Acid-Base Definitions

Arrhenius and Brønsted-Lowery Definitions

Acids and bases are fundamental concepts in chemistry, with several definitions used to describe their behavior in aqueous solutions.

• Arrhenius Acid: Dissolves in water to release H+ ions.

• Arrhenius Base: Dissolves in water to release OH- ions.

• Brønsted-Lowery Acid: A hydrogen-containing molecule that donates a proton (H+) to another molecule.

• Brønsted-Lowery Base: Any substance that can accept a proton.

In water, the H+ ion forms a bond with H2O to create the hydronium ion (H3O+).

Conjugate Acid-Base Pairs

Conjugate acid-base pairs differ by a single proton. When an acid donates a proton, it forms its conjugate base; when a base accepts a proton, it forms its conjugate acid.

• Example: HCl (aq) + H2O (l) → Cl- (aq) + H3O+ (aq)

• Polyprotic Acids: Acids that can donate more than one proton, such as carbonic acid (H2CO3, diprotic) and phosphoric acid (H3PO4, triprotic).

Additional info: Water can act as both an acid and a base (amphoteric).

Self-Ionization of Water and pH

Self-Ionization and the Ionization Constant

Water molecules can self-ionize, acting as both acid and base to form hydronium and hydroxide ions.

• Reaction: H2O (l) + H2O (l) ↔ OH- (aq) + H3O+ (aq)

• Ionization Constant: at 25°C

• Neutral Solution: [H3O+] = [OH-] = 1 × 10-7 M

pH Scale and Calculations

The pH scale is a logarithmic measure of hydronium ion concentration, ranging from acidic (low pH) to alkaline (high pH).

• pH Formula:

• Relationship:

• pOH Formula:

• Relationship:

Example: If [H+] in urine is 1 × 10-6 M, pH = 6.

Strength of Acids and Bases

Acid and Base Strength

The strength of an acid or base is determined by its ability to dissociate in water, measured by the equilibrium constant (Ka for acids).

• Strong Acids: Ka > 1, dissociate completely (e.g., HCl, HBr, HNO3, H2SO4).

• Weak Acids: Ka < 1, do not fully dissociate (e.g., acetic acid, CH3COOH).

• pKa: ; smaller pKa means stronger acid.

Reactions of Acids and Bases

Neutralization and Salt Formation

Acids react with bases in a neutralization reaction to form salts and water. This reaction is highly exothermic and goes to completion.

• General Reaction: Acid + Base → Salt + Water

• Example: HCl (aq) + NaOH (aq) → NaCl (aq) + H2O (l)

• Salt: Cation is conjugate acid of base; anion is conjugate base of acid.

Other Acid-Base Reactions

• Acids react with metals to form metal hydroxides and H2 gas.

• Acids react with metal carbonates/bicarbonates to form salt, water, and CO2 gas.

• Acids react with ammonia and amines to form ammonium salts.

Ka, pKa, and Buffers

Speciation and Buffering

The ratio of acid to conjugate base in solution depends on the pH relative to the pKa. Buffers are mixtures of weak acids and their conjugate bases that resist changes in pH.

• Henderson-Hasselbalch Equation:

• Buffering Capacity: Greatest within 1 pH unit of pKa.

• Example: For H2CO3 (pKa = 6.4) in blood (pH 7.4), ratio of HCO3- to H2CO3 is 10:1.

Buffering in the Body

Major Buffer Systems

The human body uses several buffer systems to maintain pH homeostasis, including phosphate, carbonic acid-bicarbonate, and proteins.

• Phosphate Buffer: H2PO4- ↔ HPO42- + H+ (pKa = 7.2)

• Carbonic Acid Buffer: CO2 (g) + H2O (l) ↔ H2CO3 (aq) ↔ HCO3- (aq) + H+ (aq) (pKa = 6.4)

• Protein Buffers: Ionizable groups in proteins can accept or donate protons.

Regulation of Blood pH

Blood pH is tightly regulated between 7.35 and 7.45. The lungs and kidneys play key roles in maintaining this balance.

• Lungs: Regulate CO2 levels, affecting carbonic acid concentration.

• Kidneys: Excrete excess H+ and return HCO3- to extracellular fluids.

• Red Blood Cells: Convert CO2 to HCO3- for transport.

Acidosis and Alkalosis

Disorders of pH homeostasis include acidosis (pH < 7.35) and alkalosis (pH > 7.45), which can be caused by respiratory or metabolic changes.

• Acidosis: Increased [H+] or decreased [HCO3-], or increased [CO2].

• Alkalosis: Decreased [H+] or increased [HCO3-], or decreased [CO2].

Additional info: Hyperventilation causes loss of CO2, shifting equilibrium and increasing pH (respiratory alkalosis).

Summary

• Acids donate protons; bases accept protons.

• Water can act as both acid and base.

• pH is a logarithmic measure of hydronium ion concentration.

• Strong acids/bases dissociate completely; weak acids/bases do not.

• Buffers resist changes in pH and are crucial for biological systems.

• Body buffer systems include phosphate, carbonic acid-bicarbonate, and proteins.

• Lungs and kidneys regulate blood pH; loss of homeostasis leads to acidosis or alkalosis.