Explain me about regulation of ph of acid and base
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Regulation of pH: Acid-Base Balance
1. What is pH and Why It Matters
pH is defined as the negative logarithm of hydrogen ion (H⁺) activity:
pH = −log₁₀[H⁺]
Normal arterial blood pH is 7.35–7.45 (average 7.40), corresponding to [H⁺] of ~40 nmol/L.
| Condition | pH |
|---|---|
| Acidemia | < 7.35 |
| Normal | 7.35–7.45 |
| Alkalemia | > 7.45 |
| Death (severe) | < 6.8 or > 7.8 |
Even tiny pH changes matter because most biologically important molecules (enzymes, receptors, ion channels, transporters) contain chemical groups that donate or accept H⁺. For example, phosphofructokinase activity falls ~90% with a pH drop of just 0.1; cell proliferation may fall 85% when intracellular pH drops by 0.4.
— Medical Physiology (Boron & Boulpaep)
2. The Three Lines of Defense
The body uses three interrelated systems to maintain pH:
1. Chemical Buffers (seconds)
2. Respiratory System (minutes)
3. Renal System (hours to days)
3. Chemical Buffer Systems
A buffer is any substance that reversibly consumes or releases H⁺, minimizing (not preventing) pH changes. Buffers work best within ±1 pH unit of their pKa.
A. Bicarbonate/Carbonic Acid Buffer (most important in plasma)
H⁺ + HCO₃⁻ ⇌ H₂CO₃ ⇌ CO₂ + H₂O
- pKa = 6.1 (seemingly far from plasma pH of 7.4)
- Despite the mismatch, it is the most effective buffer because:
- CO₂ is continuously exhaled by the lungs (open system)
- HCO₃⁻ is reclaimed/excreted by the kidneys
- It is present at relatively high concentrations
Normal plasma ratio: HCO₃⁻ / dissolved CO₂ = 20:1
This is described by the Henderson-Hasselbalch equation:
pH = pKa + log([HCO₃⁻] / [dissolved CO₂]) pH = 6.1 + log(25/1.25) = 6.1 + log(20) = 7.4
Buffer value (β) of bicarbonate in plasma = 55.6 mmol/L
The balance diagram below illustrates how changes in this ratio shift pH:
The teeter-totter: when HCO₃⁻/dCO₂ = 20:1 → pH 7.4 (normal). If the ratio rises → alkalosis; if it falls → acidosis. — Tietz Textbook of Laboratory Medicine
B. Phosphate Buffer System
- At pH 7.4: HPO₄²⁻ / H₂PO₄⁻ = 4:1 (pKa = 6.8)
- Reactions:
- HPO₄²⁻ + H⁺ → H₂PO₄⁻
- H₂PO₄⁻ + OH⁻ → HPO₄²⁻ + H₂O
- Accounts for ~5% of non-bicarbonate buffer value of plasma
- Most important in urine — titrates and excretes acids in renal tubules
- Organic phosphate (2,3-DPG in RBCs) accounts for ~16% of non-bicarbonate buffering in erythrocytes
C. Plasma Protein Buffer System (especially albumin)
- Non-bicarbonate buffer value of plasma ≈ 7.7 mmol/L
- Proteins (>90% albumin) dominate non-bicarbonate buffering
- Imidazole groups of histidine residues are key proton acceptors/donors
D. Hemoglobin Buffer System
- Non-bicarbonate buffer value of erythrocyte fluid ≈ 63 mmol/L
- Hb accounts for ~53 mmol/L of this (major intracellular buffer)
- Critical because Hb carries CO₂ and buffers H⁺ generated in tissues
E. Intracellular Buffers
- Negatively charged phosphates and proteins inside cells
- Bone carbonate and phosphate also contribute (especially in chronic acidosis)
4. Respiratory Regulation (Fast, Minutes)
The lungs regulate PCO₂ (the acid component):
- Acidosis → chemoreceptors detect ↑H⁺ → hyperventilation → ↑CO₂ exhaled → PCO₂ falls → pH rises
- Alkalosis → hypoventilation → CO₂ retained → PCO₂ rises → pH falls
"The pH of plasma can be considered a function of two independent variables: (1) PCO₂, regulated by the lungs; and (2) the concentration of titratable base (HCO₃⁻), regulated by the kidneys." — Tietz Textbook of Laboratory Medicine
5. Renal Regulation (Slow but Powerful, Hours–Days)
The kidneys regulate [HCO₃⁻] (the base component). Renal tubule cells contain active carbonic anhydrase, enabling H⁺ and HCO₃⁻ production from CO₂.
| Condition | Kidney Response |
|---|---|
| Acidosis | ↑H⁺ secretion into tubular fluid (exchanged for Na⁺) + ↑HCO₃⁻ reabsorption → HCO₃⁻ returned to blood |
| Alkalosis | ↓H⁺ secretion + ↓HCO₃⁻ reabsorption → HCO₃⁻ lost in urine |
For each H⁺ secreted in response to acidosis, one Na⁺ and one HCO₃⁻ are added to the blood.
— Ganong's Review of Medical Physiology
6. Compensation vs. Correction
| Disturbance | Primary Change | Compensation |
|---|---|---|
| Metabolic acidosis | ↓HCO₃⁻ | Hyperventilation → ↓PCO₂ |
| Metabolic alkalosis | ↑HCO₃⁻ | Hypoventilation → ↑PCO₂ |
| Respiratory acidosis | ↑PCO₂ | Kidneys retain HCO₃⁻ (hours-days) |
| Respiratory alkalosis | ↓PCO₂ | Kidneys excrete HCO₃⁻ (hours-days) |
Compensatory mechanisms restore the HCO₃⁻/CO₂ ratio toward 20:1 without necessarily correcting the underlying cause. Importantly, the body never overcompensates — if pH is overcorrected, a second primary disturbance must be present.
7. Clinical Definitions Summary
| Term | Meaning |
|---|---|
| Acidemia | Arterial pH < 7.35 |
| Alkalemia | Arterial pH > 7.45 |
| Acidosis | Pathologic process tending to lower pH |
| Alkalosis | Pathologic process tending to raise pH |
| Mixed disorder | Two simultaneous primary processes |
Key Takeaway
pH regulation is an integrated system. Chemical buffers provide the immediate defense; the lungs respond within minutes by adjusting CO₂; and the kidneys provide the most sustained and powerful correction over hours to days by independently controlling plasma HCO₃⁻ concentration. All three systems act together to keep arterial pH within the narrow range of 7.35–7.45.
Sources: Medical Physiology (Boron & Boulpaep); Tietz Textbook of Laboratory Medicine, 7th Ed.; Ganong's Review of Medical Physiology, 26th Ed.
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