chapter 9

Introduction

• The pH scale categorizes substances as:

  • Acidic: pH < 7

  • Neutral: pH = 7

  • Basic (Alkaline): pH > 7

• Each body fluid has a specific pH for optimal functioning.

• Homeostatic regulation of pH is referred to as acid–base balance.

• Buffer systems, respiratory functions, and renal mechanisms are involved in regulating pH and maintaining acid–base balance.

• The balance of integrated body systems is crucial for maintaining physiological pH levels.

• The body's systems are continually active and responsive to minor changes that could jeopardize pH levels.

• There is redundancy in these systems to ensure cells function optimally.

• This chapter expands on fluid and electrolyte balance and explores the mechanisms of acid–base balance.

• Clinical models in Module 2 will highlight specific alterations to facilitate application of concepts learned.

pH-Bicarbonate Diagram

• The diagram illustrates the relationship between:

  • Arterial pH

  • Bicarbonate levels

• Normal values for pH and bicarbonate are depicted within an elliptical region.

• Values left of normal indicate acidemia (causes include respiratory acidosis or metabolic acidosis).

• Values to the right of normal indicate alkalemia (causes include respiratory alkalosis or metabolic alkalosis).

• The diagram summary is based on information from Rhoades & Bell (2012).

Module 1: Acid–Base Imbalance

• Regulation of acids and bases is essential for metabolic activities.

• A narrow physiological pH range is vital for cellular function.

Regulation of Acid and Base

• Acids: Substances that donate hydrogen ions (H+).

• Bases: Substances that accept hydrogen ions.

• Weak Acids in Plasma: Examples include albumin and inorganic phosphorus.

• Strong Ions: Dissociate almost completely in solution:

  • Cations (strong): Na+, K+, Ca2+, Mg2+

  • Anions (strong): Cl−, C3H5O3 (lactate)

• The extracellular fluid maintains a regulated ratio of acids and bases which is measured clinically by pH.

• Hydrogen Ion Concentration: Inversely related to pH:

  • Low pH = High [H+]

  • High pH = Low [H+]

• Anion Gap: A clinical calculation indicating acid–base balance, defined as:

  • $ext{Anion Gap} = [ ext{Na}^+] - ([ ext{Cl}^-] + [ ext{HCO}_3^-])$

• The normal cation concentration of sodium is $140 ext{ mEq/L}$ and anions are $102 ext{ mEq/L}$ (Cl−) and $26 ext{ mEq/L}$ (HCO3−).

• Example calculation of anion gap:

  • Anion Gap = $140 ext{ mEq/L} - (102 ext{ mEq/L} + 26 ext{ mEq/L})$ = $12 ext{ mEq/L}$

• Normal range for anion gap is between $10 ext{ mEq/L}$ and $14 ext{ mEq/L}$.

• Changes in the anion gap can indicate acid-base disturbances:

  • Increased organic acids replace bicarbonate in metabolic acidosis.

  • Rise in Cl− leading to normal anion gap due to hyperchloremic acidosis.

Buffer Systems

• Buffer systems react to prevent significant changes in pH by exchanging stronger acids and bases for weaker ones.

• Types of buffer systems:

  1. Plasma Buffer System: Reaction time in seconds for hydrogen ion concentration response.

  2. Respiratory Buffer System: Operates within minutes to excrete CO2 through changing respiratory rates.

  3. Renal Buffer System: Reactive time spans hours to days; manages acids, bases, and ions.

Plasma Buffer Systems

• Primary plasma buffer systems:

  • Bicarbonate Buffer System

  • Protein Buffer System

  • H+/K+ Cation Exchange System

• Buffers’ effectiveness depends on their ability to manage free H+ ions by binding or releasing them, addressing pH changes promptly.

• Key reaction involving water, carbon dioxide, carbonic acid, and bicarbonate, catalyzed by carbonic anhydrase:

  • Reaction: ext{H$2$O} + ext{CO}2
    ightleftharpoons ext{H}2 ext{CO}3
    ightleftharpoons ext{HCO}_3^- + ext{H}^+

• The rapid nature of these responses preserves pH until other systems engage.

Bicarbonate Buffer System

• The primary extracellular buffer; involves:

  • Weak acid: $ext{H}<em>2 ext{CO}</em>3$ (carbonic acid)

  • Weak base: $ext{NaHCO}_3$ (sodium bicarbonate)

• Exchanges occur where strong acids (e.g., HCl) are buffered by weaker forms (e.g., $ext{H}<em>2 ext{CO}</em>3$) and vice versa.

Protein Buffer System

• Proteins serve as the principal buffer, maintaining pH:

  • Response mechanisms can simultaneously donate and accept H+ ions, functioning as amphoteric species.

• H+ and CO2 can diffuse across plasma membranes, stabilizing conditions.

Potassium–Hydrogen Exchange

• Cation exchange between K+ and H+ is crucial in acid–base regulation:

  • Excess H+ entering cells promotes K+ movement into the extracellular space, potentially causing hyperkalemia.

  • Conversely, hypokalemia leads to efflux of K+ from cells, promoting increased extracellular H+ and lowering pH.

Respiratory Buffer System

• Key players in managing H2CO3 derived from CO2 levels:

  • Increased pH (alkalosis) leads to decreased respiratory rate to retain CO2.

  • Decreased pH (acidosis) increases respiratory rate to eliminate CO2.

• Although effective, it cannot completely restore homeostasis, allowing for renal correction.

Renal Buffer System

• Primary regulators of acid and base balance; manage non-volatile acids:

  • Mechanisms include hydrogen ion elimination/bicarbonate conservation, tubular buffer systems, and anion exchanges.

Hydrogen Ion Elimination/Bicarbonate Conservation

• The kidneys control pH by excretion of H+ ions; excess acid leads to reabsorption and regulation of bicarbonate (HCO3−).

• Maintains blood pH within the range of 7.35 to 7.45, requiring a 20:1 ratio of bicarbonate to carbonic acid:

  • Ratio determinant of pH rather than absolute concentrations.

Tubular Buffer Systems

• Phosphate Buffer System: Involves unbuffered H+ binding with HPO4²⁻ to form H2PO4⁻ and facilitate ecretion of H+.

• Ammonia Buffer System: Involves exchanges between NH4+ and HCO3− aiding in H+ excretion via urine.

Potassium–Hydrogen Exchange

• Similar to plasma buffering; hypokalemia triggers K+ release for H+ absorption, impacting blood pH.

Chloride–Bicarbonate Exchange

• Renal regulation of bicarbonate through anion exchanges with chloride, affecting acid-base balance.

Altered Acid–Base Balance

• Acid and base disorders can originate from respiratory or metabolic dysfunction.

• Focus on metabolic alterations leading to changes in HCO3− levels affecting pH:

  • Metabolic Acidosis: Reduction in HCO3− causing decreased pH.

  • Metabolic Alkalosis: Increased HCO3− leading to increased pH.

• Compensatory mechanisms from kidneys conserve HCO3− until pH returns to normal.

Metabolic Acidosis

• Characterization by HCO3− base deficit linked to increased strong anions or weak acids.

• Contributing mechanisms include:

  • Increased production of non-volatile acids (fasting, ketoacidosis).

  • Decreased renal acid secretion.

  • Loss of bicarbonate due to diarrhea.

• Clinical manifestations include:

  • Symptoms: Anorexia, nausea, weakness, confusion, coma.

  • Laboratory findings: Decreased pH (< 7.35), HCO3− (< 24 mEq/L).

Metabolic Alkalosis

• Characterization by excess HCO3− leading to increased pH.

• Contributing mechanisms include:

  • Decreased H+ ions, increased bicarbonate, or chloride loss.

• Clinical symptoms vary from asymptomatic to signs related to electrolyte imbalance.

• Laboratory findings: Increased pH (> 7.45), HCO3− (> 31 mEq/L).

Module 2: Clinical Models of Acid-Base Imbalance

• Present clinical models demonstrate fluid, electrolyte, and acid-base balance interdependencies.

HAART-Associated Acidosis

• Highly Active Antiretroviral Therapy (HAART) used for HIV has shown effective results but with adverse effects:

  • Hyperlactatemia: Hemoglobin elevation from NRTIs causing mitochondrial dysfunction.

  • Associated symptoms include mild hyperlactatemia (asymptomatic) or lactic acidosis (serious).

• Pathophysiology involves NRTIs inhibiting DNA polymerase leading to lactic acid buildup, resulting in decreased pH.

• Clinical manifestations range from mild nausea to severe coma.

• Diagnostic criteria include lactate levels and liver function tests.

• Treatment may involve halting NRTI use and hydration techniques.

Renal Tubulopathy

• Hypokalemic Salt-Losing Tubulopathies (SLTs): Autosomal recessive disorders showing alterations in metabolic alkalosis.

• Major types include Bartter Syndrome and Gitelman Syndrome characterized by renal salt reabsorption alterations leading to hypokalemic metabolic alkalosis.

• Diagnostic criteria includes renal ultrasound and electrolyte measurements; treatment focuses on normalizing fluid and electrolyte levels.