Biological buffers are chemical systems that help living organisms keep pH within a narrow, safe range. This matters because enzymes, proteins, membranes, and metabolic reactions are highly sensitive to hydrogen ion concentration. In human blood, pH is normally kept near 7.35 to 7.45 even though cells constantly produce acidic waste.
Without buffers, small additions of acid or base could quickly disrupt normal body function.
The most important blood buffer is the bicarbonate system, which links dissolved carbon dioxide, carbonic acid, bicarbonate ions, and hydrogen ions. When acid is added, bicarbonate can bind H+ to form carbonic acid, which can become CO2 and water. When base is added, carbonic acid can release H+ to replace what was removed.
The lungs and kidneys support this buffer by controlling CO2 removal and bicarbonate concentration.
Key Facts
- A buffer resists pH change by using a weak acid and its conjugate base.
- Bicarbonate buffer reaction: CO2 + H2O ⇌ H2CO3 ⇌ H+ + HCO3-.
- Henderson-Hasselbalch equation: pH = pKa + log([A-]/[HA]).
- For blood bicarbonate buffer: pH = 6.1 + log([HCO3-]/(0.03 × PCO2)).
- Normal arterial blood pH is about 7.35 to 7.45.
- Higher CO2 shifts the bicarbonate system toward more H+, lowering pH.
Vocabulary
- Buffer
- A buffer is a solution that resists large changes in pH when small amounts of acid or base are added.
- Bicarbonate
- Bicarbonate, HCO3-, is a weak base that helps neutralize excess hydrogen ions in blood.
- Carbonic acid
- Carbonic acid, H2CO3, is a weak acid formed when carbon dioxide dissolves in water.
- pH
- pH is a measure of hydrogen ion concentration, with lower pH meaning a more acidic solution.
- Henderson-Hasselbalch equation
- The Henderson-Hasselbalch equation relates pH to the pKa and the ratio of conjugate base to weak acid in a buffer.
Common Mistakes to Avoid
- Thinking buffers stop pH from changing completely is wrong because buffers only reduce pH change within a limited capacity.
- Confusing acid with base in the bicarbonate system is wrong because H2CO3 donates H+ while HCO3- accepts H+.
- Ignoring CO2 in blood pH calculations is wrong because dissolved CO2 acts as the acid side of the bicarbonate buffer equation.
- Using concentrations without checking the ratio is wrong because buffer pH depends mainly on [A-]/[HA], not just the absolute amount of one component.
Practice Questions
- 1 A buffer has pKa = 6.1, [HCO3-] = 24 mM, and the acid term equals 1.2 mM. Use pH = pKa + log([base]/[acid]) to calculate the pH.
- 2 In a blood sample, [HCO3-] = 18 mM and PCO2 = 40 mmHg. Use pH = 6.1 + log([HCO3-]/(0.03 × PCO2)) to estimate the pH.
- 3 Explain why rapid breathing can raise blood pH using the bicarbonate buffer equation.