Determination of Ka for a Weak Acid

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Abstract The interaction at equilibrium between acids and bases during a titration can be used to determine several characteristics of the acid or the base. In this experiment 0.05 M of KHP titrate with the strong base NaOH. From the plot of pH versus the volume of sodium hydroxide added was found that acid ionization constant Ka for KHP at half point was equal to 8.3176x10-6, also the dissociation constant Kb for conjugate base of weak acid was equal to 1x10-9. This value was established by observing the pH versus volume of NaOH graph. The equivalence point of titration occurred at a volume of 31 mL 0.081M NaOH (aq). The halfway volume of titration was thus 15.5 mL and the pH of the titration system at this point, as determined graphically, was 5.08. After identification of half point, Henderson–Hasselbalch equation pH=pKa, was used to determine the value of Ka. The resulting Ka value and the concentration of KHP made it possible to calculate the initial pH of the acid, which were 3.19. How it was expected the pH at equivalence point was 9.17 this is because of the domination of hydroxide ion in solution. The relationship between the pH and the amount of titrant added offered a better understanding of the equilibrium properties of the acid. Introduction Titrations are a convenient and common method of analysis. Generally titration is an experiment where a known property of one solution is used to infer an unknown property of another solution. There are several types of titrations: Acid-base titrations are based on the neutralization reaction between the analyte and an acidic or basic titrant. Redox titrations are based on an oxidation-reduction reaction between the analyte and titrant. Complexometric titrations are based on the formation of a complex between the analyte and the titrant. Ex: the chelating agent EDTA is very commonly used to

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