The Common Ion Effect is the shift in equilibrium that occurs because of the addition of an ion already involved in the equilibrium reaction.
AgCl(s) <=> Ag^{+}(aq) + Cl^{}(aq)
<Addition of NaCl Shifts this equilibrium to the left.
Addition of common ion to a weak acid/base system:
HA <=> H^{+} + A^{}
Now add A^{} ( as a salt ) and the reaction will be driven to left
and [ H^{+} ] will decrease
Example:
CH_{3}COOH <=> H^{+} + CH_{3}COO^{}
Now add NaCH_{3}COO, where acetate is the common ion.
This increases concentration of acetate ion and the reaction is driven to left and the [H^{+}] decreases. The addition of the common ion CH_{3}COO^{} to a CH_{3}COOH solution is similar to titrating the acid with NaOH since both operations reduce the [H^{+}]. However instead of converting protons to water the common ion combines with protons to form more weak acid.
Practice Problem:
What is the [H^{+}] in a solution containing both 0.50 M HF and 0.10 M NaI?
Remember that salts generally dissociate strongly in water so that the concentration of the anion is roughly the same as the initial concentration of the salt:
Now let's consider the mixture of a weak acid (HA) and its salt formed by strong base (XA) in aqueous solution in more general terms. We know that the weak acid partially dissociates to form H^{+} and A^{} and the salt completely dissociates to form X^{+} and A^{}.
From the acid equilibrium we can write the following relationship:
The amount of pure weak acid that dissociates is small. It will only be ~ 1.6% for a 0.1 M CH_{3}COOH solution (pKa = 4.74). Thus the concentration of acid HA is still about the same as the amount of acid initially added to water. Likewise, the amount of A^{} formed by dissociation of the acid will be much smaller than the amount of A^{} formed by dissociation of the salt  hence [A^{}] is approximately the same as the concentration of salt added.We can use these facts to create an equation to simplify the calculation like the one we just completed above:
LIMITATIONS:
It is assumed that the amount of acid that dissociates is small – thus this relationship applies best to weaker acids. It cannot be used with any confidence for acids with pKa < 2.0.
The assumption is also made that very little anion (A) is contributed by dissociation of weak acid. This is not the case during the initial part of a titration of a weak acid by strong base (or weak base with strong acid)  calculation of titration curves using the HendersonHasselbalch equation is subject to error in initial phase.
Using the HendersonHasselbalch Equation:
1) Calculate the concentration of H^{+} and the pH of a solution that is 0.15 M in acetic acid and 0.25 M in sodium acetate. (K_{a}=1.8e5)
We can do this easily enough using an ICE table
HC_{2}H_{3}O_{2} + 
H_{2}O 
<=> H^{+} 
C_{2}H_{3}O_{2}^{} 

0.15M 
0.25M 

x 
+x 
+x 

0.15x 
x 
0.15+x 
K_{a} = [(x)(0.25+x)]/0.15x = 1.8e5
Assuming x is very small compared to the concentration, the equation reduces to: 0.025x/0.015 = 1.8e5; and x = 1.08e5 so the pH is 4.96.
OR
We can use the HendersonHasselbalch Equation:
pH = log(1.8e5) + log(0.25/0.15) = 4.74 + .2218 = 4.96