Determination Of Ka Of A Weak Acid By Titration Lab Report

Determination of Ka of a Weak Acid by Titration Lab The purpose of the lab is to determine the dissociation constant of a weak acid through an acid-base titration. An unknown acid was added to DI water to form a solution. Phenolphthalein was added to the acid solution. The acid solution was then titrated with sodium hydroxide until it reached the endpoint. The same amount of acid solution was added to the titrated solution in order for the solution to reach its midpoint. The experiment was repeated. Theories and concepts explored in the experiment include dilution, equilibrium reactions, acid-base indicators, titration, weak acid-strong base reaction, pH, conjugate bases, equivalence point, endpoint, midpoint, dissociation, equilibrium reactions,

At the start of the experiment, an unknown acid, the solute, was dissolved in DI water, the solvent, diluting the concentration of the acid. Phenolphthalein, an acid-base indicator--or a weak acid that changes color when the equivalence point is reached--was added to the solution. The solution was then titrated, a process by which a solution with a known concentration, NaOH, is added to an unknown solution, the acid, in order to determine its concentration. The reaction consisted of an unknown weak acid and NaOH, a strong base, making the reaction a weak acid-strong base reaction. Because NaOH is a strong base, the pH of the reaction--or the measure of the acidity of the solution by taking the negative log base ten of the hydronium concentration--would increase when titrating the unknown acid. As the unknown acid was titrated, one of the products formed in the reaction was the ion of the unknown acid, or the conjugate base which, according to the Brønsted-Lowry theory, was formed as a result of the acceptance of a hydrogen ion from NaOH. When the moles of the unknown acid and the

One example of an error was fluctuating temperatures from air vents during the experiment. Because the Ka value is dependent solely on temperature, a slight increase in temperature for instance, can shift the reaction to the right, increasing the solubility of the reaction. This as a result, would have increased the concentrations of the products and the Ka value, explaining the 60.5% deviation in trial four. Another possible source of error is the loss of acid solution as it was transferred from the Erlenmeyer flask to the beaker. Although a small volume of the solution would have been lost, it is significant as a smaller volume would result in a lower volume of NaOH needed to titrate the acid, lowering the pH. A lowered pH would result in a higher Ka value, explaining the pH deviation of 33.3% from the actual pKa value in trial four and giving a reason as to why the Ka value in trial four is slightly higher than the other Ka