Acid-Base Balance and Blood pH

Why acid-base balance matters

• Enzymes are pH machines. Nearly every intracellular enzyme has a sharply peaked pH optimum because catalysis depends on the protonation state of active-site residues — the imidazole of histidine (pKa ~6.0) is the single most pH-sensitive side chain and the dominant intracellular buffer. A blood pH swing from 7.4 to 7.0 measurably slows glycolytic flux and cardiac contractility; below 6.9, myocardial depression and arrhythmia become life-threatening.
• It governs oxygen delivery. Through the Bohr effect, falling pH (rising H+ and CO2) right-shifts the hemoglobin-oxygen dissociation curve, unloading O2 in acidic, metabolically active tissue; the reverse Haldane effect helps lungs load O2 and dump CO2. Acid-base status and gas transport are one coupled system.
• It shifts potassium. Acidemia drives K+ out of cells (H+/K+ exchange) and can produce dangerous hyperkalemia; alkalemia does the reverse and drops serum K+. A blood gas often has to be read alongside the potassium, because correcting one perturbs the other.
• It is the most-ordered test in critical care. The arterial blood gas — pH, PCO2, PO2, HCO3-, base excess — is drawn millions of times a year worldwide to diagnose respiratory failure, sepsis, diabetic ketoacidosis, poisoning, and shock. Reading it correctly is a core clinical skill built directly on Henderson-Hasselbalch.
• Fixed-acid excretion is a renal specialty. The kidney's ability to make ammonium from glutamine lets it ramp acid excretion several-fold during chronic acidosis (starvation, chronic kidney disease, renal tubular acidosis). When that capacity fails, acid accumulates relentlessly — a defining feature of advanced kidney disease.
• Temperature and altitude move the setpoint. Blood pH rises as blood cools (alphastat regulation keeps histidine charge constant, relevant during hypothermic cardiac surgery), and chronic altitude exposure resets the respiratory drive as the kidney trims bicarbonate to compensate for the alkalosis of hyperventilation.

Common misconceptions

• "Bicarbonate is a weak buffer because its pKa is far from 7.4." On paper a buffer is strongest at its pKa, so bicarbonate's pKa of 6.1 (1.3 units below plasma pH) looks poor. But the system is open: the lungs continuously remove the acid leg (CO2) and the kidneys regenerate the base leg (HCO3-). That independent control of numerator and denominator makes bicarbonate the dominant extracellular buffer despite its pKa — a closed buffer at the same pKa would be far weaker.
• "Blood pH of 7.40 means acid-base status is normal." A normal pH can conceal two opposing primary disorders that cancel out (mixed metabolic acidosis and respiratory alkalosis, for example). You must inspect PCO2 and HCO3- — and the anion gap — to know whether the number is truly normal or a coincidental crossing point.
• "Carbonic acid (H2CO3) is what buffers the blood." H2CO3 is a fleeting intermediate present at only ~1/1000 the concentration of dissolved CO2. The physiologically relevant pair is dissolved CO2 and HCO3-, which is why the clinical Henderson-Hasselbalch equation uses 0.03 × PCO2 (dissolved CO2), not H2CO3, and an apparent pKa of 6.1 that lumps the hydration step in.
• "The kidney corrects acidosis by excreting free H+." Urine pH bottoms out around 4.5, which corresponds to only ~0.03 mmol/L of free H+ — nowhere near the ~70 mmol daily load. The kidney excretes acid almost entirely as buffered protons: titratable acid (mostly phosphate) plus ammonium (NH4+), the latter being the adjustable arm that scales up during chronic acid loads.
• "Respiratory compensation can fully normalize pH." Compensation blunts but never fully corrects a primary disturbance, and it never overcorrects. If a patient with metabolic acidosis has a completely normal pH, suspect a second, superimposed disorder rather than perfect compensation.
• "Giving bicarbonate is the fix for metabolic acidosis." Because the buffer is open, infused HCO3- combines with H+ to generate CO2 that must be exhaled; in a patient who cannot increase ventilation, or in lactic acidosis where the underlying process continues, bicarbonate can paradoxically worsen intracellular acidosis. Treating the cause (restoring perfusion, giving insulin) is usually primary.

How acid-base balance works, step by step

The starting problem is a relentless acid load. Aerobic metabolism produces about 15,000 mmol of CO2 per day — the "volatile acid," because CO2 hydrates to carbonic acid and can be breathed away. On top of that, protein and phospholipid metabolism generates roughly 50 to 100 mmol of nonvolatile "fixed" acid daily (sulfuric acid from sulfur-containing amino acids, phosphoric acid, and organic acids), which the lungs cannot remove and the kidney must handle. Left undefended, this would drive pH lethally low within hours.

Line one — chemical buffers (seconds). The instant H+ appears, buffers absorb it. The dominant extracellular buffer is the bicarbonate/CO2 pair, working around CO2 + H2O ⇌ H2CO3 ⇌ H+ + HCO3-, with the hydration step catalyzed by carbonic anhydrase in red cells and on capillary endothelium. Intracellular and bone buffers (hemoglobin, phosphate, protein histidine, carbonate in bone) share the load. Buffering alone does not remove acid; it only cushions the pH change until the lungs and kidneys can act.

Line two — the lungs (minutes). Central chemoreceptors in the medulla (sensing CSF pH) and peripheral chemoreceptors in the carotid and aortic bodies (sensing PCO2, pH, and PO2) adjust alveolar ventilation. In acidosis, ventilation rises and blows off CO2, lowering the denominator (0.03 × PCO2) of the Henderson-Hasselbalch ratio and pushing pH back up — the deep, rapid Kussmaul respiration of diabetic ketoacidosis can drive PCO2 below 15 mmHg. In alkalosis, ventilation falls and CO2 is retained. The respiratory response is fast but has limited capacity and cannot address the underlying bicarbonate deficit.

Line three — the kidneys (hours to days). The kidney does two jobs. First, it reclaims the ~4,300 mmol of HCO3- filtered daily: proximal tubule cells secrete H+ into the lumen via apical Na+/H+ exchanger NHE3, that H+ combines with filtered HCO3- to form CO2 (via luminal carbonic anhydrase IV), CO2 diffuses in, and HCO3- is regenerated inside the cell and returned to blood through the basolateral Na+/HCO3- cotransporter NBCe1. Roughly 80–90% of filtered bicarbonate is recovered here. Second, the kidney generates new HCO3- to replace what buffered the day's fixed acid: alpha-intercalated cells of the collecting duct pump H+ into urine via an apical H+-ATPase and H+/K+-ATPase, and that secreted H+ is trapped by urinary buffers — titratable acid (chiefly HPO4²⁻ → H2PO4⁻) and, critically, ammonium (NH4+) made from glutamine in the proximal tubule. Every H+ excreted as ammonium or titratable acid returns one new HCO3- to the blood. Ammoniagenesis is the adjustable arm, able to multiply several-fold over days of chronic acidosis.

The controlling equation. All of this collapses into the ratio in the Henderson-Hasselbalch equation, pH = 6.1 + log₁₀([HCO3-] / (0.03 × PCO2)). At normal values (HCO3- = 24 mmol/L, PCO2 = 40 mmHg), the ratio is 24 / 1.2 = 20, log₁₀(20) ≈ 1.30, and pH = 6.1 + 1.30 = 7.40. Because pH tracks the ratio and not the absolutes, the lungs (denominator) and kidneys (numerator) can each move their variable to restore 20:1. Compensation is exactly this: a primary change in one term is partly offset by a directional change in the other.

Respiratory vs metabolic disturbances

Every simple acid-base disorder is classified by which Henderson-Hasselbalch term moved first. A primary change in PCO2 is respiratory; a primary change in HCO3- is metabolic. The compensating organ then moves its term in the same direction to blunt the pH shift.

Disorder	Primary change	pH	Common causes	Compensation
Respiratory acidosis	↑ PCO2 (>45 mmHg)	↓ (<7.35)	COPD, opioid overdose, neuromuscular weakness, airway obstruction	Renal: ↑ HCO3- over days
Respiratory alkalosis	↓ PCO2 (<35 mmHg)	↑ (>7.45)	Anxiety/pain, sepsis, high altitude, salicylate (early), pregnancy	Renal: ↓ HCO3- over days
Metabolic acidosis	↓ HCO3- (<22 mmol/L)	↓ (<7.35)	Lactic acidosis, DKA, toxins, diarrhea, renal tubular acidosis	Respiratory: ↓ PCO2 in minutes (Kussmaul)
Metabolic alkalosis	↑ HCO3- (>26 mmol/L)	↑ (>7.45)	Vomiting, diuretics, hyperaldosteronism, antacid/milk-alkali	Respiratory: ↑ PCO2 (hypoventilation)

The two acidoses feel the same to a chemoreceptor but demand opposite treatments. Metabolic acidosis is further split by the anion gap = Na+ − (Cl- + HCO3-), normally 8–12 mmol/L. A high gap (e.g., the classic mnemonic causes: ketoacidosis, lactate, methanol/ethylene glycol, salicylate, uremia) means unmeasured acid anions were added; a normal gap means bicarbonate was lost and replaced by chloride (hyperchloremic acidosis, as in diarrhea or renal tubular acidosis). Winters' formula (expected PCO2 = 1.5 × HCO3- + 8 ± 2) then checks whether respiratory compensation is appropriate — a measured PCO2 outside that band signals a second, mixed disorder.

Bicarbonate buffer vs other body buffers

Property	Bicarbonate / CO2	Phosphate (HPO4²⁻/H2PO4⁻)	Protein / Hemoglobin (histidine)
pKa	6.1 (apparent)	~6.8	~6.0 (imidazole)
Main location	Extracellular fluid, plasma	Intracellular, urine	Intracellular, red cells
Open or closed?	Open (lung + kidney)	Effectively closed	Closed
Key strength	Independent, physiologic control of both members	Dominant urinary "titratable acid"	Highest intracellular buffering capacity; ties into Bohr effect
Regulated by	Ventilation and renal H+/HCO3- handling	Renal phosphate excretion	Not directly; sets cellular baseline
Clinical readout	Arterial blood gas (pH, PCO2, HCO3-)	Urinary titratable acid	Base excess reflects total buffer base

Famous experiments and history

• Henderson's mass-action insight (1908). Lawrence J. Henderson, at Harvard, treated carbonic acid quantitatively and argued that the CO2/bicarbonate pair is uniquely fit as a physiological buffer precisely because its acid member is a gas the lungs can excrete — an early triumph of applying physical chemistry to physiology (later expanded in his 1913 book The Fitness of the Environment).
• Hasselbalch's logarithmic form (1916). Karl Albert Hasselbalch recast Henderson's relation into pH = pKa + log([base]/[acid]) using S.P.L. Sørensen's newly coined pH notation (1909), giving the compact clinical equation still taught today.
• The 1952 Copenhagen polio epidemic. Facing hundreds of patients in respiratory failure, anesthesiologist Bjørn Ibsen introduced positive-pressure hand ventilation and drove demand for rapid blood-gas measurement. Poul Astrup and, later, Ole Siggaard-Andersen developed the base excess concept and the Astrup blood-gas apparatus — the foundation of the modern arterial blood gas.
• The Copenhagen-versus-Boston debate (1960s). The Copenhagen school championed base excess and the CO2-bicarbonate ("Siggaard-Andersen") nomogram; the Boston school of William Schwartz and Arnold Relman argued for measured bicarbonate plus empirically derived compensation bands (the origin of rules like Winters' formula). The two frameworks still coexist in clinical practice.
• Pitts and renal ammoniagenesis. Robert Pitts and colleagues in the mid-20th century quantified how the kidney excretes fixed acid, showing that ammonium production from glutamine — not free H+ excretion — is the adaptable engine that lets renal acid excretion climb several-fold during chronic acidosis, the linchpin of long-term acid-base defense.