Physiology of Acid-Base Balance

The pH of the blood is tightly regulated within the range 7.35–7.45 to ensure proper physiologic functions. Large amounts of acid are generated each day through normal processes (aerobic/anaerobic respiration and dietary intake), and these are efficiently managed and eliminated by buffers in the blood, the respiratory system, and the renal system. When these regulatory systems are disturbed, acid–base balance disorders occur, including respiratory acidosis, respiratory alkalosis, metabolic acidosis, and metabolic alkalosis.

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pH Overview

pH is the quantitative measurement of the acidity or basicity of a solution.

  • pH = –log10[H+]:
    • [H+] = concentration of hydrogen ions (i.e., protons) in solution
    • Logarithmic scale from 1 to 14 
      • 1 = maximally acidic, 14 = maximally basic
      • 7 = neutral point: equal concentrations of H+ and OH 
  • Normal arterial blood pH is approximately 7.40. 
  • The normal range is tightly regulated to stay between 7.35 and 7.45.
    • Acidemia: more hydrogen ions (H+) in the blood = pH < 7.35
    • Alkalemia: more hydroxide ions (OH) in the blood = pH > 7.45 
  • “-emia” vs. “-osis”
    • “-emia” means “in the blood.”
    • “-osis” refers to a process: Acidosis and alkalosis refer to the processes that cause acidemia and alkalemia.

Relation between blood pH and concentration of hydrogen ions

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Table: Examples of pH values of various fluids
Gastric secretions (under conditions of maximal acidity)0.7
Chromaffin granule5.5
Neutral H2 at 37°C6.81
Cytosol of a typical cell6.0–7.4
Arterial blood plasma7.35–7.45
Mitochondrial inner matrix7.5
Pancreatic secretions8.1

Acids, Bases, and Buffers

The relative concentrations of acids and bases in the blood determine its pH. Buffers provide a short-term solution for disturbances in this balance before the lungs and kidneys can act definitively to restore balance. 


Acids are compounds that can donate protons (H+)  or accept electrons.

  • H+ are released when acids dissociate in solution → ↓ pH
  • Classified by strength and volatility:
    • Strong acids: 
      • Fully ionize in water 
      • More H+ released into water → greater effect on pH
      • Example: hydrochloric acid (HCl)
    • Weak acids:
      • Partially ionize in water 
      • Less H+ released into water → relatively less effect on pH
      • Example: carbonic acid (H2CO3)
    • Volatile acids:
      • Can change phase into a gas → removable through the lungs
      • Primary example: CO2
      • The body produces approximately 15,000 mM of volatile acid per day through aerobic metabolism.
    • Nonvolatile (fixed) acids:
      • Cannot change phase into a gas → not removable through the lungs
      • Removed by the kidneys
      • The body produces approximately 70 mM of fixed acids per day through anaerobic metabolism and the GI tract.
      • Examples: lactic acid, uric acid, sulfuric acid, phosphoric acid


Bases are compounds that can accept protons (H+) or donate electrons.

  • Hydroxide ions (OH) are released when bases dissociate into solution:
    • OH combine with free H+ to form H2O.
    • Net result is less [H+] → ↑ pH (becomes more basic)
  • Classified by strength:
    • Strong bases:
      • Fully ionize in water
      • More OH released into water → greater effect on pH
      • Example: sodium hydroxide (NaOH)
    • Weak bases:
      • Partially ionize in water
      • Less OH released into water → relatively less effect on pH 
      • Examples: bicarbonate (HCO3), ammonia (NH3)


Buffers are substances that consume or releases hydrogen ions (H+) to stabilize the pH.

  • Buffer power refers to the quantity of H+ that can be added or removed.
  • Categorized as bicarbonate and nonbicarbonate buffers:
    • Bicarbonate (HCO3):
      • Most physiologically important buffer
      • HCO3 + H+ ⇆ H2CO3 ⇆ CO2 + H2O
    • Nonbicarbonate buffers:
      • Less physiologically important
      • Examples: proteins (e.g., albumin, hemoglobin), phosphates
  • pK: pH of a buffer when it is 50% ionized
    • Example: bicarbonate
      • HCO3 + H+ ⇆ H2CO3 
      • 50% HCO3 and 50% H2CO3 occurs at pH 6.1.
      • pK of bicarbonate = 6.1
    • Defines the optimal pH range for buffering capability of a particular buffer
    • Buffers work best when pK is near the pH of the fluid they are buffering.
  • Henderson–Hasselbalch equation:
    • Formula used to determine the pH of blood
    • The pH of the blood depends primarily on the ratio between the amounts of HCO3 (base) and CO2 (acid) in the blood.
    • Equation: pH = 6.1 + log ([HCO3]/[0.03 × PCO2])
      • 6.1 = pK of HCO3 (dominant buffer in the blood)
      • [HCO3] = concentration of bicarbonate in the blood measured in mEq/L 
      • PCO2 = partial pressure of CO2 in the blood measured in mm Hg
      • 0.03 = solubility factor for CO2
    • Used to generate titration curves
Table: Henderson–Hasselbalch examples
Normal ABGAcidic ABGAlkalotic ABG
HCO324 mEq/L26 mEq/L22 mEq/L
PaCO240 mm Hg60 mm Hg20 mm Hg
ABG: arterial blood gas
HCO3: bicarbonate
PaCO2: partial pressure of CO2 in arterial blood

Titration curve for bicarbonate in the blood

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Acid Handling in the Lungs

The body produces approximately 15,000 mM of volatile and 70 mM of nonvolatile acids daily. The lungs and kidneys work in concert to eliminate this daily acid load, which prevents the buffering capability of the blood from being overwhelmed and allows it to maintain a normal pH.

  • The primary acid load produced by the body is in the form of CO2 (a volatile acid) produced via aerobic metabolism.
  • CO2 is eliminated through the respiratory tract.
  • ↑ CO2 → ↑ respiration

Factors involved in daily acid–base balance

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Acid Handling in the Kidneys

The kidneys are primarily responsible for elimination of the fixed (nonvolatile) acids, approximately 70 mM daily. They prevent excretion of bicarbonate, and also pair acid excretion with new bicarbonate generation so that the bicarbonate buffering system is always available at full capacity.

Bicarbonate reabsorption

Bicarbonate is freely filtered at the glomerulus, and 100% of it is then reabsorbed (80% in the proximal tubule) through the following process:

  1. Sodium–hydrogen ion exchanger 3 (NHE3) absorbs Na+ and secretes H+.
  2. Secreted H+ combine with the filtered HCO3 to form H2CO3 in the tubular lumen.
  3. H2CO3 is converted into H2O and CO2 by apical carbonic anhydrase IV. 
  4. CO2 diffuses freely across the apical membrane back into the cell.
  5. Intracellular carbonic anhydrase II converts CO2 + H2O back into H2CO3.
  6. H2CO3 then can dissociate into H+ and HCO3:
    • H+ are recycled through the process through NHE3.
    • HCO3 is absorbed through the basolateral membrane via:
      • Na+-HCO3 cotransporter 
      • HCO3-Cl exchanger 
  7. Net effects of the entire process:
    • Excretion of H+
    • Absorption of HCO3– 
  8. Note: Bicarbonate is not freely permeable across the apical membrane because it is a charged molecule.
  9. Locations of HCO3 reabsorption:
    • Proximal tubule: 80%
    • Thick ascending limb: 10%
    • Distal convoluted tubule: 6%
    • Collecting duct: 4%

Bicarbonate reabsorption in the proximal tubule

CA-IV: carbonic anhydrase IV
CA-II: carbonic anhydrase II

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Ammonia (NH3)

NH3 is able to help excrete fixed acids. 

  • Accounts for 60% of fixed acid excretion
  • Occurs primarily in the proximal tubule
  • NH3 can bind H+, which is then eliminated in the urine:
    • Within the tubule, NH3 can bind to H+ → becomes NH4+, which remains in the urine and is excreted 
    • NH4+ can be generated within the proximal tubule cell and excreted into the tubular fluid.
  • NH4+ (ammonium) and HCO3 are generated from the metabolism of glutamine in the mitochondria: 
    • Glutamine dehydrogenase: glutamine → glutamate + NH4+, then
    • Glutamate dehydrogenase: glutamate → 𝛼-ketoglutarate + NH4+, then
    • 𝛼-Ketoglutarate enters the Krebs cycle → 2 molecules of HCO3 generated
  • NH4+ is secreted into tubular fluid via 2 mechanisms:
    • NHE3 directly exchanges Na+ for NH4+.
    • NH3 is membrane-permeable:
      • Crosses freely from inside cell into tubular lumen
      • Combines with free H+ to form NH4+
  • HCO3 produced in the Krebs cycle is reabsorbed across the basolateral membrane into the bloodstream.
  • This process is highly adaptable: upregulates in chronic acidosis.

NH3 and NH4+ transport to the lumen for excretion

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Titratable acids

  • Accounts for 40% of fixed acid excretion
  • Occurs in the proximal tubule, distal convoluted tubule, and collecting duct
  • Titratable acids buffer a H+ and then are eliminated in the urine:
    • Intracellular H2CO3 dissociates into H+ and HCO3.
    • H+ ions are secreted into the tubular lumen by:
      • V-type H+-ATPase 
      • NHE3 
    • Titratable acid binds H+.
    • Bound compound is excreted in the urine.
    • HCO3 (from first step) is left over inside the cell and is considered “regenerated.”
  • Examples of titratable acids:
    • pK is closest to the pH of normal urine → most important under normal conditions:
      • Phosphate (pK 6.8)
    • pK further away from normal urine pH → less useful:
      • Urate (pK 5.8)
      • Creatinine (pK 5.0)
    • pK less than the minimum possible urine pH (approximately 4.4) → not useful:
      • Lactate (pK 3.9)
      • Pyruvate (pK 2.5)
  • This system does not upregulate in chronic acidosis (unlike the ammonia system).

Clinical Relevance

When a disease process overwhelms the normal capability to regulate pH, the primary acid–base disorders listed below occur. Compensatory mechanisms also occur, which help to offset the change in pH.

  • Metabolic acidosis: processes that result in the accumulation of H+ and generally cause acidemia (i.e., the lowering of blood pH): This process can be the result of either excessive loss of buffers (e.g., diarrhea) or increased production of acids (e.g., ketoacidosis and lactic acidosis).
  • Metabolic alkalosis: processes that result in the accumulation of HCO3 and generally cause alkalemia (i.e., the rise of blood pH): This process can be the result of excessive supply of buffers or increased excretion of hydrogen ions. Common causes are vomiting and calcium–alkali syndrome.
  • Respiratory acidosis: processes that result in the accumulation of arterial CO2 and generally cause acidemia (i.e., the lowering of blood pH): This process can be the result of deficient clearance of CO2 by the lungs. Common causes are chronic obstructive pulmonary disease and asthma.
  • Respiratory alkalosis: processes that result in the decrease of arterial CO2 and generally cause alkalemia (i.e., the rise of blood pH): This process can be the result of excessive clearance of CO2 from the blood by increased ventilation. Common causes include pregnancy, anxiety, and aspirin overdose.


  1. Emmett, M., Palmer B.F. (2020). Simple and mixed acid–base disorders. UpToDate. Retrieved April 1, 2021, from 
  2. Theodore, A. C. (2020). Arterial blood gases. UpToDate. Retrieved April 1, 2021, from

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