In Experiment 8, “Acid-Base Equilibria,” quantitative values ​​of acid ionization constants were found by measuring pH. The concept of buffer solutions that resist pH change was also presented. To determine the dissociation constant (Ka) of acetic acid, the titration curve was critical, so a pH meter was used. Once the pH meter electrodes were calibrated, the experiment was underway. Concentration variations were used, 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 to acid acetic acid was continued until the pH value reached 11.5. At the pH of 11.5 the key titration points were present. For example, the “half-equivalence point” where half of the HA has been consumed and the “equivalence point” where the HA has been completely consumed by OH-. The titration required 23.7 mL of NaOH to reach the equivalent point and 11.85 mL to reach the semi-equivalence point. The Ka was obtained using the pKa at half equivalence point as in this equation: Say no to plagiarism. Get a tailor-made essay on "Why Violent Video Games Shouldn't Be Banned"? Get an original assaypH = pKa → Ka = 10-pKa → Ka= 10-4.428 → Ka= 3.73 x 10-5The percent error from the Ka of the experiment 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 higher than the known Ka value of acetic acid. Subsequently, objective 2 was achieved in the laboratory, which was the practice of finding the dissociation of acetic acid. In this objective, variations in the volumes of acetic acid (HA) and sodium acetate (NaA) were made to observe the difference in the volumes of the solutions which can influence the dissociation constant. The Ka itself was obtained by calculating [H+] from the observed pH, and [HA] and [A-] were calculated by applying the value of [H+] in an “ICE table”. From the equilibrium values ​​of the ICE table, the Ka was determined by multiplying the equilibrium values ​​of [H+] and [A-] and dividing them by the equilibrium value of [HA]. The acetic acid Ka derived from this target was (1.15 x 10-5), while the accepted value is (1.76 x 10-5). With this information the percentage 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 obtained Ka value is lower than the known Ka value. The pH values ​​in the 2nd objective vary because the 25 mL of NaA with 10 mL of HA ended up with a slightly higher pH than the calculated pH, resulting in an error of 3.88%. 30 mL of NaA with 5 mL of HA resulted in a much lower pH than calculated with an error of 30.4%. Next, another goal was set, which was to test the efficiency of a buffer and its capacity when a strong acid and a strong base are added into a buffered and unbuffered solution. It was quite evident that buffered solutions actually resist change, such as changing pH. After calculating the expected pH value, it was evident that the unbuffered solutions were more prone to change, while the buffered solution was resistant to the pH change. This was especially observed in solutions mixed with HCl, where the unbuffered solution had a 48.3% error and the buffered solution had a 0.5% error in pH. Please note: this is just a sample. Get a custom paper from our expert writers now. Get a Custom Essay Overall, possible sources of error are when different people in our group of 3 pour the solution under the hood, which can cause random errors. This is perhaps the.