Question

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 \(\mathrm{OH}^{-}\) .

Short answer

The buffer made by dissolving NaHCO3 and Na2CO3 in water can neutralize added H+ ions and OH- ions as follows: When H+ ions are added: \( H^{+} (aq) + HCO_{3}^{-} (aq) \rightarrow H_{2}O (l) + CO_{2} (g) \) When OH- ions are added: \( OH^{-} (aq) + CO_{3}^{2-} (aq) \rightarrow CO_{3}^{2-} (aq) + H_{2}O (l) \) These reactions help to minimize changes in pH when acids or bases are added to the buffer solution.

Step-by-step solution

Write equations for the dissociation of NaHCO3 and Na2CO3 in water

When NaHCO3 and Na2CO3 are dissolved in water, they will dissociate as follows: NaHCO3 (aq) → Na+ (aq) + HCO3- (aq) Na2CO3 (aq) → 2Na+ (aq) + CO3^2- (aq)

Write equation for HCO3- reacting with H+ ions

When H+ ions are added to the buffer solution, they will react with HCO3- ions present in the solution. The reaction equation can be written as: H+ (aq) + HCO3- (aq) → H2O (l) + CO2 (g) In this reaction, the added H+ ions are consumed by HCO3- ions, preventing a significant change in pH of the solution.

Write equation for CO3^2- reacting with OH- ions

When OH- ions are added to the buffer solution, they will react with CO3^2- ions present in the solution. The reaction equation can be written as: OH- (aq) + CO3^2- (aq) → CO3^2- (aq) + H2O (l) In this reaction, the added OH- ions are consumed by CO3^2- ions, preventing a significant change in pH of the solution. To summarize, the buffer made by dissolving NaHCO3 and Na2CO3 in water can neutralize added H+ ions via reaction with HCO3- ions, and neutralize added OH- ions via reaction with CO3^2- ions. These reactions help to minimize changes in pH when acids or bases are added to the buffer solution.

Key concepts

Acid-Base Reactions

When discussing buffer solutions, the concept of acid-base reactions is fundamental. Buffers are specially designed to resist shifts in pH by neutralizing added acids or bases. In the case of the buffer formed from \(\mathrm{NaHCO}_3\) and \(\mathrm{Na}_2\mathrm{CO}_3\), two important reactions take place. The bicarbonate ion \(\mathrm{HCO}_3^-\), which acts as a weak acid, can neutralize added protons \(\mathrm{H}^+\) by forming water and carbon dioxide.
Another important part of this system is the carbonate ion \(\mathrm{CO}_3^{2-}\). Acting as a base, it can react with hydroxide ions \(\mathrm{OH}^-\) to form water, effectively preventing an increase in pH. - Reactivity: - \(\mathrm{HCO}_3^- + \mathrm{H}^+ \rightarrow \mathrm{H}_2\mathrm{O} + \mathrm{CO}_2\) - \(\mathrm{CO}_3^{2-} + \mathrm{OH}^- \rightarrow \mathrm{CO}_3^{2-} + \mathrm{H}_2\mathrm{O}\)
Through these reactions, buffer systems like this can maintain a stable hydrogen ion concentration, which is crucial for many chemical and biological processes in solutions. The balance of these reactions helps prevent drastic changes in pH, keeping environments stable.

pH Stability

pH stability is a key feature of buffer solutions. You might wonder, "How do buffers maintain such a stable pH?" It's because of their ability to engage in reversible reactions with both added acids and bases.
Consider a scenario where an acid is added to a buffered solution. The buffer will absorb these additional hydrogen ions instead of allowing them to accumulate and lower the pH significantly. Conversely, if a base is added, the buffer consumes the hydroxide ions, preventing a substantial increase in pH. This ability to modulate strong changes in pH is what distinguishes an effective buffer.- Mechanism of Resistance: - Buffers resist decreases in pH by reacting with \(\mathrm{H}^+\). - They resist increases in pH by neutralizing \(^-\).
The efficiency of a buffer is defined by its buffering capacity, which varies with concentration and the specific chemical properties of the buffer components. Buffer systems like the carbonate buffer are especially important in biochemical contexts, such as blood, where stable pH is essential for health.

Carbonate Buffer System

The carbonate buffer system is a classic example of a chemical buffer used to stabilize pH in many settings, including natural waters and biological systems. This buffer is made from the weak acid \(\mathrm{N}\mathrm{a}\mathrm{HCO}_3\) and the salt \(\mathrm{N}\mathrm{a}_2\mathrm{CO}_3\).
When prepared in an aqueous solution, \(\mathrm{N}\mathrm{a}\mathrm{HCO}_3\) dissociates into \(\mathrm{Na}^+\) and \(\mathrm{HCO}_3^-\), while \(\mathrm{N}\mathrm{a}_2\mathrm{CO}_3\) dissociates into \(\mathrm{2Na}^+\) and \(\mathrm{CO}_3^{2-}\). - Role of Ionic Species: - \(\mathrm{HCO}_3^-\) is important for buffering against acids. - \(\mathrm{CO}_3^{2-}\) helps to neutralize any added base.
This dual ability enables the carbonate buffer to handle both acids and bases, making it highly versatile. Its equilibrium lies between \(\mathrm{HCO}_3^-\) and \(\mathrm{CO}_3^{2-}\), which is essential for maintaining a consistent pH, particularly in environments prone to environmental changes in acidity or alkalinity.
In conclusion, the carbonate buffer system exemplifies how buffers operate to sustain pH, crucial in myriad scientific and practical applications.