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

Short answer

The buffer neutralizes added H⁺ and OH⁻ ions through the reactions: \[ HCO_3^- + H^+ \rightarrow H_2CO_3 \] and \[ CO_3^{2-} + OH^- \rightarrow HCO_3^- \]

Step-by-step solution

Write the dissociation equations for NaHCO₃ and Na₂CO₃.

When dissolved in water, NaHCO₃ and Na₂CO₃ dissociate as follows: For NaHCO₃: \[ NaHCO_3 \rightarrow Na^+ + HCO_3^- \] For Na₂CO₃: \[ Na_2CO_3 \rightarrow 2Na^+ + CO_3^{2-} \]

Identify the acidic and basic species in the buffer.

In the buffer solution, HCO₃⁻ acts as the acidic species, and CO₃²⁻ acts as the basic species.

Write the equation for neutralization of H⁺ ions by HCO₃⁻.

When H⁺ ions are added to the buffer solution, HCO₃⁻ will react with them to neutralize the H⁺ ions. The reaction can be represented as follows: \[ HCO_3^- + H^+ \rightarrow H_2CO_3 \]

Write the equation for the neutralization of OH⁻ ions by CO₃²⁻.

When OH⁻ ions are added to the buffer solution, CO₃²⁻ will react with them to neutralize the OH⁻ ions. The reaction can be represented as follows: \[ CO_3^{2-} + OH^- \rightarrow HCO_3^- \] With these equations, we have found how the buffer solution neutralizes both H⁺ and OH⁻ ions added to it.

Key concepts

Understanding Buffer Capacity

Buffer capacity refers to a buffer solution’s ability to resist changes in pH when small amounts of an acid or a base are added. Think of it like a sponge, soaking up excess hydrogen ions (\(H^+\)) or hydroxide ions (\(OH^-\)) without allowing significant changes in pH.
The greater the buffer capacity, the larger the amount of acid or base that can be added before a notable change in pH occurs.
Buffer solutions have two main components: a weak acid and its conjugate base (or a weak base and its conjugate acid). These components work together to neutralize any added acids or bases, maintaining the pH.

• When \(H^+\) is added, it's neutralized by the buffer’s weak base component.
• When \(OH^-\) is added, the buffer’s weak acid component can neutralize it.

Choosing a buffer with high capacity is key when you need to stabilize the pH in a solution. This capability is crucial for many biological and chemical processes, helping to maintain conditions necessary for life and accurate reactions.

The Process of Acid-Base Neutralization

Acid-base neutralization is a chemical reaction in which an acid and a base react to form water and a salt, typically resulting in a neutral, or pH-balanced, solution. This is essentially what happens in a buffer solution when it "neutralizes" added acids or bases.
Neutralizing an acid (\(H^+\)) in a buffer means the base component reacts to form water: \(HCO_3^- + H^+ \rightarrow H_2CO_3\). Here, bicarbonate ion (\(HCO_3^-\)) reacts with hydrogen ions to form carbonic acid (\(H_2CO_3\)), preventing the pH from dropping.
Similarly, neutralizing a base (\(OH^-\)) involves the acid component reacting to buffer the solution: \(CO_3^{2-} + OH^- \rightarrow HCO_3^-\). Carbonate ion (\(CO_3^{2-}\)) reacts with hydroxide ions to produce bicarbonate ion, preventing pH from rising.
Such reactions ensure the pH remains stable, making the buffer solution vital in experiments and life processes where controlled pH is essential.

Exploring the Bicarbonate-Carbonate Buffer System

The bicarbonate-carbonate buffer system is a perfect example of how buffers work to maintain pH balance. This buffer, formed from sodium bicarbonate (\(NaHCO_3\)) and sodium carbonate (\(Na_2CO_3\)), is particularly effective at managing shifts in pH.
The equilibrium between bicarbonate ions (\(HCO_3^-\)) and carbonate ions (\(CO_3^{2-}\)) allows the buffer to neutralize both acids and bases.

• Bicarbonate ion acts as a weak acid, neutralizing added base (\(OH^-\)).
• Carbonate ion serves as a weak base, neutralizing added acid (\(H^+\)).

This buffering action is crucial in natural systems, such as the human bloodstream, where it helps maintain a stable pH, ensuring optimal conditions for enzymes and chemical reactions necessary for life.
Without this system, slight changes in metabolic processes could lead to severe pH swings, demonstrating the system’s importance in homeostasis.