How to calculate acidic buffer solution pH, and how do they behave?

Example: Calculate the pH of a buffer solution which contains the weak monoprotic acid, propanoic acid (CH3CH2COOH), in concentration 0.1 moldm-3 and sodium propanoate in concentration 0.05 moldm-3. Ka of propanoic acid is 1.26×10-5 moldm-3. What happens when an acid is added to the solution? What happens when a base is added? This is an example of an acidic buffer solution, consisting of an acid (propanoic acid) and one of its salts (sodium propanoate). This gives the solution plenty of the acid (propanoic acid) and its anion (CH₃CH₂COO^-). Propanoic acid is a weak acid, so position of equilibrium for its dissociation lies well to the left. Adding propanoate ions (given from sodium propanoate) pushes this further left according to Le Chatelier's Principle. An assumption can therefore be made that [CH₃CH₂COO^-]=[added sodium propanoate]. A buffer solution system works to minimalise any change to the pH. If an acid is added, that means that there is an influx of protons. This pushes the equilibrium to the left to produce propanoic acid, by reacting the protons with the reservoir of propanoate ions. If a base is added, the added OH^- ions can be dealt with in two ways: They can react with propanoic acid to form propanoate ions and water, or react with protons to form water. As the proton concentration would drop, the equilibrium shifts to the right, dissociating more propanoic acid to upkeep proton concentration and therefore pH.