Increasing the pH of a solution containing a weak acid and its conjugate base will shift the equilibrium position of the acid-base reaction. To understand this, let's consider the equilibrium reaction of a weak acid HA and its conjugate base A- :HA aq + H2O l H3O+ aq + A- aq The equilibrium constant for this reaction is given by the acid dissociation constant, Ka:Ka = [H3O+][A-] / [HA]Now, let's consider the effect of increasing the pH of the solution. Increasing the pH means that the concentration of H3O+ ions in the solution is decreasing. According to Le Chatelier's principle, when a system at equilibrium is subjected to a change in concentration, the system will adjust itself to counteract the change and re-establish equilibrium.In this case, as the concentration of H3O+ decreases, the equilibrium will shift to the left to produce more H3O+ ions. This means that more HA will be formed, and the concentration of A- will decrease.To further support this, we can use the Henderson-Hasselbalch equation, which relates the pH, pKa, and the ratio of the concentrations of the conjugate base and the weak acid:pH = pKa + log [A-] / [HA] As the pH increases, the ratio of [A-] / [HA] must decrease to maintain the equation's balance. This means that the concentration of the conjugate base A- decreases, and the concentration of the weak acid HA increases.In conclusion, increasing the pH of a solution containing a weak acid and its conjugate base will shift the equilibrium position of the acid-base reaction to favor the formation of the weak acid HA and decrease the concentration of the conjugate base A- .