The reaction between acetic acid CHCOOH and sodium acetate CHCOONa can be represented as follows:CHCOOH + Na + CHCOO CHCOO + H + Na + CHCOOHHowever, since the sodium ion Na does not participate in the reaction, we can simplify the equation to:CHCOOH + CHCOO 2 CHCOO + HThis reaction represents a buffer system, where acetic acid acts as a weak acid and its conjugate base, acetate ion CHCOO , acts as a weak base. The equilibrium constant for this reaction, known as the acid dissociation constant Ka , can be written as:Ka = [CHCOO][H] / [CHCOOH]The pH of a solution is related to the concentration of hydrogen ions H in the solution:pH = -log[H]In a buffer system, the Henderson-Hasselbalch equation can be used to relate the pH, pKa where pKa = -log Ka , and the concentrations of the weak acid and its conjugate base:pH = pKa + log [CHCOO] / [CHCOOH] Now, let's analyze the effect of changing the pH on the equilibrium position of the reaction.1. If the pH is increased i.e., the solution becomes more basic , the concentration of H ions decreases. According to Le Chatelier's principle, the equilibrium will shift to counteract this change, meaning the reaction will shift to the right to produce more H ions. This results in the consumption of CHCOOH and the production of CHCOO.2. If the pH is decreased i.e., the solution becomes more acidic , the concentration of H ions increases. According to Le Chatelier's principle, the equilibrium will shift to counteract this change, meaning the reaction will shift to the left to consume H ions. This results in the consumption of CHCOO and the production of CHCOOH.In summary, changing the pH of the solution affects the equilibrium position of the reaction between acetic acid and sodium acetate. Increasing the pH shifts the equilibrium to the right, favoring the production of acetate ions and consumption of acetic acid. Conversely, decreasing the pH shifts the equilibrium to the left, favoring the production of acetic acid and consumption of acetate ions.