Question

The bicarbonate buffer system operates in the body's extracellular fluid by the following process: $$ \begin{array}{r} \mathrm{CO}_{2}(g)+\mathrm{H}_{2} \mathrm{O}(l) \Longrightarrow \mathrm{H}_{2} \mathrm{CO}_{3}(a q) \Longrightarrow \\ \mathrm{HCO}_{3}^{-}(a q)+\mathrm{H}^{+}(a q) \end{array} $$ What is expected to happen to the blood pH if the lungs are unable to expel \(\mathrm{CO}_{2}\) at the same rate as the body produces it? Describe what happens in terms of Le Chatelier's principle.

Short answer

If the lungs are unable to expel CO2 at the same rate it's being produced, more H+ ions are produced as a response, thus increasing the acidity and lowering the blood pH. This happens because of Le Chatelier's principle leading to shift in equilibrium position to minimize the change.

Step-by-step solution

Understand the Bicarbonate Buffer System

The bicarbonate buffer system is a chemical system that regulates the pH of blood and other biological fluids. The chemical reactions are as follows: \( \mathrm{CO}_{2}(g)+\mathrm{H}_{2} \mathrm{O}(l) \Longrightarrow \mathrm{H}_{2}\mathrm{CO}_{3}(a q) \Longrightarrow \ \mathrm{HCO}_{3}^{-}(a q)+\mathrm{H}^{+}(a q)\)

Apply Le Chatelier's Principle

Le Chatelier's principle states that if a dynamic equilibrium is disturbed by changing the conditions, the position of equilibrium shifts in the direction that'll minimize the change. Applying this principle to the bicarbonate buffer system, we can say that if CO2 is not expelled at the same rate as it's produced, it will accumulate and the reaction will be pushed to the right to minimize the change.

Understand the implications

When the reaction is pushed to the right, more \( \mathrm{H}_{2}\mathrm{CO}_{3}(a q) \) and then \( \mathrm{HCO}_{3}^{-}(a q) \) and \( \mathrm{H}^{+}(a q) \) are produced. The accumulation of \( \mathrm{H}^{+}(a q) \) ions increases the acidity, thus decreasing the pH of blood.

Key concepts

Le Chatelier's Principle and the Buffer System

Le Chatelier's principle is a fundamental concept in understanding how equilibrium systems react to changes. In the context of chemistry, it explains how a system at equilibrium responds to disturbances. If a change is introduced to a system, such as a concentration change, pressure change, or temperature change, the system will adjust to counteract the disturbance and restore equilibrium.

This principle is directly applicable to the bicarbonate buffer system in the human body. Here, the reversible reactions involved in this system maintain the blood pH around 7.4, essential for homeostasis. When there's an increase in carbon dioxide (CO₂) levels, as noted in the equation:

• CO₂(g) + H₂O(l) ⇌ H₂CO₃(aq) ⇌ HCO₃⁻(aq) + H⁺(aq)

According to Le Chatelier's principle, an increase in CO₂ will drive the reaction towards the right. This leads to more production of bicarbonate ions (HCO₃⁻) and hydrogen ions (H⁺), leading to changes in blood pH.

The system's response is crucial for maintaining physiological balance, emphasizing the importance of equilibrium concepts in biological systems.

Blood pH Regulation through Bicarbonate Buffering

The regulation of blood pH is a vital physiological function, primarily managed by the bicarbonate buffer system. Maintaining a stable pH is essential for the survival of organisms as many biochemical processes are pH-dependent. The normal pH range of blood is 7.35 to 7.45, slightly alkaline.

The bicarbonate buffer system's equation provides insight into how this regulation is achieved:

• CO₂(g) + H₂O(l) ⇌ H₂CO₃(aq) ⇌ HCO₃⁻(aq) + H⁺(aq)

When there is an accumulation of CO₂ due to inadequate expulsion by the lungs, as might occur during respiratory distress, more carbonic acid (H₂CO₃) forms. Subsequently, this leads to more bicarbonate (HCO₃⁻) and hydrogen ions (H⁺).

This increase in H⁺ leads to a lower pH, indicating acidosis. The body tries to counteract this by increasing respiration to expel more CO₂, thus pushing the reaction to the left and helping bring the pH back to normal. This mechanism shows the interplay between chemical equilibrium and biological functions necessary for life.

Carbon Dioxide Equilibrium's Role in pH Balance

Carbon dioxide equilibrium plays a crucial role in maintaining the acid-base balance in the body. CO₂ is produced as a byproduct of cellular respiration and transported in the blood, where it has a significant impact on pH. The reversible reaction involving CO₂ is central to the bicarbonate buffer system:

• CO₂(g) + H₂O(l) ⇌ H₂CO₃(aq) ⇌ HCO₃⁻(aq) + H⁺(aq)

If CO₂ builds up in the blood due to conditions such as hypoventilation, this equilibrium is disturbed. As more CO₂ is dissolved into the blood, more H₂CO₃ and subsequently H⁺ ions are formed, leading to acidosis.

The body's response involves several mechanisms:

• Increase in breathing rate to expel more CO₂
• Renal compensation to excrete H⁺ ions and reabsorb HCO₃⁻

These actions illustrate the dynamic nature of biological buffers, particularly under circumstances that challenge the normal balance. Understanding CO₂ equilibrium helps in grasping how the body maintains its pH despite environmental or physiological changes, underscoring the complexity and efficiency of biological systems.