Lab 11 - Acid-Base Studies

# The greater the extent of this reaction the larger

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Unformatted text preview: of water. The greater the extent of this reaction, the larger the equilibrium constant is for the reaction. This equilibrium constant is defined as the acid dissociation constant or Ka of the acid. The larger the Ka, the stronger the acid is. The general form of the acid dissociation reaction and Ka are shown below in equation 1. (1) Ka = HA + H2O A- + H3O- [A]+[H3O+] [HA] The basicity or alkalinity of a Brønsted base is a measure of the extent to which a Brønsted base reacts with water to produce its conjugate acid and OH- ions, the conjugate base of water. The equilibrium constant for the reaction of a base with water is given the symbol Kb . The general form for this chemical reaction and Kb are shown below in equation 2. (2) Kb = A- + H2O HA + OH- [HA][OH-] [A-] The acidity of an aqueous solution is, therefore, measured by the concentration of H3O+, and the basicity by the concentration of OH-. Most solutions we deal with are aqueous solutions, and these concentrations are important characteristics. We have seen that water behaves both as an acid (Reaction 1) and a base (Reaction 2), so it should not be surprising that it can react with itself. (3) H2O + H2O H3O+ + OH- Kw = [H3O+][OH-] Hydronium and hydroxide ion concentrations can be very small. For example, the Kw = 1.0 × 10-14 at 298 K, so in pure water at 298 K, [H3O+] = [OH-] = 1.0 × 10-7 M. Negative exponents are avoided by defining the pH: (4) pH = - log[H3O+] The pH of pure water is - log(1.0 × 10-7) = 7.0. Solutions in which [H3O+] = [OH-] are said to be neutral, so neutral solutions have a pH of 7.0 at 298 K. Solutions in which [H3O+] > [OH-] are acidic and pH < 7. Solutions in which [OH-] > [H3O+] are basic or...
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## This document was uploaded on 01/22/2014.

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