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General Chemistry Ka Lab

General Chemistry Ka Lab - Determination of Ka for Weak...

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Determination of K a for Weak Acids Introduction This experiment was performed in order to determine the acid dissociation constant, K a , of an unknown white substance which could possibly be lysergic acid, a precursor to LSD. This white substance was compared to a real sample of lysergic acid to verify that the substance smuggled across the U.S.-Mexico border is illegal. The numerical value in question, K a , or the acid dissociation constant, refers to the amount of dissociation an acid undergoes in aqueous solution. The ionization of HA, a default weak acid, is shown below: HA +H 2 O H 3 O + + A - The reaction moves in the forward direction and the reverse direction, meaning that it reaches equilibrium and can be shifted in either direction. Being an equilibrium reaction and the fact that HA is a weak acid, it dissociates only partially into separate ions. One of these ions, H + , becomes bound to the water molecules, forming H 3 O + , hydronium ion. The other ion, A - , is the conjugate base of HA. Because it acts as a proton donor, HA is the Bronsted-Lowry acid in the equation above, and A - is its conjugate base. One equation that could be used to calculate the acid dissociated constant is shown below: K a = [H 3 O + ] [A - ] _______________ [HA] As shown above, the acid dissociation constant is equal to the equilibrium concentrations of the ion products divided by the equilibrium concentration of the weak acid reactant. The stronger the acid, the more the acid dissociates into ions; this results in a larger concentration in products and a smaller concentration in the acid reactant. Therefore, K a is directly related to the strength of the acid. This particular equation was not used for practical purposes in this experiment because the equilibrium concentrations of the species could not be determined. Instead, the Henderson-Hasselbach equation, which is shown below, was used. pH = pK a + log( [base] / [acid] ) When the acetic acid and unknown acid were titrated, the pH of the solution was monitored with a MeasureNet probe. During these titrations, the NaOH neutralizes the acid because the hydroxide ions neutralize the hydronium ions. The Henderson-Hasselbach equation was simplified by using the concentrations of the Bronsted- Lowry acid and its conjugate base at the half equivalence point, because this is the point during titration where the
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OH - and H + ions are stoichiometrically equal. This is because exactly half of the acid had taken the form of its conjugate base.
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