Acid-base chemistry is the silent infrastructure running beneath every biochemistry topic you will ever study — enzyme kinetics, amino acid behaviour, oxygen delivery, and drug pharmacology all depend on understanding how protons move and how biological systems resist those movements. Master the Henderson-Hasselbalch equation and physiological buffering once, and you gain a tool that explains arterial blood gases, pKa-based drug absorption, and isoelectric points simultaneously.
1 pH and the Brønsted-Lowry Model
The pH of a solution is a compact numerical shorthand for how many free protons (H⁺) are dissolved in it at any given moment. Mathematically, it is defined as the negative base-10 logarithm of the molar proton concentration:
pH = –log₁₀[H⁺]
Because this is a logarithmic scale, a drop of one pH unit corresponds to a ten-fold increase in proton concentration. Pure water at 25 °C has a proton concentration of 10⁻⁷ mol/L, giving a pH of 7.0 — the reference point we call “neutral.” The pH of healthy arterial blood is tightly maintained between 7.35 and 7.45. Even a shift to pH 7.0 (still barely acidic on an absolute scale) is life-threatening — a reminder of just how biologically sensitive proton concentrations are.
Under the Brønsted-Lowry framework, an acid is any species capable of donating a proton, and a base is any species that accepts one. Every acid produces a conjugate base after it loses its proton, and every base yields a conjugate acid after it accepts one. These paired species — known as a conjugate acid-base pair — are the cornerstone of buffering chemistry:
HA ⇌ H⁺ + A⁻
Here, HA is the weak acid and A⁻ is its conjugate base. Acetic acid and acetate are a textbook example: CH₃COOH ⇌ H⁺ + CH₃COO⁻. The ammonium ion and ammonia are another: NH₄⁺ ⇌ H⁺ + NH₃. What makes these pairs powerful biologically is not their individual identities — it is their dynamic equilibrium, always ready to absorb or release protons in response to environmental changes.
2 Weak Acids, Ka, and pKa
A strong acid like HCl dissociates completely in water — it donates every proton it has, flooding the solution with H⁺. A weak acid, by contrast, dissociates only partially. The position of this equilibrium is captured by the acid dissociation constant Ka:
Ka = [H⁺][A⁻] / [HA]
A larger Ka means more complete dissociation and therefore a stronger acid. Because Ka values span many orders of magnitude, biochemists work with the pKa — the negative logarithm of Ka — in exactly the same way pH compresses proton concentrations:
pKa = –log Ka
A low pKa indicates a strong(er) acid; a high pKa indicates a weak(er) one. The pKa is not merely a mathematical convenience — it has a direct physical meaning: the pKa equals the pH at which exactly half the acid molecules are dissociated, meaning [A⁻] = [HA]. This 50:50 point is the sweet spot for buffering, as we will see below.
| Acid / Group | Conjugate Base | pKa | Biological Relevance |
|---|---|---|---|
| Acetic acid (CH₃COOH) | Acetate (CH₃COO⁻) | 4.8 | Model weak acid; lab buffer |
| Carbonic acid (H₂CO₃) | Bicarbonate (HCO₃⁻) | 6.1 | Blood/extracellular fluid buffer |
| Dihydrogen phosphate (H₂PO₄⁻) | Monohydrogen phosphate (HPO₄²⁻) | 6.82 | Intracellular buffer |
| α-Carboxyl group (amino acids) | Carboxylate (–COO⁻) | ~2.0–2.9 | Protein charge; isoelectric point |
| α-Amino group (amino acids) | Free amine (–NH₂) | ~8.8–10.8 | Protein charge; isoelectric point |
| Imidazolium (Histidine side chain) | Imidazole | ~6.0 | Key active-site buffer in enzymes |
3 The Henderson-Hasselbalch Equation — Derivation and Meaning
The Henderson-Hasselbalch equation is the cornerstone of biological acid-base reasoning — a single expression linking pH, pKa, and the relative concentrations of a weak acid and its conjugate base. It connects pH, pKa, and the ratio of base to acid concentrations in a form that is both mathematically simple and physiologically interpretable. Here is how it emerges from first principles.
Start from the Ka expression and solve for [H⁺]:
[H⁺] = Ka × [HA] / [A⁻]
Take the negative logarithm of both sides. The left side becomes pH; the right side splits into –log Ka (= pKa) and –log([HA]/[A⁻]) = +log([A⁻]/[HA]):
pH = pKa + log([A⁻] / [HA])
This is the Henderson-Hasselbalch equation. Reading it intuitively:
- When [A⁻] = [HA] (equal amounts of base and acid), log(1) = 0, so pH = pKa.
- When [A⁻] > [HA], the log term is positive — pH rises above pKa (more basic).
- When [A⁻] < [HA], the log term is negative — pH falls below pKa (more acidic).
A useful rule of thumb: at a 10:1 ratio of base to acid, pH = pKa + 1; at a 1:10 ratio, pH = pKa – 1. This defines the effective buffering window (see Section 4).
“More Base = Higher pH, More Acid = Lower pH”
Base ↑: log([A⁻]/[HA]) is positive → pH > pKa. Acid ↑: log([A⁻]/[HA]) is negative → pH < pKa. Equal: log(1) = 0 → pH = pKa exactly. Remember: the base form is always the numerator in the log term.
4 Buffers — How They Work and Why They Matter
A buffer is a solution capable of absorbing added acid or base without allowing more than a minimal shift in pH. It does this by containing both members of a conjugate acid-base pair in significant concentrations, so it can “absorb” protons or “donate” protons as needed.
When you add a strong base (OH⁻) to an acetate buffer, the conjugate acid mops it up:
CH₃COOH + OH⁻ → CH₃COO⁻ + H₂O
When you add a strong acid (H⁺), the conjugate base neutralises it:
CH₃COO⁻ + H⁺ → CH₃COOH
In both cases, a strong acid or base is converted into water and the weak species — a far less dramatic change in [H⁺] than if no buffer were present.
Buffering capacity — a measure of how much acid or base a solution can absorb before pH shifts significantly — peaks when pH = pKa, and the system stays useful across roughly one pH unit either side of the pKa. Outside this window, the buffer is overwhelmed: there is too much of one species and too little of the other to maintain the equilibrium effectively.
A buffer is most effective when pH = pKa (50:50 acid:base ratio). It loses practical effectiveness outside pKa ± 1 pH unit. For a buffer to work in blood (pH 7.4), its pKa should ideally fall between 6.4 and 8.4.
The bicarbonate system (pKa 6.1) technically falls just outside optimal range — yet it is the dominant blood buffer because its components (CO₂ and HCO₃⁻) are independently regulated by the lungs and kidneys, giving the body far more control than a simple chemical buffer could provide.
5 Physiological Buffer Systems
The body runs three major buffer systems, each operating in different compartments and over different timescales.
Bicarbonate Buffer System Extracellular Fluid / Blood
CO₂ + H₂O ⇌ H₂CO₃ ⇌ H⁺ + HCO₃⁻ (carbonic anhydrase)
The dominant buffer of blood plasma, with a normal [HCO₃⁻]/[H₂CO₃] ratio of about 20:1 at pH 7.4. Despite its pKa of 6.1 (seemingly far from 7.4), the system is extraordinarily powerful because CO₂ is continuously exhaled by the lungs (controlling the acid side) and HCO₃⁻ is reabsorbed or excreted by the kidney (controlling the base side). This dual physiological regulation makes the bicarbonate system the most clinically important buffer in the body.
Phosphate Buffer System Intracellular Fluid / Urine
H₂PO₄⁻ ⇌ H⁺ + HPO₄²⁻ (pKa = 6.82)
Phosphate’s pKa of 6.82 sits closer to intracellular pH (~7.0–7.2), making it the primary intracellular buffer. It is also highly effective in renal tubular fluid, where it accepts secreted protons and contributes to urinary acidification. In plasma, its absolute concentration is too low to buffer effectively — but inside cells and in urine, it is critical.
Protein Buffers (Haemoglobin, Plasma Proteins) Blood / Intracellular
Protein–NH₃⁺ ⇌ H⁺ + Protein–NH₂
Proteins contribute substantially to total blood buffering. Haemoglobin is the most important protein buffer in blood — it accounts for more than half of the non-bicarbonate buffering capacity of whole blood. Histidine residues are particularly important: their imidazole side chains have a pKa near 6.0, which sits within the physiological range and allows histidine-containing proteins to act as buffers. When haemoglobin releases its oxygen it simultaneously gains an increased appetite for protons — deoxyhaemoglobin binds H⁺ more avidly than its oxygenated counterpart — and this shift underpins the Bohr effect (see Clinical Pearl below).
“Big Cats Prowl”
Bicarbonate — blood/extracellular, lung + kidney regulated. Carbonate is close enough — don’t confuse it with phosphate! Phosphate — intracellular + urine. Proteins (Haemoglobin) — blood and cells. Three systems, three compartments, three timescales.
In metabolically active tissues, CO₂ production is high. Carbonic anhydrase converts CO₂ to carbonic acid, which dissociates into H⁺ and HCO₃⁻, lowering local pH. This local acidification has a direct effect on haemoglobin: a lower pH reduces haemoglobin’s affinity for oxygen, causing it to offload O₂ more readily to the tissues — exactly where it is needed. This is the Bohr effect.
The reverse happens in the lungs: CO₂ diffuses out, pH rises, and haemoglobin’s O₂ affinity increases, facilitating loading. The bicarbonate buffer system and the Bohr effect are thus inseparable in clinical physiology — understanding one deepens understanding of the other.
6 Titration Curves and the pKa in Practice
A titration curve plots pH on the y-axis against equivalents of base added on the x-axis as a weak acid is progressively neutralised. The shape is characteristic: a sigmoid (S-shaped) curve with a flat plateau region where pH changes very little per unit of base added. This plateau is the buffer region, and its centre — the inflection point of the curve — corresponds exactly to the pKa.
Consider the titration of acetic acid (pKa 4.8): as NaOH is added, the pH barely budges between approximately 3.8 and 5.8. Then, as the supply of acetic acid becomes depleted, the curve steepens sharply — the buffer is exhausted. The symmetry of the plateau around the pKa is a direct consequence of the Henderson-Hasselbalch equation.
Amino acids yield more complex titration curves because they carry two titratable groups — the α-carboxyl (pKa₁ ≈ 2.0–2.9) and the α-amino group (pKa₂ ≈ 8.8–10.8) — producing two distinct buffer regions and two inflection points. Amino acids with ionisable side chains add a third titratable group and a third inflection point.
7 Amino Acid Ionisation, Zwitterions, and the Isoelectric Point
At very low pH, both the carboxyl and amino groups of a free amino acid are fully protonated: the carboxyl carries –COOH and the amino group carries –NH₃⁺, giving the molecule an overall charge of +1. As pH increases, the carboxyl group (lower pKa) is the first to lose its proton, producing –COO⁻. At this point the molecule carries both –COO⁻ and –NH₃⁺ simultaneously — this is the zwitterion form, electrically neutral with a net charge of zero.
This neutral charge state — where positive and negative contributions cancel perfectly — defines the isoelectric point (pI). For amino acids with only two titratable groups (like alanine), pI is simply the average of the two pKa values:
pI = (pKa₁ + pKa₂) / 2
For alanine: pI = (2.3 + 9.1) / 2 = 5.7. At pH values below the pI, the amino acid carries a net positive charge; above the pI, it carries a net negative charge.
| pH Condition | Dominant Species (Alanine) | Net Charge | Henderson-Hasselbalch Implication |
|---|---|---|---|
| pH < pKa₁ (~2.3) | –COOH, –NH₃⁺ (Form I) | +1 | Both groups fully protonated |
| pH = pKa₁ (2.3) | 50% Form I, 50% Form II | +0.5 avg | [–COOH] = [–COO⁻] |
| pH = pI (5.7) | –COO⁻, –NH₃⁺ (Form II, zwitterion) | 0 | Electrically neutral |
| pH = pKa₂ (9.1) | 50% Form II, 50% Form III | –0.5 avg | [–NH₃⁺] = [–NH₂] |
| pH > pKa₂ (~9.1) | –COO⁻, –NH₂ (Form III) | –1 | Both groups fully deprotonated |
“Acid has Low pI, Base has High pI”
Acidic AAs (Asp, Glu): pI = average of the two lowest pKa values (pKa₁ and pKa_side chain) → pI is low (~3). Basic AAs (Lys, Arg, His): pI = average of the two highest pKa values (pKa_side chain and pKa₂) → pI is high (~8–11). Neutral AAs: pI = average of pKa₁ and pKa₂ → pI near physiologic range (~5–6).
8 Clinical Application — Drug Absorption and the Henderson-Hasselbalch Equation
The vast majority of pharmaceutical compounds behave as either weak acids or weak bases in solution, and whether they carry an ionic charge at a given pH determines whether they can cross a lipid bilayer. Only the uncharged (non-ionised) form of a drug is lipid-soluble enough to diffuse across cell membranes. By applying the Henderson-Hasselbalch equation, you can calculate the exact fraction of a drug that exists in its uncharged versus ionised form at any given compartment pH.
For a weak acid drug (e.g. aspirin, pKa 3.5):
HA ⇌ H⁺ + A⁻
In the stomach (pH 1.5), the Henderson-Hasselbalch equation gives log([A⁻]/[HA]) = 1.5 – 3.5 = –2.0, so [HA]/[A⁻] = 100:1. Aspirin is overwhelmingly in its uncharged protonated form (–COOH) at gastric pH — meaning it is well absorbed from the stomach.
For a weak base drug (e.g. morphine), the protonated form (BH⁺) is the charged species and the free base (B) is uncharged. At the alkaline pH of the small intestine, weak bases are more likely to exist in their uncharged form and are better absorbed there than in the acidic stomach.
When a drug crosses a membrane into a compartment with a different pH, it may become ionised and “trapped” — unable to diffuse back out. This is clinically exploited in aspirin overdose management: alkalinising the urine (with IV sodium bicarbonate) converts aspirin (a weak acid) into its ionised A⁻ form in the renal tubule, preventing reabsorption and accelerating excretion. The Henderson-Hasselbalch equation is the pharmacokinetic engine driving this intervention.
9 High-Yield Exam Summary
pH = –log[H⁺]. Physiological blood pH is 7.35–7.45. A change of 1 pH unit = 10-fold change in [H⁺].
Henderson-Hasselbalch: pH = pKa + log([A⁻]/[HA]). At pH = pKa, exactly 50% dissociation. Effective buffering range = pKa ± 1.
pKa definition: pH at which 50% of the weak acid is dissociated ([HA] = [A⁻]). Stronger acid = lower pKa.
Physiological buffers: Bicarbonate (pKa 6.1, blood — regulated by lungs/kidneys); Phosphate (pKa 6.82, intracellular/urine); Proteins/Haemoglobin (histidine imidazole, pKa ~6.0).
Bicarbonate system: pH = 6.1 + log([HCO₃⁻] / 0.03 × pCO₂). Respiratory acidosis = ↑CO₂ → ↓pH. Metabolic acidosis = ↓HCO₃⁻ → ↓pH.
Amino acid pI: For neutral AAs: pI = (pKa₁ + pKa₂)/2. For acidic AAs: pI = average of the two lower pKa values. For basic AAs: pI = average of the two higher pKa values.
Drug absorption: Uncharged form crosses membranes. Weak acids absorbed better in acidic environments (stomach). Weak bases absorbed better in alkaline environments (small intestine). Ion trapping used therapeutically in overdose management.
Bohr effect: ↓pH (↑H⁺, ↑CO₂) → ↓O₂ affinity of haemoglobin → right shift of dissociation curve → O₂ released to tissues. Mediated by proton binding to histidine residues of deoxyhaemoglobin.
10 Mnemonic Summary Wall
“More Base = Higher pH, More Acid = Lower pH”
Base ↑: log([A⁻]/[HA]) is positive → pH > pKa. Acid ↑: log([A⁻]/[HA]) is negative → pH < pKa. Equal: log(1) = 0 → pH = pKa exactly. The base form is always the numerator in the log term.
“Big Cats Prowl”
Bicarbonate — blood/extracellular, lung + kidney regulated (pKa 6.1). Carboxylate (from Phosphate) — intracellular + urine (pKa 6.82). Proteins/Haemoglobin — blood and cells (His imidazole, pKa ~6.0).
“Acid has Low pI, Base has High pI”
Acidic AAs (Asp, Glu): pI = average of two lowest pKa values → pI ~3. Basic AAs (Lys, Arg, His): pI = average of two highest pKa values → pI ~8–11. Neutral AAs: pI = (pKa₁ + pKa₂)/2 → pI ~5–6.
“Acids Absorbed in the Stomach, Bases absorbed Below (intestine)”
Weak Acids (e.g. aspirin, warfarin) → uncharged in acidic stomach → Absorbed in stomach. Weak Bases (e.g. morphine, codeine) → uncharged in alkaline intestine → absorbed Below (small intestine). Non-ionised = membrane-permeable = absorbed.
Kennelly, P. J., Botham, K. M., Weil, P. A., & Murray, R. K. (2018). Harper’s illustrated biochemistry (32nd ed., pp. 1–12). McGraw-Hill Education.
Hames, B. D., & Hooper, N. M. (2011). BIOS instant notes in biochemistry (3rd ed., pp. 28–34). Taylor & Francis.
The content on this page is intended for educational purposes only and is not a substitute for professional medical advice, clinical judgement, or the guidance of a qualified healthcare provider. Always refer to current clinical guidelines and consult appropriate sources before applying information in a patient care setting.