A certain buffer is made by dissolving \(\mathrm{NaHCO}_{3}\) and \(\mathrm{Na}_{2} \mathrm{CO}_{3}\) in some water. Write equations to show how this buffer neutralizes added \(\mathrm{H}^{+}\) and \(\mathrm{OH}^{-}\).

To understand how a buffer works, we need to start by looking at the dissociation of the buffering agents in water. Sodium bicarbonate (\texttt{NaHCO}_3) is a common buffering agent that, when dissolved in water, separates into sodium ions (\texttt{Na}^+) and bicarbonate ions (\texttt{HCO}_3^-).

This dissolution is crucial for the buffering action as the bicarbonate ion plays a significant role in neutralizing added acids (H⁺ ions) or bases (OH⁻ ions) in a solution. The dissociation can be represented by the chemical equation:
\[\texttt{NaHCO}_3(s) \rightarrow \texttt{Na}^{+}(aq) + \texttt{HCO}_3^{-}(aq)\]

Dissociation allows \texttt{NaHCO}_3 not only to release bicarbonate ions but also to maintain a certain level of pH by moderating changes caused by the addition of strong acids or bases.

Moving on from sodium bicarbonate, we turn our attention to sodium carbonate (\texttt{Na}_2\texttt{CO}_3). When it dissociates in water, it produces two sodium ions (\texttt{Na}^+) for each formula unit and carbonate ions (\texttt{CO}_3^{2-}).

The carbonate ion serves a similar purpose as bicarbonate—helping to buffer the solution. However, it has the ability to absorb two protons (H⁺), making it doubly effective against acids. The decomposition in water can be written as:
\[\texttt{Na}_2\texttt{CO}_3(s) \rightarrow 2\texttt{Na}^{+}(aq) + \texttt{CO}_3^{2-}(aq)\]

The presence of both \texttt{NaHCO}_3 and \texttt{Na}_2\texttt{CO}_3 in the buffer solution offers a two-tiered defense mechanism against pH changes, with both bicarbonate and carbonate ions working in tandem to neutralize different amounts of acid or base.

Bicarbonate (\texttt{HCO}_3^-) and carbonate (\texttt{CO}_3^{2-}) ions are the cornerstone of the buffer system in the discussed solution.

When an acid is added to the buffer, the bicarbonate ion can react with the hydrogen ions (H⁺) to form carbonic acid (\texttt{H}_2\texttt{CO}_3), which is relatively weak and does not significantly lower the pH of the solution. On the other hand, carbonate ions can react with hydrogen ions to form bicarbonate, which can then further react with more hydrogen ions in the manner just described.

Neutralizing Acids

For bicarbonate, the reaction with added acid can be shown as:
\[\texttt{HCO}_3^{-} + \texttt{H}^{+} \rightarrow \texttt{H}_2\texttt{CO}_3\]

For carbonate, the reaction can accommodate two protons:
\[\texttt{CO}_3^{2-} + \texttt{H}^{+} \rightarrow \texttt{HCO}_3^{-}\]
\[\texttt{HCO}_3^{-} + \texttt{H}^{+} \rightarrow \texttt{H}_2\texttt{CO}_3\]

Neutralizing Bases

Conversely, when a base is added, the carbonic acid formed can dissociate to yield hydrogen ions, thereby neutralizing the hydroxide ions (OH⁻):
\[\texttt{H}_2\texttt{CO}_3 \rightarrow \texttt{HCO}_3^{-} + \texttt{H}^{+}\]\[\texttt{H}^{+} + \texttt{OH}^{-} \rightarrow \texttt{H}_2\texttt{O}\]

The ability of bicarbonate and carbonate ions to engage in these reversible reactions is key to maintaining a stable pH in the buffer solution, protecting it against drastic changes when acids or bases are introduced.