We already know that equilibrium systems are affected by changes in concentration. LeChâtelier's Principle generalizes this effect. When we studied solubility equilibrium we used the phrase "common ion effect" to describe one aspect of equilibrium change. Of course this effect is not limited to solubility equilibrium. In the lab we saw how the addition of a common ion to a weak acid can change pH and neutralization behavior. Understanding exactly why those changes occur is next.
The universal indicator we used is sensitive
enough to distinguish between strong and weak acids of the same
Molarity. So the 0.10 M HCl had a pH of about 1.0 while the 0.10
M CH3COOH showed a pH of about 2.5. We should
expect this difference since the acetic acid is weak and
therefore will not contain as much free H3O+.
It is perhaps easy enough to see why the pH of the HCl is 1.0 based on the
definition of pH. The reason behind the value of 2.5 for the acetic
acid is more obscured by Ka. A quick calculation
reveals the "why" 
