olutions that contain a weak acid, HA, and its conjugate base, A−, are called buffer solutions because they resist drastic changes in pH. When a solution contains a weak acid and its conjugate base or a weak base and its conjugate acid, it will be a buffer solution. For buffers containing a weak acid, the principal reaction is HA(aq)+H2O(l)⇌H3O+(aq)+A−(aq) Unbuffered, a weak acid would ionize and the net reaction would proceed forward to reach equilibrium. However, when a significant amount of conjugate base is already present, the extent of the net reaction forward is diminished. Thus, equilibrium concentrations of HA and A− are approximately equal to their initial concentrations. Visualizing buffers You can visualize a buffer solution containing approximately equal concentrations of HA and A− in action using the following equation, where Ka is the equilibrium constant: Ka[H3O+]==[H3O+][A−][HA]Ka[HA][A−] When the concentrations of HA and A− are relatively large, the addition of a small amount of acid or base changes their concentration negligibly. The result is that the ratio [HA]/[A−] maintains a value approximately equal to 1, and the [H3O+] is approximately equal to the Ka value under buffer conditions. Part A Calculate the pH of the solution made by adding 0.50 mol of HOBr and 0.30 mol of KOBr to 1.00 L of water. The value of Ka for HOBr is 2.0×10−9. Express your answer numerically