Acid–base Balance

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Introduction

  • Acid–base balance is the process of maintaining the body’s pH within a normal range.
  • The normal arterial blood pH is 7.35–7.45, which is slightly alkaline.
  • Proper acid–base balance is essential for normal cellular and metabolic functions.
  • The body continuously produces acids during metabolism, which must be neutralized or eliminated.
  • Buffer systems, lungs, and kidneys work together to regulate the body’s pH.
  • Disturbances in acid–base balance can lead to acidosis or alkalosis, affecting organ function.
  • Arterial blood gas (ABG) analysis is commonly used to assess acid–base status in clinical practice.

Definition of Acid–Base Balance

  • Acid–base balance is the physiological process of maintaining the hydrogen ion (H⁺) concentration and body fluid pH within a narrow normal range (arterial blood pH: 7.35–7.45).
  • It is regulated by chemical buffer systems, the respiratory system, and the kidneys to ensure normal cellular metabolism and proper functioning of the body.

Importance of Acid–Base Homeostasis

Maintaining acid–base homeostasis is essential for normal physiological functions.

  • Maintains the normal pH (7.35–7.45) of body fluids.
  • Ensures optimal enzyme activity and metabolic reactions.
  • Maintains the normal structure and function of proteins.
  • Regulates electrolyte balance, especially potassium, sodium, and calcium.
  • Supports normal nerve impulse transmission and muscle contraction.
  • Facilitates efficient oxygen transport and release by hemoglobin.
  • Maintains normal cellular metabolism and ATP production.
  • Essential for the proper functioning of the heart, lungs, kidneys, and brain.
  • Prevents acidosis and alkalosis, which can impair organ function.
  • Helps maintain overall physiological homeostasis and good health.

Normal pH of Body Fluids

  • The pH of body fluids varies depending on their composition and physiological function.
  • Maintaining the normal pH of these fluids is essential for proper enzyme activity, metabolism, and cellular function.

Normal pH Values of Body Fluids

Body Fluid Normal pH
Arterial blood 7.35–7.45
Venous blood 7.31–7.41
Intracellular fluid  7.0–7.2
Gastric juice 1.5–3.5
Saliva 6.2–7.6
Urine 4.5–8.0 (Average: 6.0)
Pancreatic juice 7.8–8.3
Bile 7.6–8.6
Cerebrospinal fluid (CSF)  7.31

 

Acids and Bases: Definitions

  • Acids and bases are substances that influence the concentration of hydrogen ions (H⁺) in body fluids, thereby regulating the body’s pH.

Acid

  • An acid is a substance that releases hydrogen ions (H⁺) when dissolved in water.

Examples

  • Hydrochloric acid (HCl)
  • Carbonic acid (H₂CO₃)
  • Lactic acid
  • Sulfuric acid (H₂SO₄)

Base

  • A base is a substance that accepts hydrogen ions (H⁺) or releases hydroxyl ions (OH⁻) in solution.

Examples

  • Bicarbonate (HCO₃⁻)
  • Ammonia (NH₃)
  • Sodium hydroxide (NaOH)

Sources of Acids in the Body

  • Acids are continuously produced in the body as a result of normal metabolism.
  • These acids must be neutralized by buffer systems or excreted by the lungs and kidneys to maintain acid–base balance.

1. Volatile Acid (Carbonic Acid)

  • Produced from the metabolism of carbohydrates and fats.
  • Carbon dioxide (CO₂) combines with water to form carbonic acid (H₂CO₃).
  • Carbonic acid dissociates into H⁺ and HCO₃⁻.
  • Carbon dioxide is eliminated through the lungs.

Reaction:

CO₂ + H₂O ⇌ H₂CO₃ ⇌ H⁺ + HCO₃⁻

2. Non-Volatile (Fixed) Acids

These acids cannot be excreted through the lungs and are eliminated mainly by the kidneys.

a. Sulfuric Acid (H₂SO₄)

  • Produced during the metabolism of sulfur-containing amino acids (methionine and cysteine).

b. Phosphoric Acid (H₃PO₄)

  • Formed during the metabolism of phospholipids, phosphoproteins, and nucleic acids.

c. Organic Acids

  • Produced under normal or pathological conditions.

Examples

  • Lactic acid (during anaerobic glycolysis)
  • Ketoacids (during prolonged fasting, uncontrolled diabetes mellitus, or starvation)
  • Uric acid (from purine metabolism)

Summary Table

Type of Acid Source Route of Excretion
Carbonic acid (Volatile acid) Metabolism of carbohydrates and fats Lungs
Sulfuric acid Sulfur-containing amino acids Kidneys
Phosphoric acid Phospholipid and nucleic acid metabolism Kidneys
Lactic acid Anaerobic glycolysis Metabolized by liver and kidneys
Ketoacids Fat metabolism during starvation or diabetes Kidneys and metabolism after correction

Clinical Significance

  • Excess lactic acid may lead to lactic acidosis.
  • Increased ketoacid production causes ketoacidosis, commonly seen in uncontrolled diabetes mellitus.
  • Failure of the lungs or kidneys to eliminate acids results in acid–base disorders such as respiratory or metabolic acidosis.

Buffer Systems

  • A buffer system is a mixture of a weak acid and its corresponding salt (conjugate base) or a weak base and its corresponding acid that resists changes in pH when small amounts of acid or base are added.
  • Buffer systems provide the first line of defense against changes in body pH and help maintain the normal blood pH of 7.35–7.45 until the lungs and kidneys restore acid–base balance.

Major Buffer Systems in the Body

The four major physiological buffer systems are:

  1. Bicarbonate Buffer System
  2. Phosphate Buffer System
  3. Protein Buffer System
  4. Hemoglobin Buffer System

1. Bicarbonate Buffer System

  • The bicarbonate buffer system is the most important extracellular buffer system in the human body.
  • It consists of a weak acid (carbonic acid, H₂CO₃) and its conjugate base (bicarbonate ion, HCO₃⁻).
  • This buffer system plays a vital role in maintaining the normal blood pH (7.35–7.45) by neutralizing excess acids and bases.

Components of the Bicarbonate Buffer System

The bicarbonate buffer system has two main components:

  • Weak acid: Carbonic acid (H₂CO₃)
  • Conjugate base: Bicarbonate ion (HCO₃⁻)

The ratio of bicarbonate to carbonic acid in normal blood is approximately 20:1, which is essential for maintaining the normal blood pH.

Chemical Reaction

The bicarbonate buffer system is based on the following reversible reaction:

CO₂ + H₂O ⇌ H₂CO₃ ⇌ H⁺ + HCO₃⁻

Where:

  • CO₂ = Carbon dioxide
  • H₂O = Water
  • H₂CO₃ = Carbonic acid
  • H⁺ = Hydrogen ion
  • HCO₃⁻ = Bicarbonate ion

The reaction is catalyzed by the enzyme carbonic anhydrase, which is abundant in red blood cells and renal tubular cells.

Mechanism of Action

A. When Excess Acid is Added

When excess hydrogen ions (H⁺) are produced in the body:

  • Bicarbonate ions (HCO₃⁻) combine with H⁺.
  • Carbonic acid (H₂CO₃) is formed.
  • Carbonic acid rapidly breaks down into carbon dioxide (CO₂) and water.
  • Carbon dioxide is exhaled through the lungs.

Reaction:

H⁺ + HCO₃⁻ → H₂CO₃ → CO₂ + H₂O

Result: Excess acid is neutralized, preventing a fall in blood pH.

B. When Excess Base is Added

When excess hydroxyl ions (OH⁻) are present:

  • Carbonic acid (H₂CO₃) releases hydrogen ions (H⁺).
  • H⁺ combines with OH⁻ to form water.
  • The remaining bicarbonate helps maintain the buffer equilibrium.

Reaction:

OH⁻ + H₂CO₃ → HCO₃⁻ + H₂O

Result: Excess base is neutralized, preventing a rise in blood pH.

Regulation of the Bicarbonate Buffer System

The effectiveness of this buffer system depends on the coordinated action of the lungs and kidneys.

Role of the Lungs

  • Remove excess carbon dioxide (CO₂) through respiration.
  • Increased breathing removes more CO₂, reducing carbonic acid and increasing blood pH.
  • Decreased breathing retains CO₂, increasing carbonic acid and lowering blood pH.

Role of the Kidneys

  • Reabsorb filtered bicarbonate (HCO₃⁻).
  • Generate new bicarbonate ions when required.
  • Excrete excess hydrogen ions (H⁺) in urine.
  • Provide long-term regulation of acid–base balance.

Functions of the Bicarbonate Buffer System

  • Maintains the normal blood pH (7.35–7.45).
  • Acts as the primary extracellular buffer system.
  • Neutralizes excess acids and bases produced during metabolism.
  • Works closely with the lungs and kidneys to regulate acid–base balance.
  • Helps maintain normal enzyme activity and cellular metabolism.
  • Protects tissues from sudden changes in pH.

Advantages

  • Rapid and highly efficient buffer system.
  • Present in large amounts in extracellular fluid.
  • Easily regulated by the respiratory and renal systems.
  • Essential for maintaining acid–base homeostasis.

Clinical Significance

  • The bicarbonate buffer system is the most important buffer assessed during Arterial Blood Gas (ABG) analysis.
  • A decrease in bicarbonate (HCO₃⁻) is commonly seen in metabolic acidosis, such as diabetic ketoacidosis, lactic acidosis, and renal failure.
  • An increase in bicarbonate occurs in metabolic alkalosis, often due to prolonged vomiting or excessive diuretic therapy.
  • Respiratory disorders alter carbon dioxide levels, indirectly affecting the bicarbonate buffer system.
  • Disturbances in this system can lead to life-threatening acid–base disorders if not corrected promptly.

2. Phosphate Buffer System

  • The phosphate buffer system is an important physiological buffer that helps maintain the acid–base balance of the body.
  • It is most effective in the intracellular fluid (ICF) and renal tubular fluid, where it regulates pH by buffering excess hydrogen ions (H⁺) or hydroxyl ions (OH⁻).
  • Although it plays a smaller role in blood than the bicarbonate buffer system, it is essential for intracellular buffering and renal excretion of acids.

Components of the Phosphate Buffer System

The phosphate buffer system consists of:

  • Weak acid: Dihydrogen phosphate (H₂PO₄⁻)
  • Conjugate base: Hydrogen phosphate (HPO₄²⁻)

These two forms exist in equilibrium and act as an effective buffer around the physiological pH.

Chemical Reaction

The phosphate buffer system is represented by the following reversible reaction:

H₂PO₄⁻ ⇌ H⁺ + HPO₄²⁻

This reversible reaction allows the system to either accept or release hydrogen ions depending on the body’s pH.

Mechanism of Action

A. When Excess Acid is Added

When excess hydrogen ions (H⁺) are produced:

  • Hydrogen phosphate (HPO₄²⁻) combines with H⁺.
  • Dihydrogen phosphate (H₂PO₄⁻) is formed.
  • This removes free H⁺ from the solution and minimizes the fall in pH.

Reaction:

HPO₄²⁻ + H⁺ → H₂PO₄⁻

Result: Excess acid is neutralized, helping maintain normal pH.

B. When Excess Base is Added

When excess hydroxyl ions (OH⁻) are present:

  • Dihydrogen phosphate (H₂PO₄⁻) releases H⁺.
  • The released H⁺ combines with OH⁻ to form water.
  • Hydrogen phosphate (HPO₄²⁻) is regenerated.

Reaction:

H₂PO₄⁻ + OH⁻ → HPO₄²⁻ + H₂O

Result: Excess base is neutralized, preventing a rise in pH.

Role in the Kidneys

The phosphate buffer system plays a major role in the renal regulation of acid–base balance.

  • Buffers hydrogen ions (H⁺) secreted into the renal tubules.
  • Converts H⁺ into dihydrogen phosphate (H₂PO₄⁻), which is excreted in urine.
  • Facilitates the excretion of excess acids.
  • Helps conserve and regenerate bicarbonate (HCO₃⁻).
  • Contributes to maintaining normal blood pH during metabolic acid production.

Functions of the Phosphate Buffer System

  • Buffers the intracellular fluid (ICF).
  • Maintains the pH of renal tubular fluid.
  • Promotes the excretion of hydrogen ions in urine.
  • Helps regenerate bicarbonate during renal acid excretion.
  • Supports overall acid–base homeostasis.

Advantages

  • Highly effective in intracellular fluid due to its higher phosphate concentration.
  • Plays a vital role in renal acid excretion.
  • Works efficiently near physiological pH.
  • Complements the bicarbonate and protein buffer systems.

Limitations

  • Present in relatively low concentration in extracellular fluid.
  • Has a limited role in buffering blood compared with the bicarbonate buffer system.
  • Its buffering capacity depends on the availability of phosphate ions.

Clinical Significance

  • The phosphate buffer system is essential for renal acid excretion and helps prevent the accumulation of excess hydrogen ions.
  • It plays an important role in patients with metabolic acidosis, where increased urinary acid excretion is required.
  • Impaired kidney function reduces phosphate buffering, contributing to acid–base disturbances.
  • The amount of phosphate excreted as titratable acid is an important indicator of renal acid excretion.

3. Protein Buffer System

  • The protein buffer system is one of the most abundant and effective buffer systems in the human body.
  • It consists of plasma proteins and intracellular proteins, which help maintain the normal pH of body fluids by accepting or donating hydrogen ions (H⁺).
  • Due to the presence of amino acid side chains, proteins can act as both weak acids and weak bases, making them excellent physiological buffers.

Components of the Protein Buffer System

The protein buffer system is mainly composed of:

  • Plasma proteins (especially albumin)
  • Intracellular proteins present in body cells
  • Amino acid side chains containing carboxyl (-COOH) and amino (-NH₂) groups

These functional groups allow proteins to bind or release hydrogen ions depending on changes in pH.

Mechanism of Action

Proteins are amphoteric, meaning they can act as either an acid or a base.

A. When Excess Acid is Added

When the concentration of hydrogen ions (H⁺) increases:

  • The amino (-NH₂) groups of proteins accept H⁺.
  • Amino groups are converted to -NH₃⁺.
  • This removes excess H⁺ from the solution and minimizes the fall in pH.

Reaction:

Protein-NH₂ + H⁺ → Protein-NH₃⁺

Result: Excess acid is buffered, helping maintain normal pH.

B. When Excess Base is Added

When excess hydroxyl ions (OH⁻) are present:

  • The carboxyl (-COOH) groups of proteins donate H⁺.
  • The released H⁺ combines with OH⁻ to form water.
  • The protein is converted to its negatively charged form (-COO⁻).

Reaction:

Protein-COOH + OH⁻ → Protein-COO⁻ + H₂O

Result: Excess base is neutralized, preventing an increase in pH.

Types of Protein Buffers

1. Plasma Protein Buffer

  • Found in the blood plasma.
  • Albumin is the most important plasma protein buffer because of its high concentration.
  • Helps maintain the pH of extracellular fluid.

2. Intracellular Protein Buffer

  • Present inside body cells.
  • Buffers acids produced during cellular metabolism.
  • Plays a major role in maintaining intracellular pH.

Functions of the Protein Buffer System

  • Maintains the pH of intracellular and extracellular fluids.
  • Buffers hydrogen ions produced during metabolism.
  • Prevents sudden changes in blood and tissue pH.
  • Supports normal enzyme activity and cellular metabolism.
  • Works together with the bicarbonate, phosphate, and hemoglobin buffer systems to maintain acid–base balance.

Advantages

  • Highly abundant throughout the body.
  • Provides rapid buffering against changes in pH.
  • Effective in both intracellular and extracellular compartments.
  • Can buffer both acids and bases due to its amphoteric nature.

Limitations

  • Less effective than the bicarbonate buffer system in regulating blood pH.
  • Buffering capacity decreases when plasma protein levels are low, such as in severe liver disease or malnutrition.

Clinical Significance

  • Albumin is the principal plasma protein responsible for buffering hydrogen ions.
  • Hypoalbuminemia (low serum albumin) reduces the buffering capacity of blood and may contribute to acid–base disturbances.
  • The protein buffer system is particularly important in buffering acids produced within cells during normal metabolism.
  • Alterations in plasma protein concentration can influence the interpretation of acid–base status and the anion gap in clinical practice.

4. Hemoglobin Buffer System

  • The hemoglobin buffer system is one of the most important intracellular buffer systems in the human body.
  • It is present within red blood cells (RBCs) and plays a vital role in maintaining blood pH by binding and releasing hydrogen ions (H⁺) during the transport of oxygen (O₂) and carbon dioxide (CO₂).
  • It works closely with the bicarbonate buffer system to maintain acid–base balance.

Components of the Hemoglobin Buffer System

The hemoglobin buffer system consists of:

  • Hemoglobin (Hb) – oxygen-carrying protein in red blood cells.
  • Deoxyhemoglobin (HHb) – hemoglobin that has released oxygen and can bind hydrogen ions.
  • Oxyhemoglobin (HbO₂) – hemoglobin bound to oxygen.

Hemoglobin acts as a weak acid and can either accept or donate hydrogen ions depending on the body’s pH.

Mechanism of Action

A. In Peripheral Tissues

In metabolically active tissues:

  • Cells produce carbon dioxide (CO₂) during metabolism.
  • CO₂ diffuses into red blood cells and combines with water to form carbonic acid (H₂CO₃).
  • Carbonic acid dissociates into H⁺ and HCO₃⁻.
  • Deoxyhemoglobin (Hb) binds the released H⁺ to form HHb, preventing a fall in blood pH.
  • Bicarbonate (HCO₃⁻) moves into the plasma in exchange for chloride ions (Cl⁻), a process known as the chloride shift.

Reaction:

Hb + H⁺ → HHb

Result: Excess hydrogen ions are buffered, maintaining normal blood pH.

B. In the Lungs

In the pulmonary capillaries:

  • Oxygen binds to deoxyhemoglobin to form oxyhemoglobin (HbO₂).
  • Hemoglobin releases the bound H⁺.
  • H⁺ combines with bicarbonate (HCO₃⁻) to form carbonic acid (H₂CO₃).
  • Carbonic acid is converted into CO₂ and H₂O by the enzyme carbonic anhydrase.
  • Carbon dioxide is exhaled through the lungs.

Reaction:

HHb + O₂ → HbO₂ + H⁺

H⁺ + HCO₃⁻ → H₂CO₃ → CO₂ + H₂O

Result: Carbon dioxide is eliminated, helping maintain acid–base balance.

Functions of the Hemoglobin Buffer System

  • Buffers hydrogen ions (H⁺) within red blood cells.
  • Maintains the normal blood pH (7.35–7.45).
  • Facilitates the transport of carbon dioxide from tissues to the lungs.
  • Assists in the transport and release of oxygen to body tissues.
  • Works in coordination with the bicarbonate buffer system and respiratory system.

Advantages

  • Highly effective because hemoglobin is present in high concentration within red blood cells.
  • Provides rapid buffering during tissue metabolism.
  • Links acid–base balance with oxygen and carbon dioxide transport.
  • Functions continuously during respiration.

Clinical Significance

  • The hemoglobin buffer system is essential for maintaining acid–base balance during normal respiration.
  • Anemia reduces hemoglobin concentration, decreasing the blood’s buffering capacity.
  • Respiratory diseases such as chronic obstructive pulmonary disease (COPD) can impair carbon dioxide elimination, leading to respiratory acidosis.
  • Hemoglobin buffering plays an important role in maintaining normal arterial blood gas (ABG) values.

Comparison of Major Buffer Systems

Buffer System Location Buffer Components Major Function
Bicarbonate Buffer Extracellular fluid H₂CO₃ / HCO₃⁻ Primary extracellular buffer; regulates blood pH
Phosphate Buffer Intracellular fluid and kidneys H₂PO₄⁻ / HPO₄²⁻ Buffers intracellular fluid and urine
Protein Buffer Plasma and cells Plasma proteins, albumin, intracellular proteins Buffers intracellular and extracellular pH
Hemoglobin Buffer Red blood cells Hemoglobin (Hb) Buffers H⁺ and facilitates CO₂ transport

Regulation of Acid–Base Balance

The body maintains the normal arterial blood pH (7.35–7.45) through three major regulatory mechanisms that act at different speeds to prevent significant changes in hydrogen ion (H⁺) concentration.

1. Chemical Buffers (First Line of Defense)

Chemical buffer systems act within seconds to resist sudden changes in pH by binding or releasing hydrogen ions (H⁺).

Major Buffer Systems

  • Bicarbonate buffer – Major extracellular buffer.
  • Phosphate buffer – Important in intracellular fluid and kidneys.
  • Protein buffer – Present in plasma and body cells.
  • Hemoglobin buffer – Major buffer in red blood cells.

Function: Provides immediate protection against changes in blood pH.

2. Respiratory Regulation (Second Line of Defense)

The lungs regulate acid–base balance by controlling the elimination of carbon dioxide (CO₂), which is in equilibrium with carbonic acid (H₂CO₃).

  • Acidosis: Increased breathing (hyperventilation) removes more CO₂, increasing blood pH.
  • Alkalosis: Decreased breathing (hypoventilation) retains CO₂, lowering blood pH.

Response time: Minutes

3. Renal Regulation (Third Line of Defense)

The kidneys provide long-term regulation of acid–base balance by controlling hydrogen ion (H⁺) and bicarbonate (HCO₃⁻) levels.

The kidneys:

  • Reabsorb filtered bicarbonate (HCO₃⁻).
  • Excrete excess hydrogen ions (H⁺) in urine.
  • Generate new bicarbonate to replace that used during buffering.

Response time: Hours to days

Henderson–Hasselbalch Equation

  • The Henderson–Hasselbalch equation is used to determine the pH of blood by relating the concentration of bicarbonate (HCO₃⁻) and the partial pressure of carbon dioxide (PCO₂).
  • It is one of the most important equations for understanding acid–base balance and interpreting arterial blood gas (ABG) reports.

Henderson–Hasselbalch Equation

pH = 6.1 + log [ HCO₃⁻ / (0.03 × PCO₂) ]

Where:

  • pH = Blood pH
  • 6.1 = pKa of carbonic acid
  • HCO₃⁻ = Bicarbonate concentration (22–26 mEq/L)
  • PCO₂ = Partial pressure of carbon dioxide (35–45 mmHg)
  • 0.03 = Solubility coefficient of CO₂ in plasma

Normal Values

Parameter Normal Value
Blood pH 7.35–7.45
HCO₃⁻ 22–26 mEq/L
PCO₂ 35–45 mmHg
HCO₃⁻ : H₂CO₃ Ratio 20 : 1

Clinical Significance

  • Used to assess acid–base balance.
  • Helps in the interpretation of ABG reports.
  • Assists in diagnosing metabolic acidosis, metabolic alkalosis, respiratory acidosis, and respiratory alkalosis.
  • Shows that the kidneys regulate HCO₃⁻, while the lungs regulate PCO₂.

Anion Gap

  • The anion gap (AG) is the difference between the measured cations (positively charged ions) and measured anions (negatively charged ions) in the blood.
  • It is used to identify the cause of metabolic acidosis and assess acid–base disorders.

Formula

Anion Gap (AG) = Na⁺ − (Cl⁻ + HCO₃⁻)

Note: Potassium (K⁺) is usually omitted because its concentration in plasma is very low.

Normal Value

Normal anion gap: 8–12 mEq/L

Clinical Significance

Increased Anion Gap (>12 mEq/L)

An increased anion gap indicates the presence of unmeasured acids in the blood.

Common causes include:

  • Lactic acidosis
  • Diabetic ketoacidosis (DKA)
  • Renal failure (uremia)
  • Poisoning (e.g., methanol or ethylene glycol)

Normal Anion Gap (8–12 mEq/L)

Also known as hyperchloremic metabolic acidosis.

Common causes include:

  • Diarrhea
  • Renal tubular acidosis
  • Excessive saline infusion

Uses of Anion Gap

  • Helps diagnose metabolic acidosis.
  • Differentiates high anion gap from normal anion gap metabolic acidosis.
  • Assists in identifying the underlying cause of acid–base disorders.
  • Commonly used with arterial blood gas (ABG) and electrolyte analysis.

Acid–Base Disorders

 

1. Metabolic Acidosis

  • Metabolic acidosis is an acid–base disorder characterized by a decrease in blood pH (<7.35) due to a primary decrease in plasma bicarbonate (HCO₃⁻).
  • It occurs when the body produces excess acids, loses excessive bicarbonate, or the kidneys are unable to excrete hydrogen ions effectively.

Causes

Common causes of metabolic acidosis include:

  • Diabetic ketoacidosis (DKA)
  • Lactic acidosis
  • Chronic kidney disease (renal failure)
  • Severe diarrhea (loss of bicarbonate)
  • Renal tubular acidosis
  • Poisoning with methanol or ethylene glycol

Clinical Features

Patients with metabolic acidosis may present with:

  • Deep, rapid breathing (Kussmaul respiration)
  • Fatigue and weakness
  • Nausea and vomiting
  • Headache
  • Confusion or drowsiness
  • Severe cases may lead to coma

Arterial Blood Gas (ABG) Findings

Parameter Finding
Blood pH ↓ (<7.35)
HCO₃⁻ ↓
PCO₂ ↓ (respiratory compensation)

Compensation

  • The body compensates by increasing the rate and depth of breathing (hyperventilation), which removes more CO₂ through the lungs.
  • This helps reduce acidity and partially restores the blood pH.

Clinical Significance

  • Metabolic acidosis is a medical emergency if severe and requires prompt diagnosis and treatment.
  • It is commonly seen in diabetic ketoacidosis, renal failure, lactic acidosis, and severe diarrhea.
  • Arterial Blood Gas (ABG) and serum electrolyte analysis (including anion gap) are essential for diagnosis.
  • Treatment focuses on correcting the underlying cause and restoring normal acid–base balance.

2. Metabolic Alkalosis

  • Metabolic alkalosis is an acid–base disorder characterized by an increase in blood pH (>7.45) due to a primary increase in plasma bicarbonate (HCO₃⁻).
  • It occurs when the body loses excessive acid or gains excess bicarbonate.

Causes

Common causes of metabolic alkalosis include:

  • Prolonged vomiting (loss of gastric acid)
  • Excessive use of diuretics
  • Excess intake of bicarbonate or antacids
  • Excessive loss of potassium (hypokalemia)
  • Hyperaldosteronism

Clinical Features

Patients with metabolic alkalosis may present with:

  • Muscle weakness
  • Muscle cramps
  • Tingling or numbness of the hands and feet
  • Dizziness
  • Slow or shallow breathing (hypoventilation)
  • Confusion in severe cases

Arterial Blood Gas (ABG) Findings

Parameter Finding
Blood pH ↑ (>7.45)
HCO₃⁻ ↑
PCO₂ ↑ (respiratory compensation)

Compensation

  • The body compensates by decreasing the rate of breathing (hypoventilation), which retains CO₂.
  • Increased CO₂ forms more carbonic acid, helping lower the blood pH toward normal.

Clinical Significance

  • Metabolic alkalosis commonly occurs due to prolonged vomiting or diuretic therapy.
  • It may lead to electrolyte imbalances, especially hypokalemia.
  • Arterial Blood Gas (ABG) and serum electrolyte analysis are important for diagnosis.
  • Treatment involves correcting the underlying cause, restoring fluid and electrolyte balance, and replacing potassium when required.

3. Respiratory Acidosis

 

  • Respiratory acidosis is an acid–base disorder characterized by a decrease in blood pH (<7.35) due to a primary increase in the partial pressure of carbon dioxide (PCO₂).
  • It occurs when the lungs are unable to eliminate sufficient carbon dioxide because of hypoventilation.

Causes

Common causes of respiratory acidosis include:

  • Chronic obstructive pulmonary disease (COPD)
  • Severe asthma
  • Pneumonia
  • Airway obstruction
  • Respiratory depression due to sedatives or opioid overdose
  • Neuromuscular disorders affecting respiration

Clinical Features

Patients with respiratory acidosis may present with:

  • Shortness of breath
  • Headache
  • Drowsiness
  • Confusion
  • Fatigue
  • Severe cases may lead to coma

Arterial Blood Gas (ABG) Findings

Parameter Finding
Blood pH ↓ (<7.35)
PCO₂ ↑
HCO₃⁻ ↑ (renal compensation)

Compensation

The kidneys compensate by:

  • Increasing the reabsorption of bicarbonate (HCO₃⁻).
  • Increasing the excretion of hydrogen ions (H⁺) in urine.

This helps raise the blood pH toward normal, but compensation requires hours to days.

Clinical Significance

  • Respiratory acidosis commonly occurs in COPD, severe asthma, and respiratory depression.
  • It results from CO₂ retention due to inadequate ventilation.
  • Arterial Blood Gas (ABG) analysis is essential for diagnosis.
  • Treatment focuses on improving ventilation and managing the underlying respiratory disorder.

4. Respiratory Alkalosis

  • Respiratory alkalosis is an acid–base disorder characterized by an increase in blood pH (>7.45) due to a primary decrease in the partial pressure of carbon dioxide (PCO₂).
  • It occurs when excessive CO₂ is lost through hyperventilation.

Causes

Common causes of respiratory alkalosis include:

  • Anxiety or panic attacks
  • Hyperventilation
  • High altitude
  • Fever
  • Early stages of sepsis
  • Mechanical overventilation

Clinical Features

Patients with respiratory alkalosis may present with:

  • Dizziness
  • Light-headedness
  • Tingling or numbness of the hands and feet
  • Muscle cramps or spasms
  • Palpitations
  • Confusion in severe cases

Arterial Blood Gas (ABG) Findings

Parameter Finding
Blood pH ↑ (>7.45)
PCO₂ ↓
HCO₃⁻ ↓ (renal compensation)

Compensation

The kidneys compensate by:

  • Decreasing the reabsorption of bicarbonate (HCO₃⁻).
  • Reducing hydrogen ion (H⁺) excretion, resulting in greater bicarbonate loss in urine.

This helps lower the blood pH toward normal. Renal compensation develops over hours to days.

Clinical Significance

  • Respiratory alkalosis is commonly caused by hyperventilation due to anxiety, high altitude, or fever.
  • It may produce neuromuscular symptoms because of reduced ionized calcium levels.
  • Arterial Blood Gas (ABG) analysis is essential for diagnosis.
  • Treatment focuses on correcting the underlying cause and restoring normal ventilation.

Comparison of Acid–Base Disorders

Disorder Primary Change pH HCO₃⁻ PCO₂
Metabolic Acidosis Decreased HCO₃⁻ ↓ ↓ ↓ (Compensation)
Metabolic Alkalosis Increased HCO₃⁻ ↑ ↑ ↑ (Compensation)
Respiratory Acidosis Increased PCO₂ ↓ ↑ (Compensation) ↑
Respiratory Alkalosis Decreased PCO₂ ↑ ↓ (Compensation) ↓

 

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