Question

One of the important pH-regulating systems in the blood consists of a carbonic acid-sodium hydrogen carbonate buffer: $$ \begin{gathered} \mathrm{H}_{2} \mathrm{CO}_{3}(a q) \rightleftharpoons \mathrm{H}^{+}(a q)+\mathrm{HCO}_{3}^{-}(a q) \\ \mathrm{NaHCO}_{3}(a q) \longrightarrow \mathrm{Na}^{+}(a q)+\mathrm{HCO}_{3}^{-}(a q) \end{gathered} $$ Explain how this buffer resists changes in \(\mathrm{pH}\) when excess acid, \(\mathrm{H}^{+}\), gets into the bloodstream.

Short answer

The buffer system neutralizes added H⁺ ions by converting them to H₂CO₃, preventing significant pH change.

Step-by-step solution

Understanding the Buffer System

A buffer system consists of a weak acid and its conjugate base, which helps in maintaining the pH of a solution by neutralizing added acids or bases. In this case, the system is made up of carbonic acid (H₂CO₃) and sodium hydrogen carbonate (NaHCO₃).

Identify the Key Reaction

The key equilibrium reaction for this buffer system is: H₂CO₃ (aq) ⇌ H⁺ (aq) + HCO₃⁻ (aq).

Addition of Excess Acid

When excess acid, represented by H⁺ ions, enters the bloodstream, it potentially lowers the pH. The increased concentration of H⁺ ions shifts the equilibrium towards the left, as given by Le Chatelier's Principle.

Reaction with Bicarbonate Ion

The added H⁺ ions are neutralized by the available bicarbonate ions (HCO₃⁻), forming more carbonic acid (H₂CO₃). The reaction is: H⁺ (aq) + HCO₃⁻ (aq) → H₂CO₃ (aq).

Resulting pH Stability

The formation of H₂CO₃ from H⁺ and HCO₃⁻ reduces the number of free H⁺ ions, decrease in H⁺ ion concentration prevents a significant drop in pH, thus stabilizing the pH of the blood.

Key concepts

pH regulation

The pH level, a measure of how acidic or basic a solution is, is critical in the human body. Blood must maintain a pH of about 7.4 for proper physiological functions. Even small deviations from this pH level can be harmful. The body uses several mechanisms to keep the pH stable, one of which is buffer systems. These systems neutralize excess acids or bases to prevent significant changes in pH. Without these buffers, the body would be vulnerable to severe pH imbalances and related health issues like acidosis or alkalosis.

carbonic acid-sodium hydrogen carbonate buffer

The buffer system in the blood that helps regulate pH involves carbonic acid (H₂CO₃) and sodium hydrogen carbonate (NaHCO₃), also known as sodium bicarbonate. This buffer system works through two main components:

• Carbonic Acid (H₂CO₃): A weak acid that partially dissociates into hydrogen ions (H⁺) and bicarbonate ions (HCO₃⁻) in an aqueous solution.
• Sodium Bicarbonate (NaHCO₃): A salt that fully dissociates in water to provide sodium ions (Na⁺) and bicarbonate ions (HCO₃⁻).

The equilibrium reaction can be represented as:
$$\text{H₂CO₃ (aq) ⇌ H⁺ (aq) + HCO₃⁻ (aq)}$$
This equilibrium allows the system to neutralize excess H⁺ ions or OH⁻ ions, maintaining the stable pH of the blood.

Le Chatelier's Principle

Le Chatelier's Principle states that if a dynamic equilibrium is disturbed by changing the conditions, the system responds to partially counteract the change and restore equilibrium. Here's how this principle applies to the carbonic acid-bicarbonate buffer system:

• Addition of Acid (H⁺): When excess H⁺ ions enter the bloodstream, the equilibrium shifts towards the formation of carbonic acid (H₂CO₃) to decrease the concentration of H⁺ ions.
• Addition of Base (OH⁻): When OH⁻ ions are added, they react with H⁺ ions to form water, causing a shift in equilibrium towards producing more H⁺ ions from H₂CO₃. This maintains the balance of H⁺ and OH⁻ ions and helps keep the pH stable.

By following Le Chatelier's Principle, the buffer system effectively resists changes in blood pH.

bicarbonate ion

The bicarbonate ion (HCO₃⁻) plays a crucial role in the carbonic acid-sodium hydrogen carbonate buffer system. Its primary function is to neutralize excess H⁺ ions. When an influx of H⁺ ions occurs, the bicarbonate ions combine with them, forming carbonic acid (H₂CO₃).
$$\text{H⁺ (aq) + HCO₃⁻ (aq) → H₂CO₃ (aq)}$$
This reaction effectively reduces the concentration of free H⁺ ions in the bloodstream, preventing a substantial drop in pH and maintaining the blood's acid-base balance.

blood pH stability

The stability of blood pH is essential for the proper functioning of enzymes and metabolic processes. The carbonic acid-sodium hydrogen carbonate buffer system is vital in maintaining this balance. Here's how it works:

• Neutralizing Excess Acid: When there are too many H⁺ ions, the buffer system reacts to form more carbonic acid, which dissociates less than bicarbonate, thus reducing the H⁺ concentration.
• Neutralizing Excess Base: When OH⁻ ions are added, they neutralize H⁺ ions, causing the system to produce more H⁺ by dissociating carbonic acid, thus maintaining the pH.

Overall, this buffering action ensures that any pH fluctuations are minimal, keeping the blood within its narrow pH range of around 7.4, necessary for health and survival.