Conclusion:
Throughout the course of the lab, the titration of 0.1 M NaOH was utilized on an unknown weak acid as to determine its identity. The titration process involved the initial dropping of 1 mL NaOH increments before switiching to 0.2 mL increments as the rate of pH change increased. The equilibrium point of the acid-base reaction was determined to be a pH of 10.05 upon 8 mL of titration, while the pKa value was determined to be 7.20. In total, 12 mL of the 0.1M NaOH was titrated, and through the utilization of stoichiometry and molarity equations, the molar mass of our .1010 g weak acid was determined to be 126.25 g. Given that pKa is equavalent to -log(Ka), it was determined that the Ka of the reaction was 6.3E-8. In reference to the table of acids provided with the lab documents, the acid was determined to be potassium hydrogen sulfite (KSO3), which has a molar mass of 120.16 g, a pKa of 7.19, and a Ka2 of 6.4E-8. Given these values, it was determined that the percent error for molar mass was 5.07%, while the percent errors for pKa and Ka were 0.139% and 1.56% respectively. The molar mass of the weak acid was slightly greater in comparison to the identified acid - such a discrepancy likely caused from inadequate cleaning of the the pH probe, resulting in inaccurate measurements of pH throughout the lab. Additional sources of error include improper weighing of the initial acid mass and inaccuracies in the mixture of the weak acid with distilled water. Regardless, this lab allowed for the further understanding of acid-base equilibrium reactions along with the respective Ka and Kb values related to such a reaction.