pH and buffer

INTRODUCTION

To achieve acid-base balance, there must be a balance between the intake or production of hydrogen ions and net removal of hydrogen ions from the body.

An acid is defined as a substance that releases protons or hydrogen ions (H⁺) e.g:

A base is a substance that accepts protons or hydrogen ions e.g:

PROPERTIES OF WATER

Water is the solvent in which most biochemical reactions take place.

Two properties of water are especially relevant:

Water is a POLAR MOLECULE

Water molecule is bent and not linear so the distribution of charge is asymmetric.

Oxygen nucleus draws electrons away from the two hydrogen nuclei, which leaves the region around each hydrogen atom with a net positive charge.

Water molecule is thus an electrically polar structure.

Water is HIGHLY COHESIVE

Water molecules interact strongly with one another through hydrogen bonds.

These interactions are apparent in the structure of ice.

Networks of hydrogen bonds hold the structure together and link molecules in liquid water.

pH

When weak acids or bases dissolve in water to contribute H⁺ by ionizing or consume H⁺ by being protonated.

These processes are governed by equilibrium constant (Kₑq).

The total H⁺ concentration is expressed as the pH of the solution.

Water molecule has slight tendency to undergo reversible ionization to yield hydrogen ion and hydroxide ion.

H₂O ↔ H⁺ + OH⁻

Cont…

The concept of pH was first introduced by the Danish Chemist Sorenson in 1909.

P means “puissance d” and H means “Hydrogen”.

The French word means strength/power of hydrogen.

pH can be define as the negative logarithm of the hydrogen ion concentration.

pH = -log [H⁺]

pH is a unit that measures the degree of acidity or alkalinity of a solution.

It is measured on a scale of 0 to 14.

pH VALUE

The pH value of a substance is directly related to the ratio of the hydrogen ion and hydroxyl ion concentrations.

Low pH values correspond to high concentrations of H⁺ and high pH values correspond to low concentrations of H⁺.

If the H⁺ concentration is higher than OH⁻, the material is acidic and the pH is less than 7.

If the OH⁻ concentration is higher than H⁺, the material is basic and the pH is more than 7.

If the H⁺ concentration is equal to OH⁻, the material is neutral and the pH is equal to 7.

BUFFER

A buffer solution is a solution which resists changes in pH when a small amount of acid or base is added.

Typically a mixture of a weak acid and a salt of its conjugate base or weak base and a salt of its conjugate acid.

TYPES OF BUFFERS ACIDIC BUFFERS

Solution of a mixture of a weak acid and a salt of this weak acid with a strong base.

E.g.

CH₃COOH + CH₃COONa

(Weak acid) — (Salt)

BASIC BUFFERS

Solution of a mixture of a weak base and a salt of this weak base with a strong acid.

E.g.

NH₄OH + NH₄Cl

(Weak base) — (Salt)

HOW BUFFERS WORK ACIDIC BUFFER: Equilibrium between acid and base Example: ACETATE BUFFER

CH₃COOH (aq) ⇌ CH₃COO⁻ (aq) + H⁺ (aq)

CH₃COONa (aq) ⇌ CH₃COO⁻ (aq) + Na⁺ (aq)

If more H⁺ is added to this solution, the equilibrium shifts to the left, absorbing H⁺, so the [H⁺] changes only slightly and the pH remains nearly constant.

If more OH⁻ is added, H⁺ is removed to form water, then the equilibrium shifts to the right, releasing H⁺ to keep the pH constant.

HOW BUFFERS WORK BASIC BUFFER: Equilibrium between acid and base Example: AMMONIUM BUFFER

NH₄OH (aq) ⇌ NH₄⁺ (aq) + OH⁻ (aq)

NH₄Cl (aq) ⇌ NH₄⁺ (aq) + Cl⁻ (aq)

If more H⁺ is added, it reacts with OH⁻ to form water resulting in the removal of OH⁻ from the equilibrium. This shifts the equilibrium to the right, to replace the OH⁻ that was consumed.

If more OH⁻ is added, the NH₄⁺ ions react with the excess OH⁻ to NH₄OH. This shifts the equilibrium to the left, so the [OH⁻] changes only slightly and the pH remains nearly constant.

HENDERSON-HASSELBALCH EQUATION

Lawrence Joseph Henderson wrote an equation, in 1908, describing the use of carbonic acid as a buffer solution.

Karl Albert Hasselbalch later re-expressed that formula in logarithmic terms, resulting in the Henderson–Hasselbalch equation.

Henderson-Hasselbalch Equation Derivation

Kₐ = [H⁺][A⁻] / [HA]

Take the -log on both sides:

-log Kₐ = -log [H⁺] - log [A⁻]/[HA]

Therefore:

pH = pKₐ + log [A⁻]/[HA]

Also:

pH = pKₐ + log [Proton acceptor]/[Proton donor]

HA ⇌ H⁺ + A⁻

pKₐ = pH - log [A⁻]/[HA]

Apply:

p(x) = -log(x)

and finally solve for pH…

log xy = log x + log y

The greater the buffer capacity the less the pH changes upon addition of H⁺ or OH⁻.

CHOOSING THE RIGHT BUFFER

Choose a buffer whose pKₐ is closest to the desired pH.

pKₐ is the pH at which an acid is 50% dissociated (ionized) and 50% undissociated (unionized).

pH should be within pKₐ ± 1.

BUFFER SYSTEM IN BODY FLUIDS

ACIDS IN THE BODY VOLATILE ACIDS

Produced by oxidative metabolism of CHO, Fat, Protein.

Excreted through LUNGS as CO₂ gas.

FIXED ACIDS

Acids that do not leave solution, once produced they remain in body fluids until eliminated by KIDNEYS.

E.g: Sulfuric acid, phosphoric acid, Organic acids.

Are generated during catabolism of:

• Amino acids

• Phospholipids

• Nucleic acids

CARBONIC ACID–BICARBONATE BUFFER SYSTEM Carbon Dioxide

Most body cells constantly generate carbon dioxide.

Most carbon dioxide is converted to carbonic acid, which dissociates into H⁺ and a bicarbonate ion.

Prevents changes in pH caused by organic acids and fixed acids in ECF.

Cannot protect ECF from changes in pH that result from elevated or depressed levels of CO₂.

Functions only when respiratory system and respiratory control centers are working normally.

Ability to buffer acids is limited by availability of bicarbonate ions.

CARBONIC ACID–BICARBONATE BUFFER SYSTEM

The carbonic acid bicarbonate buffer is the most important buffer system.

Carbonic acid, H₂CO₃, acts as the weak acid.

Hydrogen carbonate, HCO₃⁻, acts as the conjugate base.

Increase in H⁺(aq) ions is removed by HCO₃⁻(aq).

The equilibrium shifts to the left and most of the H⁺(aq) ions are removed.

Increase in OH⁻(aq) ions results in reaction with H⁺(aq) ions to form water.

H₂CO₃ dissociates, shifting the equilibrium to the right, restoring most of the H⁺(aq) ions.

THE HEMOGLOBIN BUFFER SYSTEM

CO₂ diffuses across RBC membrane.

No transport mechanism required.

As carbonic acid dissociates:

Bicarbonate ions diffuse into plasma

In exchange for chloride ions (chloride shift).

Hydrogen ions are buffered by hemoglobin molecules.

Is the only intracellular buffer system with an immediate effect on ECF pH.

Helps prevent major changes in pH when plasma PCO₂ is rising or falling.

PHOSPHATE BUFFER SYSTEM

Consists of anion H₂PO₄⁻ (a weak acid) (pKₐ-6.8).

Works like the carbonic acid–bicarbonate buffer system.

Important in buffering pH of ICF.

The phosphate buffer system (HPO₄²⁻/H₂PO₄⁻) plays a role in plasma and erythrocytes.

H₂PO₄⁻ + H₂O ↔ HPO₄²⁻ + H₃O⁺

Acid will react with monohydrogen phosphate to form dihydrogen phosphate (equilibrium shifts to left).

H₂PO₄⁻ + H₂O ← HPO₄²⁻ + H₃O⁺

Base will be neutralized by dihydrogen phosphate (equilibrium shifts to right).

H₂PO₄⁻ + H₂O → HPO₄²⁻ + H₃O⁺

PROTEINS BUFFER SYSTEM

Proteins contain –COO⁻ groups, which, like acetate ions (CH₃COO⁻), can act as proton acceptors.

Proteins also contain –NH₃⁺ groups, which, like ammonium ions (NH₄⁺), can donate protons.

If acid comes into blood, hydronium ions can be neutralized by the –COO⁻ groups.

COO⁻ + H₃O⁺ → COOH + H₂O

If base is added, it can be neutralized by the –NH₃⁺ groups.

NH₃⁺ + OH⁻ → NH₂ + H₂O

FOUR BASIC TYPES OF ACID-BASE IMBALANCE

1. Metabolic Acidosis

2. Metabolic Alkalosis

3. Respiratory Acidosis

4. Respiratory Alkalosis

METABOLIC ACIDOSIS

Decrease in blood pH due to a decrease in bicarbonate (HCO₃⁻) or an increase in acid.

Causes

• Lactic Acidosis

• Ketoacidosis (Diabetic, Alcoholic, Starvation)

• Renal failure (Acute And Chronic)

• Toxins (ethylene glycol, methanol, salicylates, propylene glycol)

Body’s Compensation

Lungs increase breathing rate (hyperventilation) to remove CO₂ and reduce acidity.

Signs and Symptoms

• Deep and rapid breathing

• Weakness

• Fatigue

• Nausea

• Vomiting

• Coma in severe cases

METABOLIC ALKALOSIS

Increase in blood pH due to excess bicarbonate (HCO₃⁻) or loss of hydrogen ions (acid).

Causes

• Prolonged vomiting

• Excessive antacid use

• Diuretic therapy

• Potassium depletion

• Gastric suction

Body’s Compensation

Lungs decrease breathing rate (hypoventilation) to retain CO₂ and increase acidity.

Signs and Symptoms

• Muscle cramps

• Muscle twitching

• Confusion

• Cardiac arrhythmias

RESPIRATORY ACIDOSIS

Decrease in blood pH due to retention of carbon dioxide (CO₂).

Causes

• Chronic obstructive pulmonary disease

• Severe asthma

• Pneumonia

• Drug abuse

• Airway obstruction

Body’s Compensation

Kidneys retain bicarbonate (HCO₃⁻) and excrete more hydrogen ions.

Signs and Symptoms

• Slow breathing

• Headache

• Drowsiness

• Confusion

• Coma in severe cases

RESPIRATORY ALKALOSIS

Increase in blood pH due to excessive loss of carbon dioxide (CO₂).

Causes

• Anxiety or panic attack

• Fever

• Pain

• High altitude

• Over ventilation

Body’s Compensation

Kidneys excrete bicarbonate (HCO₃⁻) and retain more hydrogen ions.

Signs and Symptoms

• Dizziness

• Tingling of fingers and lips

• Muscle spasms

• Fainting

ACID BASE DISORDERS SUMMARY