Acid-base Equilibria - an Examination

In experiment 8, “Acid-Base Equilibria,” quantitative values of acid ionization constants were taken by measuring the pH. The concept of buffer solutions resisting change in pH was also presented. When determining the dissociation constant (Ka) of Acetic Acid, the Titration curve was fundamental, thus a pH meter was used. Once the pH meter electrodes were calibrated, the experiment was in progress. Concentration variations were utilized, but ultimately in objective 1 the burette was filled with 0.1 M NaOH and a beaker was filled with 25.0 mL of 0.10 M acetic acid. The titration of NaOH into the acetic acid continued until the pH value reached 11.5. With the pH of 11.5, key points in the titration were present. For example, the “half-equivalence point” where half of the HA has been consumed, and the “equivalence point” where HA has been completely consumed by OH-. The titration required 23.7 mL of NaOH to reach equivalence point and 11.85 mL to reach half-equivalence point. The Ka was obtained by using the pKa at half equivalence point such as in this equation:

pH = pKa → Ka = 10-pKa → Ka= 10-4.428 → Ka= 3.73 x 10-5

The percent error from the experiment’s Ka is:

|((3.73x 10^(-5)) - (1.76 x 10^(-5)))/((1.76 x 10^(-5)))|x 100%= 111.93%

This implies that the measured Ka was more than the known Ka value of acetic acid.

Further into the lab, objective 2 was made, which was the practice of finding the dissociation of Acetic Acid. In this objective, variations of volumes of acetic acid (HA) and sodium acetate (NaA) were made in order to observe the difference in volumes of solutions can affect the dissociation constant. The Ka itself was obtained by calculating the [H+] from the observed pH and the [HA] and [A-] was calculated by applying the value from [H+] into an “ICE table.” From the equilibrium values from the ICE table, the Ka was determined by multiplying the equilibrium values of [H+] and [A-] and dividing it by the equilibrium value of [HA]. The acetic acid Ka derived from this objective was (1.15 x 10-5), whereas the accepted value is (1.76 x 10-5). With this information the percent error can be calculated through the equation:

|((1.15 x 10^(-5)) - (1.76 x 1〖0^ 〗^(-5)))/((1.76 x 10^(-5)))|x 100%= 34.66%

The value of Ka obtained is smaller than the known value of Ka. The pH values in the 2nd objective vary because the 25 ml NaA with 10 mL HA ended with a slightly higher pH that the calculated pH, resulting in a 3.88% error. 30 ml NaA with 5 mL HA ended up with a much lower pH than calculated with 30.4% error.

Later, another objective was made, which was testing the efficiency of a buffer and its capacity when a strong acid and strong base are added into a buffered and unbuffered solution. It was quite evident that buffered solutions actually resist change, such as change in pH. After calculating the expected pH value, it was evident that the unbuffered solutions were more prone to change, whereas the buffered solution was resistant towards the change in pH. This was especially seen the solutions mixed with HCl, where the unbuffered had a 48.3% error and the buffered solution had a 0.5% error in pH.

Overall, possible sources of error is when different people in our group of 3 would pour the solution under the hood, which can cause random error. This is maybe why there is a large percent error in the 2nd solution in the last set of solutions in Part 2,. Another major source of error in objective 3 is not stirring the solution with the unbuffered or buffered solution, thus the electrode could’ve measured pH of an unevenly blended overall solution, thus affecting our results. In part 1, sometimes different people would read the burette from different angles, and from different heights since we are not tall enough to perfectly see the top, thus affecting our graph results.