Ppotentiometric Titration of Benzoic Acid With 0.1M Sodium Hydroxide [PDF]

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Western Mindanao State University Normal Rd. Baliwasan, Zamboanga City College of Science and Mathematics



EXPERIMENT NO. 8 POTENTIOMETRIC TITRATION OF BENZOIC ACID WITH 0.1M SODIUM HYDROXIDE



Cristine D.Z.Concepcion 2018-02771 BS CHEMISTRY II



December 11, 2019



ABSTRACT



Potentiometric titration is a technique useful in characterizing acids. It is the preferred approach for assessing acidity and stability constants in chemistry and is similar to direct titration of a redox reaction. This experiment aims to determine the equivalence point for the titration of 20mL benzoic acidwith 0.1M sodium hydroxide. Standardizatiojnof NaOH was done before titrating to be used on the potentiometric titration. There were four trials of KHP as the analyte. and with these. the experimental cocnentration value was determined, which is 0.090019487M. This concentration was used in the potentiometric titration of the mentioned titrant and analyte. Two trials were established, one for each group. The standard deviation is lower than 1.00%, or is too small - means that the measurements of the analyte are precise. A titration curve is established with an equivalence point result to approximately 15.70mL. After which, the data from the titration curve was used for the first derivative plot which resulted to an equivalence point of approximetely 15.70mL and with the second derivative plot which is 15.80mL. The resulting equivalence point was approximately between 15.70mL to 15.80mL, or 15.75mL.



CHAPTER I INTRODUCTION



1.1 Background of the Study Titrimetry is one of the oldest analytical methods. It showcases an important role in different fields regarding in analytical chemistry aspect. Titrimetry is known for having its superior speed, its precision in simplicity, and a little sacrifice of accuracy, giving justice as to why it is one of the most significant methods used in analytical chemistry. The potentiometric method, with a glass electrode, is most likely used in some analytical laboratories in quantitatively determining substances with acid-base properties. Potentiometric titration is the preferred approach for assessing acidity and stability constants in chemistry and biochemistry in many different situations. In potentiometric titrations, the equivalence point is commonly determined by calculating and locating the point of maximum slope, or the inflection point, of the titration curve. In some cases, very sharp breaks in the titration curve are bound to be the basis, making it at ease in locating the equivalence point. Nevertheless, when very weak acids are titrated with strong bases, the curves are more difficult to evaluate. It is then a necessity to locate the equivalent point by plotting the differential curves, ΔpH/ΔV or ΔE/ΔV against volume of titrant added. The peak shown on the curves correlates to the point of maximum slope of the normal titration curve. Such differential methods need values of potential corresponding to very small change in volume of titrant added near the end point for good result. A typical cell for measuring the pH of a certain sample comprise of a glass indicator electrode and a saturated calomel reference electrode that are consumed in the solution whose pH



is unknown. The indicator electrode contains a thin, pH sensitive glass membrane that is sealed into one end of a heavy-walled glass or a plastic tube. A little portion of hydrochloric acid that saturated with silver chloride, is contained in the tube. A silver wire in this solution forms a silver/silver chloride inner-reference electrode, which is connected to one of the terminals of the potential-measuring device, pH-meter.



Figure 1.1. Combination glass electrode (http://users.metu.edu.tr/chem223/potentiometry.pdf)



1.2 Objectives of the study At the end of this experiment, the students can: i.



standardize a strong base to be used in the potentiometric titration analysis



ii.



titrate a weak acid with a strong base



iii.



explain the principles and concepts of potentiometric titration



iv.



determine precent error and standard deviation of the analyte



v.



establish a titration curve, first derivative plot curve, and second derivative plot curve for a better determination of equivalence point



vi.



determine equivalence point of the potentiometric titration of benzoic acid with 0.1M sodium hydroxide



1.3 Significance of the Study Potentiometric titration is applied in measuring the change in the electrical potential when a neutralizing agent or the titrant is added to a chemical solution. Potentiometric titration can also be used as a technique in characterizing acid/base and complexing properties of hydrous particle surfaces. This method provides an accurate data over wide concentration ranges and, with help of automatic measuring and data collecting systems, it has become possible to collect large amounts of data within a reasonably short time period.



CHAPTER II METHODOLOGY This chapter shows the system of methods done in the activity and provides an outline for the research methodology.



2.1 Standardization of strong base (Sodium hydroxide) 2.1.1 Preparation of sodium hydroxide solution A burette was properly washed and cleaned by rinsing with several portions of distilled water. This rinse water was poured down the drain. An amount of 4.00g of sodium hydroxide was obtained and was used from standardizing the strong base. This amount was dissolved with about 50mL of carbon-free distilled water. A 1000mL volumetric flask was procured in the experiment. The dissolved sodium hydroxide was poured into the flask, and a distilled carbon-free water was then drained into the flask until it reached the certain line in the flask. The burette was rinsed with three portions of the NaOH solution. The rinsed portion was drained and discarded. The burette was then filled with NaOH slightly the zero mark and the burette was clamped up vertically. The air bubbles from the tip of the burette were removed by draining the NaOH into a small beaker. The NaOH level value was read and recorded within + 0.02 mL.



2.1.2 Preparation of KHP analyte solution A clean 250mL beaker was obtained. An amount of 0.24g of a dried acidic salt compound, potassium hydrogen phthalate or KHP, was poured into the beaker. The KHP was then dissolved with a 20mL carbon-free distilled water by constantly stirring the



solution with a glass rod. The solution was then transferred into a clean 250mL Erlenmeyer flask.



2.1.3 General standardization procedure



Two to three drops of phenolphthalein indicator was drained to the KHP solution in the Erlenmeyer flask. The flask containing the acid solution and indicator was placed under the burette set-up. NaOH from the burette was slowly added to the flask with swirling until the color of the solution in the flask is a faint pink. This faint pink color should last only 45 to 60 seconds. There should be a one-drop difference between when the solution is colorless and when it is pink. If too much base is added discard the solution and repeat the titration. A white piece of paper placed under the flask will aid in the color detection. The final volume of the NaOH was read and recorded. The molarity was then calculated from the titration trials.



2.2 Preparation of the sample analyte (benzoic acid) An amount of 0.18g of benzoic acid was obtained in a clean 250mL beaker. The benzoic acid solids were poured into a 20mL carbon-free distilled water. The solution was mixed thoroughly. The solution was then transferred into a 250mL Erlenmeyer flask, ready for titration. 2.3 General Titration Procedure 2.3.1 Preparation of potentiometric titration set-up AEutech pH meter was used in the analysis. The pH meter was readily calibrated. The iron stand with the burette that has the strong base was placed beside the pH meter.



2.3.2 Potentiometric titration The Erlenmeyer with the analyte was properly placed under the burette. The pH electrode was thoroughly mixed between measurements with distilled water to prevent carryover contamination of the tested solutions. The electrode was gently blotted on a laboratory cleaning tissue to remove the excess rinse water. The pH electrode was dipped into the testing solution or suspension. A magnetic stirring bar was placed in the analyte. The NaOH was dispensed slowly but surely per 0.50mL and 1.00mL. The titration process was occurred, done with one trial each per group.



CHAPTER III RESULTS AND DISCUSSION This chapter discusses the data and result gathered from doing the activity. 3.1. Standardization of 0.1M NaOH Table 3.1 Summary data table for the standardization of NaOH Trial



Mass of KHP, g



Volume of NaOH dispensed, mL



Moles of KHP, mol



Concentration, M



1



0.24g



13.75mL



1.17514x10-3



0.087272727



2



0.25g



13.60mL



1.22413x10-3



0.088235294



3



0.24g



13.00mL



1.17514x10-3



0.090395520



4



0.25g



13.00mL



1.22413x10-3



0.090019487



Average



0.090019487



Standard Deviation



±0.002658



The Table above shows the summary data for the standardization procedure of sodium hydroxide. Four trials were done in the process, of which two are of the same mass of KHP. Each trial has different volume of sodium hydroxide that has dispensed. In the first trial, the mass of KHP used is 0.24g, with a volume dispensed by the sodium hydroxide to be 13.75mL. By the aid of the calculations, the resulting mole of the KHP is 1.17514x10 -3, with a concentration of 0.087272727M. For the second trial, 0.25g of KHP was used. The volume dispensed is 13.60mL. With calculations, the mole of the KHP in the second trial is 1.22413x10 -3 with a concentration of 0.088235294M. For the third and fourth trials, the mole and concentration of the KHP are the



same as of the first and second respectively. The average of the concentrations of the four trials is 0.090019487, which is the final concentration of the NaOH to be used in the potentiometric titration of benzoic acid. 3.2 Standard deviation and percent error of the concentration of the analyte Table 3.2 Summary table for the standard deviation and percent error of the concentration of the analyte Experimental Concentration of NaOH



0.090019487



Theoretical Concentration of NaOH



0.1000



Standard Deviation



0.002255



Percent Error



0.09980



The data in Table 3.2 shows the experimental and theoretical concentrations of the titrant. The calculated value for concentration of sodium hydroxide was determined by standardizing the NaOH with KHP in four trials. After getting the concentrations from these trials, the average was then determined, resulting to an experimental value of 0.090019487M. Moreover, the theoretical value is 0.1000M. The standard deviation of the analyte tells how close the values are. It is a statistical measurement of the precision of a group. The smaller the s is, the more precise the measurement is. The computed value for the standard deviation of the analyte is 0.002255 or 0.02%. The standard deviation is lower than 1.00%, or is too small - means that the measurements of the analyte are precise. The repeated values are close with each other. The percent error is calculated by subtracting the experimental value to the theoretical value, and the difference was divided by the theoretical value, multiplied by 100. The percent error is 0.09980 or 9.98%. The accepted



value of percent error is 5%, or lower, but in this case the percent error is 9.98% - slightly acceptable. 3.3 Quantitative Analysis by potentiometric titration of benzoic acid with 0.1M sodium hydroxide The potentiometric titration was done with two trials, each of different groups. Phenolphthalein was used as the indicator. Using Excel, data were plotted for each titration. Two groups titrated individually, each had one set list of pH corresponding to how many volume the titrant had dispensed. The values obtained were used to calculate for the average pH from the two trials.



TITRATION CURVE OF BENZOIC ACID WITH 0.1M SODIUM HYDROXIDE



14.00 12.00 10.00



pH



8.00 6.00 4.00 2.00 0.00 0.00



2.00



4.00



6.00



8.00



10.00



12.00



14.00



16.00



18.00



20.00



22.00



24.00



26.00



NaOH volume



Figure 3.1 titration curve of 20mL benzoic acid with 0.1m sodium hydroxide



The figure above shows the titration curve of 20mL benzoic acid with 0.1m sodium hydroxide. On the x-axis lies the volume of the titrant dispensed while the y-axis shows the H reading. The highest slope in the titration curve is characterized as an inflection point and that it



is regarded to be the equivalence point. Thus an inflection point has to be present in the titration curve in order to locate the equivalence point. From the data shown above, it is safe to assume that the equivalence point is located approximately at 15.75mL.



14.00



TITRATION CURVE OF BENZOIC ACID WITH 0.1M SODIUM HYDROXIDE



12.00



pH



10.00 8.00 6.00 4.00 2.00 0.00 0.00



2.00



4.00



6.00



8.00



10.00



12.00



14.00



16.00



18.00



20.00



22.00



24.00



26.00



NaOH volume



Figure 3.2 titration curve of 20mL benzoic acid with 0.1m sodium hydroxide with corresponding points



The titration curve shows different points located on the curve. The red point on the titration curve indicated the point where no NaOHwas added. So, the pH of the analyte is low. The chemical equation is denoted as: C6H5COOH + H2O ↔ C6H5COO- + H3O+. As the sodium hydroxide is added in a drop-wise motion, the H3O+ is slowly starting to be consumed by the OH- that is produced by the dissociation of the NaOH. Moreover, the analyte is still acidic due to the predominance of the H3O+ ions. The green point in the titration curve represents the pH recorded at a point in time just before the completion of neutralization is taking place. Lastly, the yellow point in the curve is the point where the NaOH is in excess, the point beyond the equivalence point.The point of maximum slope, or the inflection point, is where the equivalence



point is located. At this point, the moles of the titrant and the analyte are the same. The H 3O+ ions are now completely neutralized by the OH - ions. The solution only contains C6H5COONa salt and H2O. This reaction is denoted as: C6H5COOH + NaOH ↔ C6H5COONa + H2O. From this reaction, the solution contains C6H5COONa salt. This salt dissociates intoC6H5COO- and Na+. According to the principle of weak acid/strong base titration, C6H5COO- is the conjugate base of C6H5COOH. So C6H5COO- is the strong base that will react with H2O to produce hydroxide ions OH-, thus there is an increase in pH. 3.3.1 Estimation of pK a at half-titration point The equivalence point is located to be at approximately 15.75mL with the help of the first titration curve. The estimated acid dissociation constant of the acid can now be calculated, this is called the buffer point, or the half-equivalence point. Generally, the equilibrium for a weak acid is denoted as: HA (aq) + H2O ↔ H3O+ + A-. The half-titration point is the conversion of half of the acid into its conjugate base.



Half-titration curve



4.80 4.60 4.40



pH



4.20 4.00 3.80 3.60 3.40 3.20 3.00 0.00



0.80



1.60



2.40



3.20



4.00



4.80



5.60



6.40



7.20



Volume of NaOH



8.00



8.80



9.60 10.40 11.20 12.00



Figure 3.3. Half-titration curve of 20mL benzoic acid with 0.1m sodium hydroxide



The data from Figure 3.3 shows the half-titration curve of 20mL benzoic acid with 0.1m sodium hydroxide. The equivalence point from the previous determination was determined, occurring at 15.75mL. The half equivalence of this point is 7.88mL with a pKaof approximately 4.10. 3.3.2 First derivative plot The first derivative plot is an aid to obtain a better and more improved approximation of the volume at the equivalence point. The plot in the first derivative graph is the change of pH divided by the change in volume versus the volume of NaOH.



8



7.52



7



6 ΔpH/ΔV



5 4 3 2



1.5



1 0 13.00



0.26 13.40



13.80



14.20



0.32 14.60



1.34



0.54 15.00



0.45 15.40



15.80



16.20



16.60



17.00



0.3 17.40



0.2 17.80



18.20



Average Volume of NaOH



Figure 3.4 First derivative plot



The data from Figure 3.4 shows first derivative plot, where the slope of the titration curve changes while approaching the equivalence point. The first derivative shows a spike at the equivalence point, approximately 15.70mL, due to slope that had reached its maximum value at the inflection point.



3.3.3 Second derivative plot Other than having a more improved approximation of the volume at the equivalence point with the first derivative, the second derivative can clearly depict the volume as it intersect directly to the x-axis. Moreover, the second derivative curve gives the end point more precisely.



Second Derivative 15



Δ(ΔpH/ΔV)/ΔV



10 5 0 14.00 14.30 14.60 14.90 15.20 15.50 15.80 16.10 16.40 16.70 17.00 17.30 17.60 -5 -10



-15



Average Volume of NaOH



Figure 3.5 Second derivative plot The data from the Figure 3.5 shows the second derivative plot for this titration around the inflection point. It can be asserted that the inflection point of the original function is located nearly 15.80mL.



CHAPTER IV SUMMARY AND CONCLUSION



This chapter shows the summary and conclusion of the study. This aims to cover-up the end result of the study. The experiment seeks to determine the equivalence point of the potentiometric titration of 20mL benzoic acid with 0.1M sodium hydroxide. The standardization of sodium hydorxide was done first. The analyte used was the potassium hydorgen phthalate or KHP. Four trials were conducted, of which two of the four have the sae mass of KHP. With these mass, the moles of the KHP were determined, and as well as the four concentrations of NaOH. The average of hese four concentrations were calculated as the experimental concenrtation of the titrant, which is 0.090019487M, that woud be used in the potentiometric titration. The data were gathered patinetly, as



the reading is recorded 0.5mL, 1mL, and 2mL difference. The NaOH was



dispensed ata certain range, thus having numerouds readings. The standard deviation is lower than 1.00%, or is too small - means that the measurements of the analyte are precise. A titration curve was established from the data gathered. It was found out that the highest slope on the inflection point lies approximately 16.00mL. The data from the first curve was used to geth the first derivative plot curve, which resulted to an equivalence point to approximetely 15.70mL, and with the second derivative plot curve that resulted to an equivalence point to approximetely 15.80mL. From these informations, it can be concluded that the equivalence point of the potentiometric titration of benzoic acid with 0.1m NaOH is between 15.70mL to 15.80mL, or 15.75mL.



REFERENCES



[1] Checchetti A., Lanzo J. (2015) Qualitative Measurement of pH and Mathematical Methods for the Determination of the Equivalence Point in Volumetric Analysis. World Journal of Chemical Education, Vol. 3, No. 3, pp. 64-69.



[2] Martín J., Ruiz, D.B., Asuero, A. (2018) . Determination of the End Point in Potentiometric Titrations: Gran and Schwartz Methods. Journal of Laboratory Chemical Education, 6(4), pp. 77-90. doi:10.5923/j.jlce.20180604.02



[3] Perveen S. &Mohiuddin S. (2016) Multiproticity of Weak Acids: Inflection Point vs. Equivalence Point. World Journal of Chemical Education, Vol. 4, No. 1, pp. 21-24 . DOI:10.12691/wjce-4-1-4



[4]Potentiometric Titration experiment. (online) Retreived https://www.bc.edu/content/dam/bc1/schools/mcas/Chemistry/pdf/undergrad/labcourses/gen/spring/pHtitration.pdf



from:



[5]Potentiometric Titration of a weak acid. Retrieved from: http://fliphtml5.com/qlgo/wesg/basic



[6] Steenbock H. (1912) quantitative determination of benzoic, hippuric, and phenaceturic acids in urine. Laboratory of the Department of Agricultural Chemistry of the University of Wisconsin.



[7] Sjöberg S., &Lövgren Lars., (1993) The application of potentiometric techniques to study complexation reactions at the mineral/water interface. Aquatic sciences, Vol. 55, No.4, pp 324–335



APPENDIX



Raw data Titration of 20mL benzoic acid with 0.1M sodium hydroxide Volume dispensed (mL) 0.00 1.00 2.00 3.00 4.00 5.00 6.00 7.00 8.00 9.00 10.00 11.00 12.00 12.50 13.00 13.50 14.00 14.50 15.00 15.50 16.00 16.50 17.00 17.50 18.00 18.50 19.00 19.50 20.00 22.00 24.00



Trial 1 2.87 3.26 3.54 3.72 3.87 3.99 4.10 4.21 4.32 4.42 4.53 4.67 4.79 4.85 4.93 5.02 5.15 5.29 5.48 5.85 9.64 10.62 10.87 11.04 11.15 11.23 11.30 11.36 11.41 11.56 11.66



pH Trial 2 2.68 2.92 3.17 3.34 3.52 3.65 3.78 3.92 4.01 4.13 4.25 4.41 4.53 4.63 4.71 4.84 4.97 5.15 5.50 6.63 10.36 10.72 10.92 11.05 11.14 11.22 11.28 11.34 11.39 11.52 11.61



Ave. 2.78 3.09 3.36 3.53 3.70 3.82 3.94 4.07 4.17 4.28 4.39 4.54 4.66 4.74 4.82 4.93 5.06 5.22 5.49 6.24 10.00 10.67 10.90 11.05 11.15 11.23 11.29 11.35 11.40 11.54 11.64



Half-titration Volume of NaOH 4.00 5.00 6.00 7.00 8.00 9.00 10.00 11.00



pH 3.70 3.82 3.94 4.07 4.17 4.28 4.39 4.54



First derivative data Volume of NaOH 13.50 14.00 14.50 15.00 15.50 16.00 16.50 17.00 17.50 18.00



ΔpH



pH 4.93 5.06 5.22 5.49 6.24 10.00 10.67 10.90 11.05 11.15



ΔV 0.13 0.16 0.27 0.75 3.76 0.67 0.225 0.15 0.1



Avg. Vol 0.50 0.50 0.50 0.50 0.50 0.50 0.50 0.50 0.50



ΔpH/ΔV



13.75 14.25 14.75 15.25 15.75 16.25 16.75 17.25 17.75



0.26 0.32 0.54 1.5 7.52 1.34 0.45 0.3 0.2



Second derivative data Avg. Vol 13.75 14.25 14.75 15.25 15.75 16.25 16.75 17.25 17.75



ΔpH/ΔV 0.26 0.32 0.54 1.5 7.52 1.34 0.45 0.3 0.2



Δ(ΔV) 0.50 0.50 0.50 0.50 0.50 0.50 0.50 0.50



Δ(ΔpH/ΔV) 0.06 0.22 0.96 6.02 -6.18 -0.89 -0.15 -0.1



Ave. Vol 14.00 14.50 15.00 15.50 16.00 16.50 17.00 17.50



Δ(ΔpH/ΔV)/ΔV 0.12 0.44 1.92 12.04 -12.36 -1.78 -0.3 -0.2



Documentaries



Potentiometric titration set-up



Sample pH meter readings



Sample Calculations



Moles of KHP:



Moles KHP 1 = (0.24g) (



1mol ) = 1.17514x10−3 204.2g



Moles KHP 2 = (0.25g) (



1mol ) = 1.22413x10−3 204.2g



Volume of NaOH:



Volume of NaOH 1 = (13.75mL) (



1L ) = 0.01375L 1000mL



Volume of NaOH 1 = (13.60mL) (



1L ) = 0.01360L 1000mL



NaOH concentration:



NaOH concentration 1 = (



1.17514x10−3 ) = 0.087272727M 0.01375L



Experimental concetration of NaOH:



0.087272727 + 0.088235294 + 0.090395520 + 0.090019487 ) 4 = 0.090019487



M exp 𝑜𝑓 𝑁𝑎𝑂𝐻 = (



First derivative plot: ∆𝒑𝑯



∆𝑝𝐻 = 𝑝𝐻 2 − 𝑝𝐻 1 ∆𝑝𝐻 = 5.06 − 4.93 ∆𝑝𝐻 = 0.13



∆𝑝𝐻 = 𝑝𝐻 3 − 𝑝𝐻 2 ∆𝑝𝐻 = 5.22 − 5.06 ∆𝑝𝐻 = 0.16



ΔV



∆𝑉 = 𝑉 2 − 𝑉 1 ∆𝑉 = 14.00 − 13.50 ∆𝑉 = 0.50



∆𝑉 = 𝑉 3 − 𝑉 2 ∆𝑉 = 14.50 − 14.00 ∆𝑉 = 0.50



Average V



𝑉2 + 𝑉1 2 14.00 + 13.50 𝐴𝑣𝑒. 𝑉 = 2 𝐴𝑣𝑒. 𝑉 = 13.75 𝐴𝑣𝑒. 𝑉 =



𝑉3 + 𝑉2 2 14.50 + 14.00 𝐴𝑣𝑒. 𝑉 = 2 𝐴𝑣𝑒. 𝑉 = 14.25 𝐴𝑣𝑒. 𝑉 =



ΔpH/ΔV



0.13 0.50 ∆𝑝𝐻/𝑉 = 0.26 ∆𝑝𝐻/𝑉 =



0.16 0.50 ∆𝑝𝐻/𝑉 = 0.32 ∆𝑝𝐻/𝑉 =



Second derivative plot: Δ(ΔV)



Δ(ΔV) = 𝐴𝑣𝑒. 𝑉 2 − 𝐴𝑣𝑒. 𝑉 1 Δ(ΔV) = 14.25 − 13.75 Δ(ΔV) = 0.50



Δ (ΔpH/ΔV)



Δ(ΔpH/ΔV) = 0.32 − 0.26 Δ(ΔpH/ΔV) = 0.06



`



Average V



𝑉2 + 𝑉1 2 14.25 + 13.75 𝐴𝑣𝑒. 𝑉 = 2 𝐴𝑣𝑒. 𝑉 = 14.00 𝐴𝑣𝑒. 𝑉 =