The Importance of Buffers in Natural Systems
Dot-Point 7: describe the importance of buffers in natural systems (ACSCH098, ACSCH102)
Buffers play a critical role in maintaining the pH of various natural systems, ensuring the stability required for life and environmental processes. This post explores the importance of buffers in the ocean, blood, and intracellular fluids, highlighting how they counteract pH fluctuations and preserve functionality.
Ocean Buffers: Balancing Marine pH
The Challenge:
Rising levels of carbon dioxide (CO₂) in the atmosphere lead to increased dissolution of CO₂ in the ocean. This process can release hydronium ions (H₃O⁺), potentially lowering the ocean’s pH and threatening marine ecosystems.
Key Reactions:
CO₂ (g) + H₂O (l) ⇌ H₂CO₃ (aq)
H₂CO₃ (aq) + H₂O (l) ⇌ HCO₃⁻ (aq) + H₃O⁺ (aq)
HCO₃⁻ (aq) + H₂O (l) ⇌ CO₃²⁻ (aq) + H₃O⁺ (aq)
Buffer System:
The ocean's buffer relies primarily on bicarbonate (HCO₃⁻, 88%) and carbonate (CO₃²⁻, 11%), which interact to regulate pH.
When acidic conditions arise, hydronium ions (H₃O⁺) increase, causing the equilibrium to shift left, reducing the hydronium concentration.
When basic conditions arise, hydroxide ions (OH⁻) increase, and the equilibrium shifts left in a complementary reaction, maintaining balance.
Significance:
The ocean's pH, typically between 8.0 and 8.3, supports diverse marine life. Coral reefs, for example, require a pH of 8.2-8.3 for proper skeleton growth. Buffering prevents rapid acidification, which could hinder marine biodiversity and disrupt ecosystems.
Blood Buffers: Maintaining Homeostasis
The Challenge:
The human body must keep blood pH tightly regulated between 7.35 and 7.45 for physiological processes to function optimally.
Key Buffer System:
The carbonic acid-bicarbonate buffer system stabilizes blood pH:
CO₂ (g) + H₂O (l) ⇌ H₂CO₃ (aq)
H₂CO₃ (aq) + H₂O (l) ⇌ HCO₃⁻ (aq) + H₃O⁺ (aq)
HCO₃⁻ (aq) + H₂O (l) ⇌ CO₃²⁻ (aq) + OH⁻ (aq)
Responses to pH Changes:
Blood pH < 7.35 (Acidic):
Equation 2 shifts left to reduce hydronium ions.
Increased exhalation removes CO₂, shifting Equation 1 left and further reducing acidity.
Blood pH > 7.45 (Basic):
Equation 3 shifts left to decrease hydroxide ions.
Slower breathing retains CO₂, shifting Equation 1 right. This produces more H₂CO₃, which dissociates to restore H₃O⁺ levels.
Significance:
This buffer ensures that blood remains within its narrow pH range, critical for enzyme activity and overall physiological function.
3. Intracellular Buffers: Stabilizing Cellular Environments
The Challenge:
Enzymes and biochemical processes inside cells require a stable pH between 7.0 and 7.4. Deviations can denature enzymes and disrupt cellular function.
Key Buffer System:
The intracellular buffer system uses dihydrogen phosphate (H₂PO₄⁻) and hydrogen phosphate (HPO₄²⁻):
H₂PO₄⁻ (aq) + H₂O (l) ⇌ HPO₄²⁻ (aq) + H₃O⁺ (aq)
HPO₄²⁻ (aq) + H₂O (l) ⇌ H₂PO₄⁻ (aq) + OH⁻ (aq)
Responses to pH Changes:
In acidic conditions, H₂PO₄⁻ reduces H₃O⁺ by shifting equilibrium left.
In basic conditions, HPO₄²⁻ reduces OH⁻, maintaining stability.
Significance:
A stable intracellular pH preserves enzyme structure and activity, ensuring efficient cellular processes and overall health.
