Acid-Base Balance and pH Regulation in Human Physiology

Acid-Base Balance in Human Physiology

Acid-Base Reactions and Examples

Acid-base balance is crucial for maintaining homeostasis in the human body. The body uses various chemical reactions to regulate the concentration of acids and bases, especially in the blood and tissues.

• Example 1: Ammonium ion formation

  • $\mathrm{NH_3} + \mathrm{H^+} \rightarrow \mathrm{NH_4^+}$

  • This reaction helps remove excess hydrogen ions (acid) from the body.

• Example 2: Bicarbonate formation

  • $\mathrm{H^+} + \mathrm{HCO_3^-} \rightarrow \mathrm{H_2CO_3}$

  • Bicarbonate ions act as buffers to neutralize excess acid.

• Role of the Kidney: The kidneys utilize synthesis reactions to remove excess acid (H+) from the body, helping to maintain acid-base balance.

pH: Acid-Base Concentration

The pH scale is a measure of hydrogen ion (H+) concentration in a solution, indicating its acidity or alkalinity.

• Scale: Ranges from 0 to 14

  • pH = 7: Neutral (pure water)

  • pH < 7: Acidic (higher concentration of H+)

  • pH > 7: Basic/Alkaline (lower concentration of H+)

• Definition: pH is defined mathematically as:

  • $\mathrm{pH} = -\log_{10}[\mathrm{H^+}]$

Concept of Physiological pH

Maintaining a stable pH is essential for proper cellular function and overall health.

• Normal pH Range for Human Blood: 7.35–7.45

• Clinical Implications:

  • pH < 7.35: Acidosis (excess acid in the blood)

  • pH > 7.45: Alkalosis (excess base in the blood)

  • Both conditions can disrupt enzyme activity and metabolic processes.

Buffers

Buffers are chemicals that help regulate pH by neutralizing excess acids or bases. They are vital for maintaining the body's acid-base homeostasis.

• Definition: A buffer is a solution that resists changes in pH when small amounts of acid or base are added.

• Example: Carbonic Acid-Bicarbonate System

  • This system is the primary buffer in human blood.

  • The reaction is as follows:

  $\mathrm{CO_2} + \mathrm{H_2O} \leftrightarrow \mathrm{H_2CO_3} \leftrightarrow \mathrm{H^+} + \mathrm{HCO_3^-}$

  • Carbon dioxide (CO2) combines with water (H2O) to form carbonic acid (H2CO3), which can dissociate into hydrogen ions (H+) and bicarbonate ions (HCO3-).

  • This reversible reaction helps buffer changes in blood pH.

Summary Table: Blood pH Classification

Condition	Blood pH Range	Description
Acidosis	< 7.35	Excess acid in blood; can impair organ function
Normal	7.35–7.45	Optimal range for human physiology
Alkalosis	> 7.45	Excess base in blood; can disrupt metabolic processes

Additional info: The carbonic acid-bicarbonate buffer system is especially important in respiratory and renal physiology, as it allows the body to rapidly respond to changes in acid-base status through breathing and kidney function.