Buffer Chemistry in Blood and Physiology — Henderson–Hasselbalch at Work
You already know the Henderson–Hasselbalch equation from acetate and phosphate buffer problems. The same equation, unchanged, is the one used to describe the pH of blood. That makes it one of the clearest bridges between a syllabus formula and a real physical system, and it is a favourite context for application questions in IIT-JAM, GATE, CSIR-NET and CUET-PG papers.
This article works the blood bicarbonate buffer properly — including the two constants that look strange the first time you see them — and then says plainly where the textbook formula stops describing the real system. That last part matters: the equation is exact algebra applied to an approximate model, and knowing the approximations is what separates a full-marks answer from a half-marks one.
This is a chemistry explanation written for exam preparation. It is not clinical guidance and must never be used to interpret anyone's medical results.
The formula you already have
Written for the bicarbonate system, the conjugate pair is carbonic acid / hydrogen carbonate, H2CO3 / HCO3−:
What each term means, and why the numbers look odd
| Symbol | Meaning | Typical unit |
|---|---|---|
| pKa′ = 6.1 | The apparent acid dissociation constant of the CO2/HCO3− pair in plasma at 37 °C | dimensionless |
| [HCO3−] | Hydrogen carbonate concentration in plasma | mmol/L |
| pCO2 | Partial pressure of carbon dioxide in the gas phase in equilibrium with the plasma | mmHg |
| 0.03 | Solubility coefficient of CO2 in plasma at 37 °C — a Henry's-law constant that converts a pressure into a dissolved concentration | mmol·L−1·mmHg−1 |
Two things surprise students here.
First, why 6.1 and not 3.6? The thermodynamic pKa1 of true carbonic acid is around 3.6. But in solution only a very small fraction of the dissolved carbon dioxide is actually hydrated to H2CO3; the overwhelming majority stays as dissolved CO2. If you define the acid form as all the dissolved CO2 — which is what you can actually measure — the denominator of the ratio becomes much larger, and the effective constant shifts to about 6.1. It is the same equilibrium, described with a different (and more useful) definition of the acid. The 6.1 is an empirical constant for plasma at 37 °C, not a pure thermodynamic quantity.
Second, why a pressure inside a concentration ratio? Because Henry's law lets you swap them. Dissolved [CO2] = 0.03 × pCO2, so a pCO2 of 40 mmHg corresponds to 1.2 mmol/L of dissolved CO2. Using pressure is simply more convenient, because pressure is what a blood-gas instrument measures.
Worked example 1 — the normal value, derived rather than memorised
Take the standard reference figures: [HCO3−] = 24 mmol/L and pCO2 = 40 mmHg.
Step 1 — dissolved CO2: 0.03 × 40 = 1.2 mmol/L
Step 2 — the ratio: 24 ÷ 1.2 = 20.0
Step 3 — the logarithm: log10(20.0) = 1.301
Step 4 — add: pH = 6.1 + 1.301 = 7.40
The famous 20 : 1 base-to-acid ratio is not a separate fact to learn — it is what you get when you put the reference numbers into the equation.
Worked example 2 — what happens when CO2 is not cleared
Suppose gas exchange is impaired so that pCO2 rises to 60 mmHg while [HCO3−] has not yet changed from 24 mmol/L.
0.03 × 60 = 1.8 mmol/L
24 ÷ 1.8 = 13.33
log10(13.33) = 1.125
pH = 6.1 + 1.125 = 7.22
Raising the acid side of the pair by 50% drops the pH by 0.18 units. Note how small the pH change is compared with the concentration change — that is buffering, and the logarithm is why.
Worked example 3 — restoring the ratio, not the concentrations
Now suppose pCO2 stays at 60 mmHg but [HCO3−] is raised to 31 mmol/L.
0.03 × 60 = 1.8 mmol/L
31 ÷ 1.8 = 17.22
log10(17.22) = 1.236
pH = 6.1 + 1.236 = 7.34
The pH is pulled back most of the way towards 7.40 even though both concentrations are now abnormal. This is the single most important chemical lesson in the whole topic: the Henderson–Hasselbalch equation fixes pH from the ratio, not from the absolute amounts. Two very different solutions with the same ratio have the same pH — though, as the buffer-capacity argument below shows, they are not equally good at resisting the next disturbance.
Why a pKa of 6.1 is not a design error
A closed buffer works best when the working pH is within about one unit of its pKa, because that is where the ratio is near 1 and the buffer can absorb acid or base equally well. At pH 7.40 with pKa′ = 6.1 the ratio is 20 : 1, which is far outside that window. On paper, the bicarbonate system looks like a badly chosen buffer.
The resolution is that it is not a closed system. The acid member of the pair is a gas that can be removed by ventilation, and the base member can be adjusted by the kidneys. A buffer whose components can be independently added and removed is far more powerful than the closed-flask buffer of a titration problem, and the ordinary buffer-capacity formula from your notes simply does not apply to it. This is a genuinely important idea in physical chemistry: the "open buffer" is a different system, not a special case.
Alongside it, the body also uses buffers that are closer to closed:
| Buffer pair | Approximate pKa | Where it dominates |
|---|---|---|
| CO2 / HCO3− | 6.1 (apparent, plasma, 37 °C) | Extracellular fluid; open to the lungs and kidneys |
| H2PO4− / HPO42− | about 7.2 | Inside cells and in urine, where phosphate is concentrated |
| Protein side chains (notably imidazole of histidine) | near neutral | Inside cells and in red blood cells; large total capacity |
The phosphate pair has a pKa almost exactly at the working pH, which is why it is the textbook example of a well-matched buffer — but there is far less phosphate than bicarbonate in extracellular fluid, so its total capacity there is small. Capacity depends on both how close the pKa is and how much buffer is present.
Where the simple formula stops being valid
- Concentrations are not activities. Henderson–Hasselbalch in the form you learnt uses concentrations. Plasma has an ionic strength of roughly 0.15 mol/L, so activity coefficients are well below 1. The value 6.1 already absorbs those non-ideality corrections for plasma at 37 °C — which is exactly why it is called an apparent constant and why you must not use it for a bicarbonate solution made in a beaker at 25 °C.
- Both constants are temperature-specific. pKa′ = 6.1 and the solubility coefficient 0.03 are 37 °C values. Change the temperature and both move.
- The system is open. Any calculation that treats the pair as a fixed total amount of buffer — the standard buffer-capacity derivation, for instance — is structurally wrong here, because the acid form is continuously produced by metabolism and continuously vented as a gas.
- The equation gives pH, never the cause. Two completely different situations can give the same pH, as example 3 shows. pH alone does not tell you which member of the pair moved. That is a chemistry point as much as a physiological one: one equation with two unknowns needs a second measurement.
- The units must match the constant. The 0.03 coefficient is written for pCO2 in mmHg. If you work in kPa the coefficient is a different number (pressure units convert, so the constant must too). Silently mixing the two is the single commonest arithmetic slip in this calculation.
- Real values are measured, not computed. In practice the pH and pCO2 are measured and the bicarbonate is calculated from them — not the other way round. Treating a computed number as an independent measurement is circular.
Where this appears in exams
| Exam | How it is usually asked |
|---|---|
| IIT-JAM | Ionic equilibrium: compute pH from a buffer ratio; identify which pair buffers best at a stated pH |
| CUET-PG | Direct Henderson–Hasselbalch substitution; effect of changing one component of the pair |
| GATE | Buffer capacity, choice of buffer for a working pH, and non-ideality/activity corrections |
| CSIR-NET | Open vs closed buffer reasoning, and biochemical contexts where a near-neutral pKa matters |
Do the ratio arithmetic without slips. The free Buffer (Henderson–Hasselbalch) calculator takes the pKa and the two concentrations and returns the pH, so you can test what a change in either member of the pair does before you attempt it by hand in an exam.
Open the Buffer (Henderson–Hasselbalch) Calculator →Preparing for IIT-JAM, GATE, CSIR-NET or CUET-PG? ABC Chemistry runs dedicated competitive-exam batches at its coaching centre and online for students across India — details at abcchemistry.in.