Acid-base balance

Body fluids are constantly exposed to fluctuating concentrations of proton-releasing acids and proton-accepting bases. The proton concentration of a solution is indicated by its pH value. Many of the body's biochemical processes require a stable pH. Fluctuations in pH can lead to the denaturation and loss of function of proteins and must therefore be compensated for by the interplay of buffer systems in the blood, along with regulation by the lungs and kidneys. Failure of these mechanisms can result in metabolic acidosis, characterized by decreased blood pH, often seen in poorly controlled diabetes mellitus due to excess acid accumulation. In contrast, excessive proton loss , e.g., during hyperventilation, leads to metabolic alkalosis, marked by elevated blood pH. See: "Acid-base disorders" for more clinical examples.

Many inorganic and organic substances react as either "acidic" or "basic." There are several definitions for this property, of which the Brønsted-Lowry definition is presented here. The Lewis definition is covered in the article on redox reactions.

Brønsted-Lowry acid-base definition

• Acid: proton donor
  • E.g., HCl (= hydrochloric acid): HCl + H₂O ⇄ Cl⁻ + H₃O⁺
• Base: proton acceptor
  • E.g., C₅H₅N (= pyridine): C₅H₅N + H₂O ⇄ C₅H₆N⁺ + OH⁻
• Conjugate acid-base pairs: When an acid donates a proton, it becomes a conjugate base, which can accept a proton. The conjugate base and the original acid form a conjugate acid-base pair.
  • E.g., HCl (= hydrochloric acid): HCl [acid] + H₂O [base] ⇄ Cl⁻ [conjugate base of HCl] + H₃O⁺ [conjugate acid of H₂O]
• Amphiprotic substance: compounds that can act as either proton donors or proton acceptors, depending on the chemical environment
  • E.g., water: H₂O + H₂O ⇄ H₃O⁺ + OH^-
    • E.g., amino acids, see also:

Not all acids and bases react with equal strength; they can be categorized as strong or weak. This distinction is crucial for various applications, including pH calculations and buffer composition.

Acid strength

The strength of an acid is determined by how completely it donates protons (also known as dissociation).

• Measuring acid strength: The strength of an acid can be quantified using the equilibrium constant (K_a), derived from the law of mass action for the deprotonation reaction
  • Calculation (e.g., for hydrochloric acid): K_a = ([H₃O⁺] × [Cl⁻])/[HCl]
    • K_a > 1: indicates a strong acid, where the majority of molecules are dissociated
    • K_a < 1: indicates a weak acid, where the majority of molecules remain undissociated
  • pK_a value: The pK_a is the negative base-10 logarithm of K_a: pK_a=−log(K_a).
    • A larger K_a corresponds to a smaller pK_a, signifying a stronger acid.
• Example: HCl (hydrochloric acid), with its negative pK_a value of −6, is a stronger acid than H₃PO₄ (phosphoric acid, pK_a = 2.1)
• Acid strength of carboxylic acids: The acid strength of a carboxylic acid increases when a strong electron-withdrawing substituent is attached to the carbonyl carbon.
  • This substitution enhances acidity by stabilizing the negative charge of the acid anion, a phenomenon referred to as the −I effect (negative inductive effect).
  • An example is trichloroacetic acid, which has powerful electron-withdrawing chlorine atoms.

Base strength

The strength of a base is defined by its ability to accept protons (known as protonation).

• Measuring base strength: The strength of a base can be quantified using the equilibrium constant (K_b), based on the law of mass action for the protonation reaction.
  • Calculation (e.g., for ammonia): K_b = ([OH⁻] × [NH₄⁺])/[NH₃]
    • K_b > 1: indicates a strong base, where the majority of molecules are protonated
    • K_b < 1: indicates a weak base, where the majority of molecules remain unprotonated
    • pK_b value: The pK_b is the negative base-10 logarithm of K_b: pK_b= −log(K_b)
      • A larger K_b corresponds to a smaller pK_b, indicating a stronger base.
• Example: OH⁻ (hydroxide ion), with its negative pK_b value of −1.7, is a stronger base than NH₃ (ammonia, pK_b = 4.75)

The larger the K_a/K_b value, the smaller the pK_a/pK_b value, and the stronger the acid/base.

Weak acids and bases do not completely dissociate in solution, resulting in limited conductivity and making them poor conductors of electricity. In contrast, strong acids and bases fully dissociate, leading to high conductivity and making them excellent electrical conductors.

• Definition: The pH value is the negative base-10 logarithm of the hydronium ion (H₃O⁺) concentration in a solution. As a dimensionless quantity, pH is expressed without units.
• Formula: pH = −log[H₃O⁺]
  • Alternatively, it is commonly represented as: pH = −log[H⁺], where H⁺ denotes the concentration of hydrogen ions.
  • Example: In pure water, the concentration of hydronium ions is 10⁻⁷ mol/L. The negative base-10 logarithm of this value yields a pH of 7.
• pH scale: 0 (strongly acidic) – 7 (neutral) – 14 (strongly basic)
• Interpretation: The pH value indicates whether a solution is acidic or basic.

When an acid is added to a solution, it becomes acidic (the acid transfers its protons to H₂O molecules, so there are more H₃O⁺ ions than OH⁻ ions)!

When a base is added to a solution, it becomes basic (the base accepts protons from H₂O molecules, so there are more OH⁻ ions than H₃O⁺ ions)!

pH value calculation

Different formulas are used to calculate the pH value, depending on the strength of an acid or base. In the following, concentration values are abbreviated as c [in mol/L].

• Strong acids
  • Examples: hydrochloric acid (HCl), sulfuric acid (H₂SO₄), nitric acid (HNO₃), perchloric acid (HClO₄), trichloroacetic acid (Cl₃CCOOH)
  • pH formula: pH = −log([acid] × valence)
  • Sample calculation: What is the pH of a 0.05 molar HCl solution?
    • pH = −log (0.05 × 1)
    • pH ≈ 1.3
• Weak acids
  • Examples: carbonic acid (H₂CO₃), citric acid (C₆H₈O₇), acetic acid (CH₃COOH), ammonium ion (NH₄⁺), phosphoric acid anions (H₂PO₄⁻, HPO₄²⁻), propionic acid (CH₃CH₂COOH)
  • pH formula: pH = ½ (pK_a – log[acid])
  • Sample calculation: What is the pH of a 0.1 molar acetic acid solution (pK_a = 4.75)?
    • pH = ½ (4.75 – log(0.1))
    • pH = ½ (4.75 + 1) = 2.88
• Strong bases
  • Example: hydroxide ion (OH⁻) as in potassium hydroxide (KOH) and sodium hydroxide (NaOH)
  • pOH formula: –log ([base] × valence)
  • pH formula: pH = 14 – pOH = 14 + log ([base] × valence)
  • Sample calculation: What is the pH of a 0.01 molar potassium hydroxide solution (KOH)?
    • pH = 14 + log (0.01 × 1)
    • pH = 14 + (−2) = 12
• Weak bases
  • Examples: ammonia (NH₃), amines
  • pH formula: pH = 14 − ½ (pK_b − log[base])
  • Sample calculation: What is the pH of a 0.1 molar ammonia solution (pK_b = 4.75)?
    • pH = 14 – ½ (4.75 – log(0.1))
    • pH = 14 – ½ (4.75 + 1) = 14 – 2.88 = 11.12

pH measures the concentration of hydronium ions (H₃O⁺) in a solution and is calculated using the formula pH = −log[H₃O⁺] or pH = −log[H⁺].

Remember the relationship between pH and pOH: pH + pOH = 14. This is useful for calculating pH from pOH values.

Autoprotolysis and the pH of water

Since water is an ampholyte, one H₂O molecule can transfer a proton to a second H₂O molecule. One water molecule acts as an acid and the other as a base. This process is also called the autoprotolysis of water.

• Definition
  • Protolysis: a chemical reaction in which a proton (H⁺) is transferred from one reactant to another (proton transfer reaction)
  • Autoprotolysis of water: one H₂O molecule transfers a proton (H⁺) to a second H₂O molecule
• Reaction equation: H₂O + H₂O ⇄ H₃O⁺ + OH^-
  • From the law of mass action for this equation, the so-called ion product of water can be calculated (K_W)
    • It represents the product of OH^– and H₃O⁺ ion concentrations in aqueous solutions
    • K_W= ([H₃O⁺] × [OH⁻]) = 10⁻⁷ mol/L × 10⁻⁷ mol/L = 10⁻¹⁴ mol²/L² (at 25°C)
  • In pure water, the number of OH^– and H₃O⁺ ions is equal, both having a concentration of 10⁻⁷ mol/L (1 L of water contains approx. 55 mol of water molecules).
  • The ion product of water can be used to relate the pK values of a conjugate acid-base pair: pK_W = pK_a + pK_b = 14.

The following table illustrates the relationship between the excess of either hydronium or hydroxide ions in a solution and the pH value:

Neutralization and the pH of salt solutions

• Neutralization reaction: a chemical reaction in which an acid and a base combine to form water and a salt
  • Equivalence point: the point in a neutralization reaction where the number of acid equivalents (e.g., H⁺) equals the number of base equivalents (e.g., OH⁻)
  • pH of the salt solution: The pH at the equivalence point is only neutral (pH = 7) for a strong acid + strong base reaction.
    • Example (strong acid + strong base): HCl (acid) + KOH (base) ⇄ H₂O (water) + KCl (salt)

When salts are dissolved in water, the pH of the resulting solution is not always neutral (pH = 7), but rather depends on the acid and base strengths of the ions.

Sample calculation (sodium acetate):

The reaction of NaOH (= strong base) with acetic acid (= weak acid) produces sodium acetate (= a weakly basic salt). Therefore, the pH is calculated using the formula for weak bases:

• Given: concentration of sodium acetate = 0.1 mol/L and pK_b (acetate) = 9.25 (since pK_a of acetic acid is 4.75)
• Substitute into the formula (pH = 14 – ½ (pK_b – log[base]))
• It follows that: pH = 14 – ½ (9.25 – log 0.1) = 14 – ½ (9.25 + 1) = 14 – 5.125 = 8.875 (weakly basic pH)

Common-ion effect

• Definition: a decrease in the dissociation of a weak acid or weak base when a salt containing one of its conjugate ions is added to the solution
• Mechanism: based on Le Chatelier's principle
  • For a weak acid (HA ⇄ H⁺ + A^-): Adding a salt with the common ion A^- (e.g., NaA) shifts the equilibrium to the left.
  • This reduces the dissociation of HA and decreases the [H⁺] (increases pH).
• Application: This effect is the fundamental principle behind buffer solutions.

Buffer substances stabilize the pH of solutions, crucial for maintaining physiological functions. The body employs endogenous buffers, such as the bicarbonate buffer system, to keep blood pH largely constant.

• Definition: Buffers are aqueous solutions of weak conjugate acid-base pairs, such as acetic acid/acetate or ammonium/ammonia .
• Mechanism
  • The combination of a weak acid and its conjugate base allows a buffer to absorb minor changes in pH.
  • The pH remains relatively stable when small amounts of acid or base are added.
  • Significant additions of acid or base can overwhelm the buffer capacity, leading to noticeable changes in pH.
• Functions
  • In the body: Buffers maintain consistent pH levels, which is vital for enzymatic and metabolic processes.
  • Applications: Due to their stabilizing properties, buffers are also utilized in food preservation, pharmaceuticals, and cosmetics.
• pH calculation of a buffer solution
  • Henderson-Hasselbalch equation: pH = pK_a + log([base]/[acid])
    • pK_a = negative base-10 logarithm of the equilibrium constant of the acid, [base] = concentration of the base, [acid] = concentration of the conjugate acid
  • Example: acetic acid (HAc)/acetate (Ac^-) buffer
    • Given
      • pK_a(HAc) = 4.75
      • [HAc] = 1 mmol/L
      • [Ac⁻] = 1.3 mmol/L
    • Buffer system: HAc + H₂O ⇄ Ac⁻ + H₃O⁺
    • Calculation
      • pH = pK_a + log([Ac⁻]/[HAc])
      • pH = 4.75 + log(1.3 mmol/L/1 mmol/L)
      • pH = 4.75 + log(1.3) ≈ 4.75 + 0.11 ≈ 4.86
• Buffer capacity: corresponds to the amount of an acid or base needed to change the pH of 1 L of buffer solution by ±1
• Optimal pH (buffer): the pH value at which the buffer solution reaches its maximum buffer capacity; this occurs when equal molar amounts of acid and base are present in the mixture, in which case pH = pK_a

In many biochemistry and pharmacology experiments, maintaining narrow pH limits is essential for allowing desired reactions to proceed without interruption from sudden pH changes. A specific example is the polymerase chain reaction (PCR), which requires an optimal pH of approximately 7.8–8.4 for Taq polymerase to function effectively. To prevent abrupt shifts in pH, a buffer is added to the PCR reaction solution. This buffer should be prepared in advance using a conjugate acid-base pair in the appropriate concentration ratio. The Henderson-Hasselbalch equation is often employed to accurately calculate these required concentrations.

Titration is a procedure for experimentally determining an unknown concentration of acid or base (analyte) in a solution: To do this, a known concentration and volume of a base or acid (titrant) is added, and the change in pH is monitored, often using an indicator (color).

Definition

• There are two different forms of titration:
  • Acidimetry: addition of a known amount of acid to an unknown amount of base to calculate the concentration or amount of the base present in the solution
  • Alkalimetry: addition of a known amount of base to an unknown amount of acid to calculate the concentration or amount of the acid present in the solution

Procedure

1. A suitable indicator is chosen based on its ability to change color within a specific pH range.
   1. Often a weak organic acid or base that changes color within a specific pH range
   2. Visually indicates the completion of the reaction or the equivalence point
   3. The indicator is added to the analyte solution before starting the titration, or sometimes during
   4. Examples: phenolphthalein or methyl orange
2. The titrant (acid or base) is added drop by drop to a vessel containing the acid (or base) to be analyzed.
   1. Examples: sodium hydroxide (NaOH) or hydrochloric acid (HCl)
3. The color change indicates that the equivalence point has been reached.

Titration curve

• Creation
  • y-axis: represents the pH
  • x-axis: represents the volume of titrant added
• Special points on the curve
  • Equivalence point: The point where the amounts of the analyte (acid or base) and titrant are equal, indicating the completion of the reaction.
    • Measurement: not directly measured but calculated based on stoichiometry
    • Function: allows calculation of the unknown acid or base concentration
    • Determination: identified by a significant pH jump (inflection point) and often a color change of the indicator
    • pH value: depends on the strength of the acid and base
      • Weak acid with strong base: pH > 7 (basic)
      • Strong acid with strong base: pH = 7 (neutral)
  • Endpoint: The point in a titration at which a noticeable change is observed, typically marked by a color change of the indicator.
    • Measurement: directly observed during the titration
    • Function: serves as a visual cue to stop adding titrant
  • Buffer regions: areas in the titration curve where pH changes gradually before the equivalence point, indicating buffer capacity
  • Half-equivalence point
    • Definition: occurs when half of the acid or base has been neutralized
      • Function: important for understanding buffer systems
      • Determination: located at half the volume of the equivalence point
      • pH relation: At this point, the pH corresponds to the pK_a (for weak acids) or pK_b (for weak bases) of the substance.
  • Neutral point
    • Definition: the point at which a neutral pH (approximately 7) is reached
      • Function: indicates the volume of titrant needed to achieve a neutral solution
      • Measurement: identified when pH is measured to be 7

The equivalence point is essential for calculating unknown concentrations using stoichiometry (C₁V₁ = C₂V₂), where C is the concentration and V is the volume of the solutions.

Recognize that the steep rise around the equivalence point indicates a rapid pH change with small additions of titrant. In contrast, flatter regions, where the pH remains stable, suggest the presence of a buffer solution that effectively neutralizes the added acid or base (titrant).

Comparison of curve shapes

• Strong acid + strong base
  • Curve shape: sharp rise around pH 7 at the equivalence point
  • Characteristics: quick transition from acidic to neutral, with minimal buffering capacity
• Weak acid + strong base
  • Curve shape: equivalence point occurs at a pH > 7, with a less steep rise
  • Characteristics: initial gradual rise, indicating weak acid buffering until the equivalence point
• Strong acid + weak base
  • Curve shape: similar to weak acid and strong base, but with the equivalence point at a pH < 7
  • Characteristics: The curve will rise slowly initially and then steepen as reaching the equivalence point.

Polyprotic acid titration

• Definition: applicable to acids that can donate multiple protons; e.g., phosphoric acid (H₃PO₄), carbonic acid (H₂CO₃)
• Titration curve: shows multiple equivalence points, one for each proton donated
• Half-equivalence points
  • Each half-equivalence point corresponds to pH = pK_a for each deprotonation step.
  • Example: at the first half-equivalence point for H₃PO₄, pH = pK_a1 and [H₃PO₄] = [H₂PO₄^-]
• Application to amino acids: amino acids are a key biological example of polyprotic acids
  • Because every amino acid has at least two ionizable groups (the alpha-carboxyl and alpha-amino), its titration curve will always show at least two pK_a values and two buffer regions.
  • Amino acids with ionizable side chains (like Asp, Glu, Lys, His) are triprotic and show three pK_a values.
  • For more information, see "Amino acid titration."

Redox titration

• Definition: a method for determining the concentration of an oxidizing or reducing agent through an oxidation-reduction reaction with a titrant
• Principle: relies on the balance between the moles of oxidizing and reducing agents being equal
• Equivalence point: reached when the moles of the oxidizing agent equal the moles of the reducing agent
• Monitoring methods
  • Redox indicators: change color at specific electrical potentials, providing visual cues for the equivalence point
  • Potentiometer: directly measures the electrical potential during titration, offering a more precise determination of the equivalence point
• Applications: commonly used in determining the concentration of substances like vitamin C, iron, and other analytes

Solubility principles

Just as the acid-base properties of ions determine the pH of a salt solution, these same properties, along with the presence of other ions, also determine how well a salt dissolves. This is described by solubility principles.

Solubility product constant (K_sp)

• Definition: the equilibrium constant for the dissolution of a sparingly soluble salt in solution
• Formula: for a salt A_mB_n, K_sp = [Aⁿ⁺]^m[B^m-]ⁿ
• Ion product (Q): has the same formula as K_sp but uses the actual ion concentrations, not equilibrium concentrations
  • Q < K_sp: solution is unsaturated; more salt can dissolve
  • Q = K_sp: solution is saturated (at equilibrium)
  • Q > K_sp: solution is supersaturated; precipitation will occur

Solubility and pH relationships

• Principle: The solubility of a salt can be pH-dependent if one of its ions is a weak acid or base.
• Example (basic anion): CaF₂, where F^- is the conjugate base of the weak acid HF
  • In acidic solution (low pH): H⁺ ions react with F^- ions to form HF.
  • This removes F^- from the dissolution equilibrium (Le Chatelier's principle).
  • Equilibrium shifts right (CaF₂(s) ⇄ Ca²⁺ + 2F^-), increasing solubility
• General rule: solubility of salts with basic anions (conjugates of weak acids) increases as pH decreases (solution becomes more acidic)

Common-ion effect and solubility

• Principle: The solubility of a sparingly soluble salt is decreased when a common ion (an ion already in the salt) is added to the solution.
• Mechanism: Based on Le Chatelier's principle, adding a product ion shifts the dissolution equilibrium to the left, favoring the solid salt.
• Application (laboratory separations): This effect can be used to precipitate ions from a solution selectively.
  • Example: To separate Ca²⁺ from Na⁺ in a solution, one could add a source of F^- (like NaF). The common ion F^- will cause CaF₂ (which has a low K_sp to precipitate, while NaF remains soluble.

Complex ion formation and solubility

• Definition: A complex ion (or coordination complex) consists of a central metal cation bonded to one or more ligands.
• Effect on solubility: Formation of a soluble complex ion can dramatically increase the solubility of an otherwise insoluble salt.
• Example: AgCl is insoluble (low K_sp). Adding ammonia (NH₃), a ligand, forms the soluble complex ion [Ag(NH₃)₂]⁺_.
  • This removes Ag⁺ ions from the AgCl dissolution equilibrium.
  • The equilibrium shifts right, causing more AgCl solid to dissolve.

The function and survival of an organism are linked to the condition of a constant pH value. The main reason for this is the pH-sensitive spatial structure of proteins. The term "acid-base balance" refers to all the regulatory mechanisms that are intended to prevent or compensate for deviations from the target pH value.

Normal pH values in the body

The secretions and compartments of the body are sorted in the following list according to increasing pH value.

A decreased arterial blood pH value (pH < 7.35) is called acidosis, and an increased arterial blood pH value (pH > 7.45) is called alkalosis.

Influences on pH in the body

Various physiological factors contribute to fluctuations in the body’s pH. For example, the citric acid cycle produces volatile acids like carbon dioxide (CO₂). The body maintains a balance of acidic and basic metabolic products through various regulatory mechanisms, ensuring stable pH levels.

To maintain stable pH levels in the body, several continuously active regulatory systems are in place, which can be categorized into buffer solutions in body fluids and organ-specific mechanisms. The human body contains multiple buffer systems that effectively compensate for acute fluctuations in blood pH, keeping it approximately constant at around 7.4. The vast majority of sudden increases in proton concentration are immediately buffered by these systems.

Open buffer systems

Open buffer systems are characterized by the fact that one reactant can be removed from the system (e.g., via the lungs or kidneys), which increases the buffer capacity. The two most important open buffer systems in humans are the bicarbonate and ammonium buffer systems.

Bicarbonate buffer system

The bicarbonate buffer is the most important buffer system in the human body. It acts as an open buffer system in the excretion of acid equivalents via the lungs by exhaling CO₂. At 20–28 mmol/L, the bicarbonate buffer system accounts for about half of the total buffer capacity of the blood.

• Key reaction: H₂O + CO₂ ⇄ H₂CO₃ ⇄ HCO₃⁻ + H⁺
• Example
  • Given
    • pK_a(combined) = 6.1
    • [HCO₃⁻] = 24 mmol/L
    • [CO₂] = 1.2 mmol/L
  • Key reaction: CO₂ + H₂O ⇄ H₂CO₃ ⇄ HCO₃⁻ + H⁺
  • Calculation
    • pH = pK_a + log([base]/[acid])
    • pH = pK_a + log([HCO₃⁻]/[CO₂])
    • pH = 6.1 + log(24 mmol/L/1.2 mmol/L)
    • pH = 6.1 + log(20) = 6.1 + 1.3 = 7.4
• Properties with concentration changes
  • If the concentration of the conjugate base HCO₃⁻ increases, the pH increases.
  • If the concentration of CO₂, which acts as a weak acid, increases, the pH decreases.
• Function: buffering of blood pH by CO₂ dissolved in the blood
  • Excess of acids: HCO₃⁻ increasingly accepts protons and is exhaled as CO₂.
  • Excess of bases: Dissolved CO₂ converts to its deprotonated form (HCO₃⁻) via H₂CO₃, and less CO₂ is exhaled.
  • Increased CO₂ partial pressure in the blood: The total concentration of buffer bases remains constant.

Effectiveness of the buffer: increases in alkalosis (pH ↑) and decreases in acidosis (pH ↓)

When there is excess acid in the blood, bicarbonate neutralizes it by accepting protons and converting to carbonic acid, which is then eliminated by the lungs as CO₂. Conversely, if the blood becomes too alkaline, carbon dioxide levels can be regulated to adjust acidity.

Ammonium buffer system

The ammonium buffer system is an important regulatory system for renal acid-base excretion, with which acidic substances can be excreted permanently. In addition, the system is involved in renal gluconeogenesis, the de novo synthesis of bicarbonate, and the regulation of intracellular pH.

• Key reaction: NH₃ + H⁺ ⇄ NH₄⁺
• Function
  • Enables H⁺ excretion via the urine in the form of NH₄⁺
  • HCO₃⁻-sparing method of NH₃ excretion

Closed buffer systems

Closed buffer systems have a lower buffer capacity than open buffer systems. They are characterized by the fact that the sum of the concentrations of the acid and its conjugate base remains constant.

Protein buffer system

Proteins in the blood can function as buffers via ionizable side groups. Hemoglobin in erythrocytes and albumin play the largest role due to their high concentration.

• Reaction: H⁺ uptake via reactive groups of amino acids (see also: Acid-base properties of amino acids)
  • Basic imidazole ring of histidine: The imidazole ring of histidine has a pK_a value of 6.0, close to the physiological pH range, which is why histidine is particularly important for the buffer function of plasma proteins.
  • Amino and carboxyl groups of all amino acids
• Function: regulation of blood pH (accounts for the other half of the total buffer capacity of the blood)
• Main representatives
  • Albumin
  • Hemoglobin: deoxygenated hemoglobin (= Hb) has a lower acidity than oxygenated Hb → deoxy-Hb is more likely to accept H⁺ at the same pH than oxy-Hb
    • This also results in the relationship between pH and the Hb-oxygen-hemoglobin dissociation curve: pH↓ → favors deoxy-Hb → right shift; pH↑ → favors oxy-Hb → left shift

Phosphate buffer system

The phosphate buffer system is important for regulating both the intracellular pH of all body cells and the pH of urine.

• Key reaction: PO₄³⁻ + 3 H⁺ ⇄ HPO₄²⁻ + 2 H⁺ ⇄ H₂PO₄⁻ + H⁺ ⇄ H₃PO₄
  • Henderson-Hasselbalch equation (for the physiologically relevant pair)
    • pH = pK_a + log ([base]/[acid])
    • pH = 6.8 + log ([HPO₄²⁻]/[H₂PO₄⁻])
• Function
  • Regulation of intracellular pH
  • Regulation of urine pH in the form of HPO₄^2– and H₂PO₄^– (cf. proton secretion in the chapter on tubular transport processes)

The body's buffer systems can only compensate for pH shifts for a short time and to a limited extent. Without intact lung and kidney function, the acid-base balance collapses, and life-threatening conditions can occur.

The role of the lungs in pH regulation

The role of the lungs in the acid-base balance is to exhale the "volatile acid" CO₂, which is constantly produced in the body as a byproduct of energy metabolism. Only if the arterial CO₂ partial pressure is kept constant does the blood pH also remain unchanged.

• Mechanism: elimination of the volatile acid CO₂ via exhaled air
  • CO₂-producing reactions in the body
    • Physiological: metabolism → oxidation of the carbon atoms of carbohydrates, fats, and proteins/amino acids → CO₂
    • Pathological: during compensation of metabolic acidosis by the bicarbonate buffer system → HCO₃⁻ + H⁺ ⇄ H₂O + CO₂
  • Balance: elimination of an average of 16,000 mmol of CO₂ per day
• Regulation
  • CO₂ respiratory drive: pCO₂ increase in the blood → stimulation of respiration
  • pH respiratory drive: pH drop in the blood → stimulation of respiration
  • Mechanoreceptors: degree of physical activity is measured in skeletal muscles/joints
    • High activity, which is associated with an increased production of CO₂ → stimulation of respiration

Hyperventilation leads to a pH increase (alkalosis) by exhaling the weak acid CO₂, while in hypoventilation, more CO₂ remains in the body, leading to a pH drop (acidosis)!

The role of the kidneys in pH regulation

The kidneys regulate pH via two mechanisms: on the one hand, they excrete excess H⁺ ions largely in the form of NH₄⁺ and H₂PO₄⁻. On the other hand, they maintain the HCO₃⁻ concentration in the blood by reabsorbing it from the urine and performing de novo synthesis of HCO₃⁻. For more information, see also proton secretion and bicarbonate reabsorption.

• Excretion of acid equivalents
  • Proton secretion: via the Na⁺/H⁺ antiporter in the proximal tubule and an H⁺-ATPase as well as H⁺/K⁺-ATPase of the intercalated cells in the late distal tubule and collecting duct
    • Buffering in the urine by binding to ammonia and phosphate
  • Regulation
    • In acidosis: excretion of NH₄⁺ and H₂PO₄⁻ (titratable acid) increases due to increased ammonia production in the kidney (glutaminase activity ↑) and decreased phosphate reabsorption in the proximal tubule
      • Enzyme induction: chronic acidosis → renal glutaminase and glutamate dehydrogenase are increasingly expressed
• Production and reabsorption of base equivalents
  • De novo synthesis of HCO₃^– by the kidney: Through the conversion of one molecule of glutamine to one molecule of α-ketoglutarate (cf. NH₄⁺ excretion, process), the body gains two molecules of HCO₃^–.
  • HCO₃^– reabsorption: Indirect reabsorption in the proximal tubule via conversion to CO₂ with proton consumption
    • Regulation
      • In acidosis: HCO₃⁻ is almost 100% reabsorbed.
      • In alkalosis: HCO₃⁻ reabsorption ↓ and stimulation of the Cl⁻/HCO₃⁻ exchanger

The role of the liver in pH regulation

The role of the liver in pH regulation is directly linked to its NH₃ detoxification function. The liver has two ways of detoxifying NH₃: via the urea cycle and via glutamine synthesis. Under normal circumstances, 95% of the resulting NH₃ is metabolized in the urea cycle and 5% via glutamine synthesis. In case of blood pH deviation, for example, glutamine synthesis can be increased and the urea cycle simultaneously inhibited to save HCO₃⁻.

• Mechanisms of ammonia detoxification
  • Urea cycle
    • Process: Urea is synthesized in five reactions from one molecule of NH₃, one molecule of HCO₃⁻, and the amino group of an aspartate.
    • Importance for the organism: main mechanism of NH₃ detoxification (= toxic breakdown product of amino acid metabolism)
    • Influence on acid-base balance: requires the base HCO₃⁻
  • Glutamine synthesis
    • Process: The synthesis of the amino acid glutamine from the precursor glutamate with the incorporation of NH₃ is a way of detoxifying NH₃ without consuming HCO₃^- .
      1. In the liver: Glutamate + NH₃ → glutamine (enzyme: glutamine synthetase)
      2. Glutamine reaches the kidneys via the blood
      3. In the kidneys: cleavage of NH₃ from glutamine → NH₃ + H⁺ are excreted in the form of the acid NH₄⁺ without consuming HCO₃⁻ (for the exact process in the kidneys, cf. proton secretion)
    • Influence on acid-base balance: does not require HCO₃⁻ → mechanism of base conservation in acidosis
• Regulation of ammonia detoxification in case of pH deviation
  • Acidosis: urea synthesis ↓ and glutamine synthesis ↑
    • Urea synthesis is reduced to conserve HCO₃⁻ (e.g., for buffering blood pH)
    • In the liver, more glutamine is formed and delivered to the kidney via the blood → NH₃ is increasingly excreted renally in the form of the acid NH₄⁺
  • Alkalosis: urea synthesis ↑ and glutamine synthesis ↓
    • Urea synthesis is increased because HCO₃⁻ is sufficiently available → NH₃ is increasingly excreted in the form of urea
    • Glutamine synthetase is inhibited, so that less glutamine reaches the kidney → renal NH₄⁺ excretion decreases

pH deviations alter the spatial structure of proteins, which in turn regulate many body functions. In pH disorders, a distinction is made between a decrease in arterial pH < 7.35 (= acidosis) and an increase in pH > 7.45 (= alkalosis).

Consequences of pH deviation

Most effects of a pH shift are caused by influencing enzymes or ion-selective transmembrane channels.

A blood pH of < 6.8 or > 7.8 is generally incompatible with life.

Respiratory and metabolic disorders

A pH disorder is usually caused by only one of the two regulatory systems: the lungs (respiratory disorder) or metabolism (metabolic disorder). The intact system then tries to compensate for the pH deviation. For more information, see "Acid-base disorders."