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The strength of a weak acid is usually represented as an equilibrium constant.
The acid-dissociation equilibrium constant (Ka), which measures the propensity of an acid to dissociate, for the reaction is:
The greater [H+] x [A-] is than [HA], the greater the value of Ka, the more the formation of H+ is favored, and the lower the pH of the solution.
ICE Tables: A Useful Tool For Solving Equilibrium Problems
ICE (Initial, Change, Equilibrium) tables are very helpful tools for understanding equilibrium and for calculating the pH of a buffer solution.
They consist of using the initial concentrations of reactants and products, the change they undergo during the reaction, and their equilibrium concentrations.
Consider, for example, the following problem:
Calculate the pH of a buffer solution that initially consists of 0.0500 M NH3 and 0.0350 M NH4+.
(Note: Ka for NH4+ is 5.6 x 10-10).
The equation for the reaction is as follows:
$NH_4^+ \rightleftharpoons H^+ + NH_3$
We know that initially there is 0.0350 M NH4+ and 0.0500 M NH3.
Before the reaction occurs, no H+ is present so it starts at 0.
During the reaction, the NH4+ will dissociate into H+ and NH3.
Because the reaction has a 1:1 stoichiometry, the amount that NH4+ loses is equal to the amounts that H+ and NH3 will gain.
This change is represented by the letter x in the following table.
Therefore the equilibrium concentrations will look like this:
Apply the equilibrium values to the expression for Ka.