Overview
This lesson explains acids, bases, and buffer systems in the body, focusing on how the body maintains low hydrogen ion concentrations to prevent protein damage and preserve normal pH levels.
Ions in the Body
- The body contains numerous charged atoms (ions) including sodium, potassium, calcium, hydrogen, bicarbonate, chloride, and magnesium.
- Sodium concentration in blood is approximately 140 millimolar; hydrogen concentration is only 0.00004 millimolar.
- Each ion serves specific functions: sodium and potassium enable nerve conduction; calcium supports muscle contraction.
- Hydrogen ions are extremely reactive despite their low concentration, requiring careful regulation.
Hydrogen Ion Reactivity
- Hydrogen is the smallest element on the periodic table; in ionized form it consists of just a proton.
- Compared to larger ions like sodium and potassium, hydrogen has a massive charge-to-size ratio.
- This high reactivity causes hydrogen ions to steal electrons from proteins, potentially damaging structural and functional components.
- Maintaining low hydrogen concentrations protects enzymes and other critical proteins throughout the body.
Acids and Bases: Definitions
- Acids donate hydrogen ions (protons) to solutions; represented as HA → H⁺ + A⁻.
- Bases absorb or mop up hydrogen ions from solutions.
- When an acid donates a hydrogen ion, the remaining molecule (A⁻) is called the conjugate base.
- The conjugate base can theoretically bind hydrogen ions to reform the original acid.
Physiological Acids
- Hydrochloric acid (HCl): Strong acid with pKa around –6; dissociates completely into H⁺ and Cl⁻.
- Dihydrogen phosphate (H₂PO₄⁻): Produces hydrogen phosphate (HPO₄²⁻) when dissociated.
- Carbonic acid (H₂CO₃): Dissociates into H⁺ and bicarbonate (HCO₃⁻).
- Amino acids: Contain both a carboxyl terminus (C-terminus) and amino terminus (N-terminus) that can donate or accept protons.
- Additional acids include lactic acid, pyruvic acid, and fatty acids.
pH and pKa Relationship
- pH measures hydrogen ion concentration in a solution; normal body pH is 7.4.
- pKa indicates how likely a specific acid is to donate its hydrogen ion.
- When pKa equals pH, the acid-base system is balanced and acts as an effective buffer.
- If pKa < pH, the system produces more hydrogen ions (shifts toward dissociation).
- If pKa > pH, the system mops up more hydrogen ions (shifts toward acid formation).
Strong vs. Weak Acids
- Strong acids (e.g., HCl) have very low pKa values and dissociate completely in one direction.
- Weak acids have pKa values closer to physiological pH and can reversibly donate or accept protons.
- Only weak acids can function as buffers because they operate bidirectionally.
Buffers: Function and Importance
- Buffers resist drastic changes in pH by releasing protons when levels drop or binding them when levels rise.
- A good buffer has pKa as close as possible to the solution's pH.
- Buffers maintain stable pH despite metabolic processes that produce or consume hydrogen ions.
Three Main Chemical Buffer Systems
| Buffer System | Key Components | Primary Location | pKa | Characteristics |
|---|
| Phosphate | H₂PO₄⁻ / HPO₄²⁻ | Intracellular fluid, renal tubules | ~6.8 | Closest pKa to body pH; concentrated inside cells and in kidneys where pH drops |
| Bicarbonate | H₂CO₃ / HCO₃⁻ | Extracellular fluid, blood | ~6.1 | Most important buffer; volatile (produces CO₂ gas); regulated by lungs and kidneys |
| Protein (Amino Acid) | Amino acid carboxyl/amine groups | Intracellular fluid | Varies | Most abundant buffer; accounts for 60–70% of total buffering; dual ionizable groups |
The Bicarbonate Buffer System
- Carbon dioxide (CO₂) from cellular metabolism dissolves in blood water to form carbonic acid (H₂CO₃).
- Carbonic acid dissociates reversibly: CO₂ + H₂O ↔ H₂CO₃ ↔ H⁺ + HCO₃⁻.
- Increasing CO₂ (e.g., holding breath) raises blood acidity by producing more H⁺.
- Decreasing CO₂ (e.g., hyperventilating) lowers blood acidity by removing H⁺ precursors.
- Lungs provide short-term pH control by adjusting breathing rate to change CO₂ levels.
- Kidneys provide long-term pH control (hours to days) by reabsorbing or excreting bicarbonate.
- This dual regulation makes bicarbonate the most important buffer system despite not being the most abundant.
Phosphate Buffer Details
- Most effective intracellularly where phosphate concentration is highest.
- Less abundant in extracellular fluid, limiting its systemic buffering capacity.
- Becomes especially important in renal tubules where pH decreases and phosphate concentrations increase.
- The pKa of approximately 6.8 makes it theoretically the best buffer among the three systems.
Protein Buffer Details
- Amino acids have two ionizable groups: carboxyl end (C-terminus) and amino end (N-terminus).
- Excess acid: Amino end (NH₂) binds H⁺ to form NH₃⁺, mopping up excess hydrogen ions.
- Insufficient acid: Carboxyl end (COOH) donates H⁺ to form COO⁻, releasing hydrogen ions.
- Some amino acids have additional acidic or basic functional groups, increasing buffering capacity.
- Proteins are everywhere in the body, providing the largest total buffering capacity.
Key Terms & Definitions
- Ion: Charged atom or element (e.g., Na⁺, K⁺, Ca²⁺, H⁺).
- Proton: Synonymous with hydrogen ion (H⁺); the ionized form of hydrogen.
- Conjugate Base: The molecule remaining after an acid donates a hydrogen ion.
- Strong Acid: Completely dissociates into hydrogen ions and conjugate base; cannot buffer.
- Weak Acid: Partially dissociates; can reversibly donate or accept protons; acts as buffer.
- Volatile Buffer: A buffer system involving a gas component that can be exhaled (bicarbonate system).
- Nonvolatile Buffers: Buffer systems without gaseous components (phosphate and protein systems).