In experiment 8, “Acid-base equilibria”, quantitative values ​​of acid ionization constants were taken by measuring pH. The concept of buffer solutions resistant to pH changes was also presented. When determining the dissociation constant (Ka) of acetic acid, the titration curve was fundamental, therefore a pH meter was used. Once the pH meter electrodes were calibrated, the experiment was underway. Variations in concentration 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 into acetic acid continued until the pH value reached 11.5. With a pH of 11.5, the key points of the titration were present. For example, the “equivalence point” where half of the HA has been consumed, and the “equivalence point” where the HA has been entirely consumed by OH-. The titration required 23.7 mL of NaOH to reach the equivalence point and 11.85 mL to reach the half-equivalence point. The Ka was obtained by using the pKa at the midpoint of equivalence, as in this equation: Say no to plagiarism. Get a tailor-made essay on “Why Violent Video Games Should Not Be Banned”? Get the original experimentpH = pKa → Ka = 10-pKa → Ka= 10-4.428 → Ka= 3.73 x 10-5The percentage error of 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 greater than the known Ka value of acetic acid. Further into the laboratory, objective 2 was achieved, which was the practice of finding the dissociation of acetic acid. For this purpose, variations in volumes of acetic acid (HA) and sodium acetate (NaA) were carried out in order to observe the difference in solution volumes that could affect the dissociation constant. The Ka itself was obtained by calculating the [H+] from the observed pH and the [HA] and [A-] were calculated by applying the value of [H+] in an “ICE table”. From the equilibrium values ​​in 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), while the accepted value is (1.76 x 10-5). With this information, the percentage error can be calculated using 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 lower than the known value of Ka. The pH values ​​in the 2nd objective vary because the 25 ml of NaA with 10 ml of HA ends up with a pH slightly higher than the calculated pH, resulting in an error of 3.88%. 30 ml of NaA with 5 ml of HA resulted in a pH much lower than calculated with an error of 30.4%. Later, another goal was set: to test the effectiveness of a buffer and its capacity when a strong acid and a strong base are added in a buffered and unbuffered solution. It was quite evident that buffered solutions actually resist changes, such as changes in 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 pH change. This was particularly observed in solutions mixed with HCl, where the unbuffered solution had a 48.3% error and the buffered solution had a 0.5% pH error. Keep in mind: this is just a sample. Get a personalized paper now from our expert writers. Get a Custom Essay Overall, the possible sources of error arise when different people in our group of 3..