A buffer is prepared by dissolving \(\mathrm{HONH}_{2}\) and \(\mathrm{HONH}_{3} \mathrm{NO}_{3}\) in some water. Write equations to show how this buffer neutralizes added \(\mathrm{H}^{+}\) and \(\mathrm{OH}^{-}\).

To demonstrate how this buffer neutralizes added H+ ions, we'll focus on the weak base HONH2 present in the mixture. When H+ ions are added to the buffer solution, the weak base reacts with them to form its conjugate acid. Here's the equation representing this process: HONH2 (aq) + H+ (aq) -> HONH3+ (aq) This reaction shows the buffer neutralizing the added H+ ions by forming HONH3+ ions, as the weak base HONH2 consumes the added H+ ions.

Now let's analyze how the buffer neutralizes added OH- ions. To do this, we'll focus on the conjugate acid HONH3NO3 present in the buffer solution. When OH- ions are added to the buffer, the conjugate acid reacts with the added OH- ions to form its conjugate base, HONH2, and water. The equation for this reaction is as follows: HONH3NO3 (aq) + OH- (aq) -> HONH2 (aq) + H2O (l) This equation illustrates the process of neutralizing added OH- ions, as the conjugate acid HONH3NO3 reacts with OH- ions to produce the weak base HONH2 and water. In summary, the buffer solution composed of HONH2 and HONH3NO3 neutralizes added H+ ions by allowing the weak base HONH2 to react with them, and added OH- ions by allowing the conjugate acid HONH3NO3 to react with them. This maintains the overall pH of the buffer solution and prevents significant changes in its acidity or alkalinity.