Nathalie Dagmang
Group 8
Co-workers: Annjaneth Briones and groups 5, 6, 7 and 9
Date Performed: February 22, 2011
Results and Discussion Report 11: Potentiometric Determination of the Purity and Dissociation Constant of Potassium Hydrogen Phthalate
The three main objectives of the experiment are to (1) apply the principles of potentiometry in the determinatitaon of the equivalence point in a titration, (2) determine the purity of KHP using potentiometric titration, and (3) derive the acid dissociation constant of KHP from the potentiometric data. Aside from the commonly used volumetric analysis, another method used for determining the concentration of a substance is the potentiometry. This is used when visual indicators are unavailable and when one is aiming to get more accurate results. There are two types of potentiometry: direct and indirect. In direct potentiometry, the potentials of a cell with the indicator electrode at different analyte concentrations and a constant potential at reference electrode are Acompared. While in indirect potentiometry, the pH is measured at each addition of titrant. Potentiometric methods are very accurate compared to volumetric titrations because of the high sensitivity of the pH meter as compared to visual indicators, the possibility of human error in determining the endpoint by looking at the color change of the analyte and the effect of the indicator on the pH of the analyte. Also, indirect potentiometry use potential as the function of titrant volume, where it follows the actual change in concentration at different amounts of titrant added. Hence, one can see the pH values of the analyte at many different points, not just the point at which it reaches the equivalence point. Thus, potentiometry is often used for dilute solutions. Because the endpoint is indicated at the largest potential break, the exact potential is unnecessary and only the change in cell potential is needed. Also, the glass electrode does not need to be calibrated with a standard buffer nor does the pH meter need to be a high-end model since potential need not be read closely. However, in order to anticipate the endpoint, so that one will add smaller increments as it begins to neutralize, the approximate equivalence point may be predetermined using calculations, volumetric titrations or a run of potentiometric titration. 1
Aside from the concentration of the sample, its acid-dissociation constant (Ka) and percent purity was also determined. Three trials were done using a pH meter set-up as described in Figure 1.
Figure 1. Set-up for Potentiometric Titration
The pH meter’s electrode is immersed in the beaker containing the analyte while it is titrated with a NaOH solution. The magnetic stirrer was used so that the solution is constantly mixed and the NaOH is equally distributed throughout the analyte, Potassium Hydrogen Phthalate (KHP). At each addition of titrant, the pH was recorded and was used to determine the equivalence point, acid-dissociation constant and percent purity. For the first run, the solution was added 50 ml of titrant. For each pH recording, 1 ml of titrant was added. The approximate endpoint determined in the first run was used in the following runs so that the complete neutralization can be anticipated and smaller increments can be added as it reaches the endpoint. However, the volume increments should not be too small as there may be two or more points of the straight line of the second derivative plot that passes zero and it may become vulnerable to experimental errors due to more volume measurements. Also, the volume increments should not be too large or there would not be enough points to make accurate results and graphs. The data gathered was then plotted so that one can clearly see the equivalence point. The first plotted graph was the normal titration curve. Shown in Figure 2 is the normal titration curve for trial 2.
Figure 2. Normal Titration Curve
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For this graph, the equivalence point can be seen at the corresponding volume of the curve’s inflection point where the pH change is highest. The preceding points are the pre-equivalence points where the pH values are acidic while the points following the inflection point are the post-equivalence points where the pH of the analyte is towards basic. Another graph used was the First Derivative plot where dpH/dV is a function of the average volume. Figure 3 shows the First derivative plot of trial 2.
Figure 3. First Derivative Plot
This graph gives a clearer image for determining the equivalence point. Because the change in potential with respect to the added volume is highest at equivalence point, the graph will have a “spiked” curve and will peak at the corresponding volume. The pre- and post-equivalence points have almost equal changes in potential as certain volumes of titrant are added. Lastly, a Second Derivative Plot was graphed, which is showed in Figure 4. For this graph, one can see the endpoint at which y is equal to zero.
Figure 4. Second Derivative Plot
The KHP analyzed in the experiment follows the dissociation reaction: KHP↔KP-+H+ In order to calculate the acid-dissociation constant of KHP, the halfequivalence point was determined. At half-equivalence point, where the volume of titrant used is half of the volume used at equivalence point, the number of moles of the KP- is equal to the number of moles of KHP left. Therefore, given the pH at halfequivalence point and using the formula for the acid-dissociation constant, the Ka can be calculated. Ka=KP-[H+][KHP] At ½ equivalence point,
KP-=[KHP] Ka=KP-[H+][KHP] Ka=[H+]→pKa=pH The percent purity, on the other hand was calculated by using the formula:
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%KHP=MNaOH×VNaOH×1 mol KHP1 mol NaOH×204.2 gmol KHPg KHP sample×100% It was found out that the sample’s percent purity is 69.24% with a confidence interval of + 0.3 and a pKa of 4.87 with a confidence interval of 4.87 + 0.01. When compared to the book value of KHP’s pKa which is 5.51 at room temperature, the calculated percent error was 11.62%. The two trials yielded slightly different endpoints, hence, also different halfequivalence points and volume used in calculating the acid-dissociation constant. Also, because several very small increments of titrant were added at each point, there were errors due to volume measurements.
Sources: Skoog, et al., Fundamentals of Analytical Chemistry, Eighth edition, 2004 Day, Underwood, et al. Quantitative Analysis, 1967 Christian, G.D. Analytical Chemistry, 1986
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