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: Differ by a single proton (H+).

• Water: Can act as both an acid and a base (amphoteric).

Polyprotic Acids

Polyprotic acids can donate more than one proton. Examples include carbonic acid (H2CO3, diprotic) and phosphoric acid (H3PO4, triprotic).

• Conjugate base: Formed when an acid loses a proton.

• Conjugate acid: Formed when a base gains a proton.

Self-Ionization of Water and pH

Self-Ionization and pH Calculation

Water can self-ionize, producing equal concentrations of hydronium and hydroxide ions.

• Self-ionization reaction:

• Ionization constant of water: at 25°C

• Neutral solution:

• Acidic solution:

• Basic solution:

The pH scale is logarithmic and measures the concentration of H+ ions:

• pH formula:

• Relationship:

Relationship Between pH and pOH

• pOH formula:

• Relationship:

Strength of Acids and Bases

Acid and Base Strength

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

• Strong acids: Completely dissociate in water (Ka > 1). Examples: HCl, HBr, HNO3, H2SO4.

• Weak acids: Partially dissociate (Ka < 1). Example: Acetic acid (CH3COOH).

• pKa: ; smaller pKa means stronger acid.

Reactions of Acids and Bases

Neutralization Reactions

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

• General equation:

• Example:

Other Acid-Base Reactions

• Acids react with metals to produce hydrogen gas.

• Acids react with metal carbonates/bicarbonates to produce CO2, water, and a salt.

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

Ka, pKa, and Buffers

Buffer Systems and the Henderson-Hasselbalch Equation

Buffers are mixtures of a weak acid and its conjugate base that resist changes in pH when small amounts of acid or base are added.

• Henderson-Hasselbalch equation:

• Buffering capacity is greatest when pH is within 1 unit of pKa.

• At pH = pKa, the solution contains equal amounts of acid and conjugate base.

Buffering in the Body

Major Buffer Systems

The human body uses several buffer systems to maintain pH homeostasis, including:

• Phosphate buffer: Dihydrogen phosphate (H2PO4-) and hydrogen phosphate (HPO42-).

• Carbonic acid-bicarbonate buffer: Carbonic acid (H2CO3) and bicarbonate (HCO3-).

• Protein buffers: Ionizable groups in proteins.

Blood Buffering and Homeostasis

Blood pH is tightly regulated between 7.35 and 7.45. The carbonic acid-bicarbonate system is central to this regulation, with CO2 and HCO3- interconverted in red blood cells and lungs.

• CO2 + H2O ⇄ H2CO3 ⇄ HCO3- + H+

• Excess H+ is buffered by hemoglobin and excreted by the kidneys.

• Respiratory and metabolic changes can cause acidosis (pH < 7.35) or alkalosis (pH > 7.45).

Summary Table: Acid-Base Buffer Systems in the Body

Key Equations and Concepts

• pH calculation:

• pOH calculation:

• Relationship:

• Henderson-Hasselbalch:

Summary

• Acids donate protons; bases accept protons.

• Water can act as both an acid and a base.

• pH is a measure of acidity/basicity; buffers resist changes in pH.

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

• Homeostasis is maintained by respiratory and renal regulation of acid-base balance.