIDENCE OF DE BROGLIE WAVELENGTH Waves associated with material particles were called by de Broglie as “matter waves”. If matter waves exist for small particles, then beams of particles, such as electrons, should exhibit the properties of waves, like diffraction. Diffraction refers to various phenomena which occur when a wave encounters an obstacle or a slit. In classical physics, the diffraction phenomenon is described as the interference of waves. If the distance between objects that the waves scatter from is about the same as the wavelength of the radiation, diffraction occurs and an interference pattern occurs. Although De Broglie was credited for his hypothesis, he had no actual experimental evidence for his conjecture. In 1927, Clinton J. Davisson and Lester H. Germer, from the United States, shot electron particles onto a crystal of nickel. What they saw was the diffraction of the electron similar to waves diffraction against crystals (x-rays). In the same year, an English physicist, George P. Thomson, from Scotland, fired electrons towards thin metal foil providing him with the same results as Davisson and Germer. As a historical note, the father and son demonstrated the waveparticle duality of electrons. George P. Thomson is the son of J.J. Thomson, who won the Nobel Prize in 1906 for discovering the electron. The father, J.J. Thomson, showed that the electron is a particle and George P. Thomson, the son, showed that the electron is a wave. 154
Teacher Tip A video on diffraction and applications can be seen at https://www.youtube.com/watch? v=F6dZjuw1KUo (4 minutes) A 4-minute video on the Davisson and Germer experiment can be seen at https:// www.youtube.com/watch?v=Ho7K27B_Uu8.
ENRICHMENT (5 minutes)
Return to the question posted on the board during the introduction. Ask the learners to answer the question based on the preceding lesson. Why do elements emit different colors when heated? Pyrotechnic materials such as flares and fireworks also follow the atomic spectra concepts. Inside a mortar are different chemicals.. These chemicals are ignited through a time fuse, causing the electrons in the chemicals to be excited during the reaction in the atmosphere. As the electrons go down a lower energy level, different colors are emitted from these different chemicals. The red glow is light with the least energy and the violet glow has the most energy. What is the wave-particle duality of matter? Ask the learners to answer this in their own words according to their understanding.
EVALUATION
Put a circle around the letter corresponding to the best answer. 1. Waves are characterized by frequency and wavelength. Frequency A. is the distance between two consecutive peaks or troughs in a wave. B. is the number of cycles or complete oscillations that pass a given point per second. C. the vertical distance from the midline of a wave to the peak or trough. D. has units of J-s. E. has units of cm/s. 2. What is the relationship between energy and wavelength of a photon? A. direct relation
D. inverse relation
B. logarithmic relation
E. quadratic relation
C. cubic relation
155
Answer Key 1. B 2. D 3. E 4. C 5. D 6. C 7. B 8. C 9. D 10. D
3. Which of the following types of electromagnetic radiation will have the least energy? A. gamma rays
D. x-rays
B. visible light
E. radio waves
C. microwaves 4. What is the energy in joules of one photon of microwave radiation with a wavelength 0.122 m? A. 2.70 x 10-43 J
D. 4.07 x 10-10 J
B. 5.43 x 10-33 J
E. 2.46 x 109 J
C. 1.63 x 10-24 J 5. In the Bohr model of the H atom A. The atom is a mass of positive charge with electrons embedded in it. B. The electron energy increases as it gets closer to the nucleus. C. The electron goes around the nucleus in certain allowed circular orbits. D. Energy is absorbed when an electron goes from an orbit of high energy to an orbit of low energy E. C and D 6. Complete this sentence: Atoms emit visible and ultraviolet light __________. A. As electrons jump from lower energy levels to higher levels. B. As the atoms condense from a gas to a liquid. C. As electrons jump from higher energy levels to lower levels. D. As they are heated and the solid melts to form a liquid. E. As the electrons move about the atom within an orbit.
156
7. The line spectrum of hydrogen gives proof of the A. B. C. D. E.
Shape of the orbits of the electron Quantized nature of the H energy levels Uncertainty of the momentum of the electron Continuous emission of energy B and D
8. Calculate the energy, in joules, required to excite a hydrogen atom by causing an electronic transition from the n = 1 to the n = 4 principal energy level. Recall that the energy levels of the H atom are given by En = -2.18 10-18 J(1/n2) A. B. C. D. E.
2.07 10-29 J 2.19 105 J 2.04 10-18 J 3.27 10-17 J 2.25 10-18 J
9. Suppose that a tennis ball, a neutron, an electron, and a pingpong ball are all moving at the same speed. The wavelengths associated with them will be of the order: A. B. C. D. E. 10.
tennis ball > pingpong ball> electron> neutron pingpong ball> tennis ball > electron > neutron neutron > electron > pingpong ball > tennis ball electron > neutron > pingpong ball > tennis ball tennis ball > pingpong ball > neitron > electron
Calculate the wavelength of a neutron that has a velocity of 200. cm/s. (The mass of a neutron = 1.675 10-27 kg.) A. B. C. D. E.
1.98 10-9 m 216 nm 1.8 1050 m 198 nm 5.05 mm 157
General Chemistry 1
120 MINS
Lesson 19: Flame Test (Laboratory) Content Standard
Lesson Outline
The learners demonstrate an understanding of the quantum mechanical description of the atom and its electronic structure
Introduction
Explore Colorful Elements Laboratory Activity; Flame Test
35
Performance Standards
Instruction and Practice
The learners can illustrate the dual nature of an electron.
Enrichment
Post-Lab Activity
20
5
Materials
Learning Competencies At the end of the lesson, the learners: 1. Describe the quantum mechanical model of the atom (STEM_GC11CB-IIab-52) a. Discuss quantum theory b. Discuss the use of quantum theory in explaining the emission spectrum of the hydrogen atom Specific Learning Competencies At the end of the lesson, the learners will be able to: 1. Demonstrate the flame tests for various metal ions. 2. Calculate the energy, wavelength, and frequencies involved in the electron transitions in the hydrogen atom. 3. Relate the emission spectra to common occurrences like fireworks and neon lights.
158
Cream of tartar (potassium hydrogen tartrate), table salt (sodium chloride), moisture absorber (calcium chloride), barium chloride, lithium chloride, copper sulphate or copper chloride, boric acid (sodium tetraborate), distilled water, 50mL beaker or a clean glass container, popsicle sticks, large receptacle for used popsicle sticks, alcohol lamp, lighter/ splinter Resources (1) Chang, Raymond and Goldsby, Kenneth A. (2016). Chemistry (12th ed). New York: McGraw-Hill. (2) Petrucci, Herring, Madura, and Bissonnette (2011). General Chemistry and Modern Applications, 10th Ed. Pearson Canada, Inc. (3) Zumdahl, S.S. and Zumdahl, S.A (2013).Chemistry, 8th ed. Cengage Learning (4) Roque, et al. laboratory Manual in General Chemistry (2008). Philippine Normal University.
INTRODUCTION (5 minutes)
Teacher Tip
When substances such as metal salts are heated to high temperatures, the electrons of the metal ions are excited to higher energy levels. When these electrons return to their ground states, energy is emitted in the form of light. Since each element emits a unique set of wavelengths, the emission spectrum can be used as a tool to identify the elements.
2. Distribute the lab sheets at the start of the lesson.
Explore Colorful Elements
One method of demonstrating the emission spectrum of substances is through a qualitative analysis called the flame test. In this technique, a small amount of substance is heated. The heat of the flame excites the electrons of the metals ions, causing them to emit visible light the color of which is unique to the metal ion.
1. Prepare the classroom or laboratory, the materials, the lab sheets to be used.
3. After the introduction and motivation, explain the procedures of the activity. 4. Explain the safety precautions.
Objectives a. To be able to conduct a flame test for metal ions b. To observe the flame colors emitted by selected metal ions. c. To explain the origin of the flame colors. Teacher Tip
INSTRUCTION and PRACTICE (35 minutes) 1. Safety Precautions
a. Do this activity with teacher supervision. Follow all laboratory instructions as directed by the instructor. b. Wear laboratory gown, goggles and mask. c. Consider all metal salts as harmful materials. Do not taste the chemicals. Avoid skin contact with the chemicals. d. Do not eat or drink while doing the activity. e. Dispose of all materials according to the instructions of your teacher.
159
The activity can be performed individually or in groups. Nevertheless, caution must be observed in handling any material in the lab. Instruct the learners how to behave in the laboratory. Dispose the materials properly.
PROCEDURE
Teacher Tip The teacher should prepare the samples ahead of time. Place the salt samples in different watch glass or paper/plastic plates. Label the samples.
1. Dip the popsicle stick in water. 2. Dip the wet popsicle stick into the solid sample. 3. Heat over the flame. Observe the color change in the flame.
Guide to the flame colors
4. Repeat procedures 1-3 with the other samples. Note: It is advisable to repeat the test to ensure that the right color of the flame is observed. 5. Dispose of used popsicle sticks in a receptacle. 6. Note your observation in the data table given. DATA TABLE Sample Material
Metal Ion
Sodium chloride
sodium
Flame Color
There are many causes for the indicated color not to come out such as contaminants in the material, contaminants in the water or in the popsicle stick. Or the flame may not be hot enough. The teacher should NOT mark as wrong any observation. Encourage the students to be honest with stating the result rather than getting the supposedly “correct” answer. The purpose of the experiment is to record observations and try to explain the observations as well as possible sources of error.
ENRICHMENT (20 minutes) Post-Lab Activity
1. Have the students answer the following questions in their activity sheets. Then discuss the class results for the post-lab activity. Compare results of the different groups. a. Why do you think the chemicals have to be heated in the flame first before the colored light is emitted? 160
Metal Ion
Flame Color
Lithium
Red
Sodium
Yellow
Potassium
Lilac
Calcium
Orange / Yellow-red
Strontium
Red
Barium
Pale green
Copper
Blue green
Note Everyone sees and describes colors differently so students may describe their colors as purple instead of lilac or crimson instead of red.
b. Arrange the group of metals which produced the most easily identifiable colors. Start with those that emitted the most intense color and end with those metals with colors that are least intense. c.
Colorful light emissions are observed in everyday life. Where else have you observed light emissions? Are these light emissions an evidence of excited electrons?
d. Cite at least 2 reasons why the flame test is sometimes inaccurate. 2. Write your conclusion and recommendations for the lab activity.
EVALUATION
Use the following rubric to rate the learner’s performance in the lab activity.
CONTENT Observations (10)
PERFORMANCE
PRESENTATION
Attendance and Lab Attire (10)
Lab Results are presented well (10)
Lab Technique and Observance of Safety Procedures (20) Answers to Questions on the Report (20) Conclusions (5)
Housekeeping (5)
Participation in Oral Discussion / Oral Presentation (10)
Lab Requirements / Materials (5)
Recommendations (5)
161
TOTAL
General Chemistry 1
120 MINS
Lesson 20: Electronic Structure of the Atom Content Standard
Lesson Outline
The learners demonstrate an understanding of the quantum mechanical description of the atom and its electronic structure
Introduction
Review of the Quiz
Motivation
3D Models of the Orbitals
Performance Standard
Instruction and Practice
I. Heisenberg’s Uncertainty Principle II. The Schrondinger Equation III. The Quantum Mechanical Description of the Hydrogen Atom IV. The Quantum Numbers V. The Atomic Orbitals
80
Evaluation
Laboratory Activity
25
The learners can illustrate the distribution of the electrons in an atom. Learning Competencies At the end of the lesson, the learners: 1. Describe the electronic structure of atoms in terms of main energy levels, sublevels and orbitals and relate this to energy (STEM_GC11CB-IIa-b-53)
12 3
2. Use quantum numbers to describe an electron in an atom (STEM_GC11CB-IIa-b-54)
Material Calculator
3. (LAB) Perform exercises on quantum numbers (STEM_GC11CB-IIa-b-55)
Resources (1) Chang, Raymond and Goldsby, Kenneth A. (2016). Chemistry (12th ed). New York: McGraw-Hill. (2) Petrucci, Herring, Madura, and Bissonnette (2011). General Chemistry and Modern Applications, 10th Ed. Pearson Canada, Inc. (3) Zumdahl, S.S. and Zumdahl, S.A (2013).Chemistry, 8th ed. Cengage Learning (4) http://csi.chemie.tu-darmstadt.de/ak/immel/script/redirect.cgi? filename=http://csi.chemie.tu-darmstadt.de/ak/immel/tutorials/ orbitals/hydrogenic.html (5) http://winter.group.shef.ac.uk/orbitron/AOs/6g/
Specific Learning Competencies At the end of the lesson, the learners will be able to: 1. Explain Heisenberg’s Uncertainty Principle 2. Describe how atomic orbitals arise from the Schrodinger equation 3. Relate orbital shapes to electron density distribution 4. Qualitatively sketch the orbital shapes 5. Interpret the information obtained from a set of four quantum numbers 6. Assign the correct set of quantum numbers for an electron
162
INTRODUCTION (12 minutes)
1. Review the quiz given in the last meeting 2. Introduce the following learning objectives using any of the suggested protocol (Verbatim, Own Words, or Read-aloud): 3. At the end of the lesson, I will be able to: a. Explain Heisenberg’s Uncertainty Principle b. Describe how atomic orbitals arise from the Schrodinger equation c. Relate orbital shapes to electron density distribution d. Qualitatively sketch the orbital shapes e. Interpret the information obtained from a set of four quantum numbers f.
Assign the correct set of quantum numbers for an electron
4. Present the keywords for the concepts to be learned: a. Heisenberg’s Uncertainty Principle b. Schrodinger Equation c. Wave function d. Electron probability density e. Atomic orbital f.
Principal quantum number
g. Angular momentum quantum number h. Magnetic quantum number i.
Spin quantum number
j.
Shell
k. Subshell
MOTIVATION (3 minutes)
If available, show 3-dimensional models of the orbitals (s, p, and d) to the students to gain their attention and curiosity. If 3-D models are not available, post large illustrations on the board.
163
INSTRUCTION and PRACTICE (80 minutes) I.
HEISENBERG’S UNCERTAINTY PRINCIPLE
With the discovery that particles like electrons are wavelike (shown by De Broglie, Davisson and Germer, and Thomson), how can the ‘position’ of a wave be specified? How can the precise location of a wave be defined when a wave extends in space? Werner Heisenberg, a German physicist, formulated what is now known as Heisenberg’s Uncertainty Principle which states that “the position of a particle and its momentum cannot be simultaneously measured with arbitrarily high precision.” In other words, it is not possible to measure the exact position and the exact momentum of a particle at the same time. Mathematically, this is stated as
where "x is the uncertainty in position, "p is the uncertainty in momentum, and h is Planck’s constant. To explain this equation, let us answer the following questions 1. What is the meaning of the ≥ (greater than or equal to) sign with respect to the uncertainties? When conducting experiments, especially if conditions are crude, the uncertainties in position and momentum can be large. The product of "x"p can be greater than h/4π. However, even when you want to make very precise measurements, h/4π. The product "x"p can never be smaller h/4π. Hence, there will always be uncertainties even under good conditions. 2. How are "x and "p related? They are inversely related. Remember that the right side of the equation, h/4π, is a constant. If we want to make very precise measurement of the position (meaning "x is very small), then "p becomes large. Conversely, if we want a very small uncertainty in momentum, "p becomes small, but the uncertainty in position ("x) becomes large.
164
3. According to the Bohr model, the electron goes around the nucleus in well-defined orbits, the radius of which can be determined. How can you relate the Bohr model to Heisenberg’s Uncertainty Principle? The Bohr model violates Heisenberg’s Uncertainty Principle. Electrons do not go around the nucleus in well-defined orbits. Otherwise, we will be able to determine the exact position and momentum of the electron in the atom at the same time. A better model is needed to fully describe the atom. 4. An electron is travelling at a speed of 2.05 x 106 m/s. Assuming that the precision (uncertainty) of this value is 1.5%, with what precision can the position of the electron be measured? Uncertainty%in%velocity%=%u%=%(0.015)(%2.05%x%106%m/s)%=%3.1%x%104%m/s%% To compute for the uncertainty in momentum, "p, multiply "u by the mass of the electron "p%=%m(ru)%=%(9.109%x%10*31%kg)%(3.1%x%104%m/s)%=%2.8%x%10*26%kg*m/s%% The uncertainty in position, "x, will be "x%=
This value shows that the electron’s position is about 10 atomic diameters. Given the uncertainty of the speed, there is no way to pin down the electron’s position with any greater accuracy.
165
Teacher Tip
5. Why is the uncertainty principle not significant when applied to large objects such as a transportation vehicle? II. THE SCHRODINGER EQUATION While the Bohr model of the atom could explain the emission spectrum of hydrogen, it could not account for many observations and could not provide a complete description of the electronic behavior in atoms. In 1926, Erwin Schrodinger, an Austrian physicist, formulated a mathematical equation that describes the behavior and energies of submicroscopic particles. The Schrodinger equation incorporates particle behavior and wave behavior, treating the electron as a standing wave. The solution to the Schrodinger equation is a wave function called ψ (psi). The wave functions are also called atomic orbitals (as distinguished from the Bohr orbits). Aside from the wave functions, energies are also obtained from solving the equation. The wave function itself has no physical meaning. However, the probability of finding the electron in a particular volume element in space is proportional to ψ2. In wave theory, the intensity of light is proportional to the square of the amplitude of the wave or ψ2. Similarly, the most likely place to find the particle is where the value of ψ2 is greatest. The Schrodinger equation began a new field in physics and chemistry referred to as quantum mechanics or wave mechanics. The Schrodinger equation can be solved exactly for the hydrogen atom but not for atoms with more than one electron. For many-electron atoms, approximation methods are used to solve the Schrodinger equation. III. THE QUANTUM MECHANICAL DESCRIPTION OF THE HYDROGEN ATOM It is not possible to pinpoint the exact location of the electron in an atom but ψ2 gives the region where it can most probably be found. The electron density gives the probability that the electron will be found in a particular region of an atom. Figure (a) is a representation of the electron density distribution around the nucleus in the hydrogen atom. The darker the shade, the 166
The solution of the Schrodinger equation involves advance calculus and differential equations. The lesson will only deal with the interpretation of the solution. The Schrodinger equation for the hydrogen atom looks like this:
higher the probability of finding the electron in that region. In this case, the probability distribution is spherical. The probability can also be plotted versus the distance from the nucleus as shown in Figure (b). It can be seen that there is a probability of finding the electron even very far from the nucleus, although this probability is small. The closer to the nucleus, the higher the probability. Sources (a) and (b) Probability of Finding the Electron in the Ground State of the Hydrogen Atom at Different Points in Space, “Atomic Orbitals and Their Energies”, section 6.5 from the book Principles of General Chemistry (v. 1.0), Retrieved from http://2012books.lardbucket.org/books/ principles-of-general-chemistry-v1.0/s10-05atomic-orbitals-and-their-ener.html (2 Nov. 2016), Creative Commons by-nc-sa 3.0 license.
(a)
(b)
As mentioned earlier, ψ is the solution to the Schrodinger equation. It is also referred to as an atomic orbital. When we say that the electron is in an atomic orbital, we mean that it is described by a wave function, ψ, and that the probability of locating the electron is given by the square of the wave function associated with that orbital. Therefore, the atomic orbital has a characteristic energy as well as a characteristic electron density distribution. This electron density distribution in three-dimensions gives the shape of the atomic orbital.
167
IV. THE QUANTUM NUMBERS In the mathematical solution of the Schrodinger equation, three quantum numbers are obtained. These are the principal quantum number (n), the angular quantum number, (ℓ) ,and the magnetic quantum number (ml). They describe the atomic orbitals. A fourth quantum number, the spin quantum number (ms) completes the description of the electrons in the atoms. The Principal Quantum Number (n) a. Determines the energy of an orbital b. Determines the orbital size c. Is related to the average distance of the electron from the nucleus in a particular orbital; the larger the n value, the farther the average distance of the electron from the nucleus d. Can have the values: n = 1, 2, 3, … e. Orbitals with the same n are said to be in the same shell. The Angular Momentum Quantum Number (ℓ) a. Describes the “shape” of the orbitals b. Can have the following values: ℓ = 0, 1, 2, up to n-1. Examples n value
ℓ value
1
0
2
0, 1
3
0, 1, 2
c. Orbitals with the same n and values belong to the same subshell. d. It is usually designated by letters s, p, d, f, … which have a historical origin from spectral lines. The designations are as follows Teacher Tip The s, p, d, f designations of the orbitals refer to sharp, principal, diffuse, and fundamental lines in emission spectra.
!!!s!!!!!!!!!!!p!!!!!!!!!!d!!!!!!!!!!f!!!!!!!!!!!!g!!!!!!!!!!!h 168
The Magnetic Quantum Number (ml) a. Describes the orientation of the orbital in space b. Can have the values: - ℓ, (-ℓ + 1), … 0, … (+ ℓ -1), + ℓ The Electron Spin Quantum Number (ms) a. The first three quantum numbers describe the energy, shape and orientation of orbitals. The 4th quantum number refers to two different spin orientations of electrons in a specified orbital. b. When lines of the hydrogen spectrum are examined at very high resolution, they are found to be closely spaced doublets and called as the Zeeman effect. This splitting is called fine structure, and was one of the first experimental evidences for electron spin. The direct observation of the electron's intrinsic angular momentum was achieved in the Stern–Gerlach experiment. c. Uhlenbeck, Goudsmit, and Kronig (1925) introduced the idea of the self-rotation of the electron. The spin orientations are called "spin-up" or "spin-down" and is assigned the number ms = ½ ms = -½, respectively. d. The spin property of an electron would give rise to magnetic moment, which was a requisite for the fourth quantum number. The electrons are paired such that one spins upward and one downward, neutralizing the effect of their spin on the action of the atom as a whole. But in the valence shell of atoms where there is a single electron whose spin remains unbalanced, the unbalanced spin creates spin magnetic moment, making the electron act like a very small magnet. As the atoms pass through the in-homogeneous magnetic field, the force moment in the magnetic field influences the electron's dipole until its position matches the direction of the stronger field. The four quantum numbers compose the numbers that describe the electron in an atom. The quantum numbers shall be in the order: energy level (n), sub-level or orbital type (ℓ), the orientation of the orbital specified in ℓ (mℓ), and the orientation of the spin of the electron (ms). It is written in the order (n, ℓ, mℓ, ms ).
169
For example 1. An electron is found in the first energy level. What is the allowed set of quantum numbers for this electron? a. The energy level, n = 1. b. The orbital type is only s, its designation is 0, thus, ℓ = 0 c. From ℓ, the orbital type is s. There is only one orientation of an s orbital, designated as 0, thus, mℓ = 0.m d. An electron in the 1s orbital can have an up-spin or a down-spin. Therefore, ms could be +1/2 or -1/2. So the allowed set of quantum numbers for 1s electron are: (1,0,0,1/2) and (1,0,0,-1/2) How does (1,0,0,1/2) differ from (1,0,0,-1/2)? The first set corresponds to the electron with spin up and the second set refers to the electron with spin down. V. THE QUANTUM NUMBERS AND THE CORRESPONDING ATOMIC ORBITALS The quantum numbers and corresponding atomic orbitals are given in the following table. n
ℓ
mℓ
Number of Orbitals
Atomic Orbital Designation
1
0
0
1
1s
2
0
0
1
2s
2
1
-1, 0, 1
3
2px, 2py 2pz
3
0
0
1
3s
3
1
-1, 0, 1
3
3px, 3py 3pz
3
2
-2, -1, 0, 1, 2
5
3dxy, 3dyz, 3dxz, 3dx2-y2, 3dz2
170
Exercises 1. What is the total number of orbitals associated with the principal quantum number n=1? Answer: 1 What is the total number of orbitals associated with the principal quantum number n=2? Answer: 4 What is the total number of orbitals associated with the principal quantum number n=3? Answer: 9 We can therefore say that the total number of orbitals associated with a given principal quantum number n is n2. 2. List the values of n, ℓ , mℓ for an orbital in the 4d subshell. Answer: n=4; ℓ =2; ml can have the values of -2, -1, 0, 1, 2 The Representations of the Shapes of Atomic Orbitals What are the shapes of the atomic orbitals? Strictly speaking, an orbital does not have a definite shape because the wave function extends to infinity. However, while the electron can be found anywhere, there are regions where the probability of finding it is much higher. Figure (a) shows the electron density distribution of a 1s electron around the nucleus. Note that it does not have a well-defined boundary; the more dots, the darker the shade, the higher the probability of finding the electron in that region. Also note that the probability distribution is spherical. We can draw a boundary surface that will enclose 90% of the total electron density in the orbital as shown in Figure (c). This will result in a boundary surface diagram of the 1s orbital as shown in Figure (d).
Sources (c) Circular boundary enclosing 90 percent of electron density in a hydrogen atom 1s orbital. From Electron Waves in the Hydrogen Atom, Chemistry LibreTexts, National Science Foundation. Retrieved from http:// chem.libretexts.org/Textbook_Maps/ General_Chemistry_Textbook_Maps/Map %3A_ChemPRIME_(Moore_et_al.)/ 05The_Electronic_Structure_of_Atoms/ 5.06%3A_Electron_Waves_in_the_Hydrogen_Atom (3 November 2016), Creative Commons Attribution-Noncommercial-Share Alike 3.0 United States License. (d) The 1s, 2s, and 3s orbitals. From High School Chemistry/Shapes of Atomic Orbitals. Retrieved from https://en.wikibooks.org/wiki/ High_School_Chemistry/ Shapes_of_Atomic_Orbitals (3 November 2016), Creative Commons Attribution-ShareAlike 3.0 License.
(c)
(d) 171
Figure (d) shows that all the s orbitals are spherical in shape but differ in size, which increases as the value of n increases. The p orbitals starts when n =2 for which ℓ has a value of 1 and mℓ has values -1, 0, +1. Therefore, there are three 2p orbitals: 2px, 2py, 2pz indicating the axes along which they are oriented. For the p orbitals, the electron probability density is not spherically symmetric but has a double teardrop shape, or in some books, a dumbbell shape. The greatest probability of finding the electron is within the two lobes of the dumbbell region; it has zero probability along the nodal planes found in the axes. All three 2p orbitals are identical in shape and energy but differ in orientation as shown in Figure (e). The p orbitals of higher principal quantum numbers have similar shapes.
(e) Figure (f) shows the d orbitals occur for the first time when n = 3. The angular function in these cases possesses two angular (or planar) nodes. Four of the orbitals have the same basic shapes except for the orientation with respect to the axes. The wave functions exhibit positive and negative lobes along the axes and shows zero probability of finding the electron at the origin. The fifth wave function, dx2 , has a similar shape with that of the p-orbital with a donut-shape region along the x-axis.
Sources (e) The boundary surface diagrams of the 2p orbitals. From Atomic Orbitals and Their Energies. Retrieved from http://2012books.lardbucket.org/books/ principles-of-general-chemistry-v1.0/s10-05-atomicorbitals-and-their-ener.html (3 November 2016), ), Creative Commons by-nc-sa 3.0 license. (f) The five 3d orbitals of the hydrogen atom. From Atomic Orbitals and Their Energies. Retrieved from http://2012books.lardbucket.org/books/principles-ofgeneral-chemistry-v1.0/s10-05-atomic-orbitals-andtheir-ener.html (3 November 2016), ), Creative Commons by-nc-sa 3.0 license.
(f)
ASSESSMENT/LAB ACTIVITY (25 minutes) QUANTUM NUMBERS Worksheet
Rearrange the letters of the correct term that is described by the corresponding statement. 1. Write your answer on the space provided. ___________
a
LAPNICRIP – It is the quantum number that represents the energy level the electron is in.
___________
b
LATOBRI – It is a representation of the wave function of a hydrogen-like atom.
___________
c
ALGANUR MUTMENMO – It is a quantum number that represents the shape of orbitals.
___________
d
NOTRECLE – It is the particle that can be described by four quantum numbers
___________
e
MEGATINC – It represents the quantum number that describes the orientation of an orbital.
___________
f
NEREGY EVELL – It is being represented by n.
___________
g
RHEPES – It is the shape of the s orbital.
___________
h
ROGUND EATTS – It is the most stable state of the electron in the hydrogen atom.
2. Give the n and ℓ values for the following orbitals _______________________________ a. 1s _______________________________ b. 3p _______________________________ c. 5f _______________________________ d. 4d 3. What is the mℓ values for the following types of orbitals? _______________________________ a. s _______________________________ b. p _______________________________ c. d _______________________________ d. f 173
Answer Key 1. Rearrange the letters a. Principal b. Orbital c. Angular momentum d. Electron e. Magnetic f. Energy level g. Sphere h. Ground state 2. Give the n and ℓ values a. n=1, l = 0 b. n=3, l =1 c. n= 5, l= 3 d. n= 4, l=2 3. mℓ values a. ml= 0 b. ml= -1, 0,1 c. ml= -2, -1, 0, 1, 2 d. ml= -3, -2, -1, 0, 1, 2, 3 4. Possible Orbitals a. 32 electrons b. 50 electrons
4. How many possible orbitals and how many electrons can inhabit the energy level n _______________________________ a. 4 _______________________________ b. 5 5. State the number of possible electrons described by the following quantum numbers a. n = 3, l = 0 _______________________________ b. n = 3, l = 1 _______________________________ c. n = 3, l = 2, ml = -1 ________________________ d. n = 5, l = 0, mℓ =-2, ms =-1/2 _________________ 6. Which of the following is not a valid set of quantum numbers? Explain your answer. a. n = 2, l = 2, ml = 0, and ms = -1/2 b. n = 2, l = 1, ml = -1, and ms = -1/2 c. n = 3, l = 0, ml = 0, and ms = 1 d. n = 3, l = 2, ml = 3, and ms = ½ 7. What is the maximum electron pairs that can occupy an: _______________________________ a. s orbital _______________________________ b. the subshell of p orbitals _______________________________ c. the subshell of d orbitals _______________________________ d. the subshell of f orbitals _______________________________ e. the subshell of g orbitals 8. Do as directed. a. Sketch the shape of the orbital with the quantum numbers n=3, l=0 and mℓ = 0 b. The sketch of the shape of the subshell with the quantum numbers n=4, l=2 is c. The highest orbital possible in n = 4 d. Sketch the orientation of the allowed values of l= 1 for the shell n=2. e. Write the set of quantum numbers for the following i. It is an up-spin 4d electron with an orbital orientation of 0. ii. The electron is in the 3rd energy level, px-orbital, and down spin. iii. When n=2, l is 1, mℓ = 1, ms = ½ f. What is the value of l for a 4f electron? g. What is the orbital designation for an electron in the 3rd shell and p sublevel? h. How many electrons have the following quantum numbers: n =4, l = 2, mℓ = -2?
174
5.
Number of possible electrons a. 2 b. 6 c. 2 d. not possible
6. Valid Set of Quantum Numbers a. l =2 is not allowed, maximum is 1 b. possible c. ms should only be ½ or -1/2 d. mℓ should only be within the values of 2l+1; mℓ should only be within the values of 2l+1 7. Maximum electron pairs a. 1 b. 3 c. 5 d. 7 e. 9 8. “Do as directed” a. Sphere b. Any of the d orientations c. f orbital d. p orbitals e. Answers: i. (4, 2,0, ½) ii. (3,1,-1,-1/2) iii. (2,1,1,1/2) f. 3 g. 3p h. 2
General Chemistry 1
120 MINS
Lesson 21: Electron Configuration Content Standard The learners demonstrate an understanding of the electronic distribution in an atom. Performance Standards The learners can illustrate the distribution of the electrons in an atom. Learning Competencies At the end of the lesson, the learners:
Lesson Outline Introduction
Communicating Learning Objectives
7
Motivation
Addresses and Zip Codes
3
Instruction and Practice
I. II. III. IV. V.
Evaluation
Exercises and Activity
1. Write the electronic configuration of atoms (STEM_GC11CB-IIa-b-56) 2. Determine the magnetic property of the atom based on its electronic structure (STEM_GC11CB-IIa-b-57) 3. Draw an orbital diagram to represent the electronic configuration of atoms; (STEM_GC11CB-IIa-b-58) 4. Perform exercises on writing electronic configuration (STEM_GC11CB-IIab-59) Specific Learning Competencies
Energies of the Orbitals Electron Configuration Hand and Rule Aufbau Principle The Quantum Numbers and the Arrangement of Elements in the Periodic Table
80
30
Materials Periodic Table Resources
(a) Chang, Raymond and Goldsby, Kenneth A. (2016). Chemistry (12th ed). New York: McGraw-Hill.
At the end of the lesson, the learners will be able to:
1. Explain the unique electron distribution of the atom;
(b) Petrucci, Herring, Madura, and Bissonnette (2011). General Chemistry and Modern Applications, 10th Ed. Pearson Canada, Inc.
2. Compare and contrast the orbital energies in a hydrogen atom with that of the many-electron atom; 3. Write the electron configuration of an atom using the conventional method as well as the core noble gas configurations;
(c) Zumdahl, S.S. and Zumdahl, S.A (2013).Chemistry, 8th ed. Cengage Learning
4. Illustrate the electron distribution using orbital diagrams;
6. Determine valence configuration and valence electrons.
5. Determine magnetic properties of an atom based on its electronic configuration; and
7. Relate valence configuration of elements with position of element in the periodic table.
175
INTRODUCTION/ REVIEW (7 minutes) 1. Review orbitals and their shapes.
2. Introduce the following learning objectives using any of the suggested protocol (Verbatim, Own Words, or Read-aloud): At the end of the lesson, I will be able to: a. Explain the unique electron distribution of the atom; b. Compare and contrast the orbital energies in a hydrogen atom with that of the manyelectron atom; c. Write the electron configuration of an atom using the conventional method as well as the core noble gas configurations; d. Illustrate the electron distribution using orbital diagrams; e. Determine magnetic properties of an atom based on its electronic configuration; and f. Determine valence configuration and valence electrons. g. Relate valence configuration of elements with position of element in the periodic table. 3. Present the keywords for the concepts to be learned: a. Ground state b. Excited state c. Degenerate d. Electron configuration e. Orbital Diagram f. Pauli Exclusion Principle g. Paramagnetic h. Diamagnetic i. Hund’s Rule j. Building-Up Principle (Aufbau Principle) k. Noble gas l. Transition metals m. Valence configuration n. Valence electrons
176
MOTIVATION (3 minutes)
1. Ask a few learners to give their home addresses. 2. What are zip codes? What is the zip code of the school? 3. What is the use of zip codes? Look for the zip code of a school outside your city or province and compare with yours.
INSTRUCTION/ DELIVERY/ PRACTICE (80 minutes)
I. ENERGIES OF THE ORBITALS After understanding the shapes and sizes of atomic orbitals, it is imperative to understand the relative energies of the orbitals and how it affects the actual arrangement of electrons in atoms. 1. Orbital energy levels in a hydrogen atom The energy of an electron in a hydrogen atom depends solely on its principal quantum number, n. The energy of the electron in the hydrogen atom is given by:
where RH is equal to 2.18 x 10-18J. Therefore, the energies of the hydrogen atom increase according to the following (see Figure 1): 1s"
