Acids and Bases: Fundamental Concepts

Acid-Base Equilibria

Acid-base equilibria describe the reversible transfer of protons between substances. Substances are classified as acids or bases based on their ability to donate or accept protons.

• Brønsted-Lowry acid: A proton donor. Example: Ammonium ion (NH4+).

• Brønsted-Lowry base: A proton acceptor. Example: Hydroxide ion (OH-).

Acid and Base Strength

Definitions and Examples

Acid strength refers to the degree of dissociation in solution, not the concentration. Strong acids and bases dissociate completely, while weak acids and bases only partially dissociate.

• Strong acid: Completely dissociates into ions in solution (pH 0-1).

• Weak acid: Only slightly dissociates into ions in solution (pH 3-5).

• Strong base: Completely dissociates, typically pH 12-14.

• Weak base: Partially dissociates, typically pH 8-11.

Example: Hydrochloric acid (HCl) is a strong acid; ethanoic acid (CH3COOH) is a weak acid.

pH and Its Calculation

Definition and Formula

The pH scale measures acidity and alkalinity, ranging from 0 (most acidic) to 14 (most basic). It is a logarithmic scale based on the concentration of hydrogen ions in solution.

• pH formula:

• Hydrogen ion concentration from pH:

These equations allow calculation of pH from [H+] and vice versa. For strong acids, [H+] equals the acid concentration.

Ionic Product of Water

Equilibrium Constant (Kw)

Water self-ionizes to a small extent, forming an equilibrium between H+ and OH- ions. The equilibrium constant for this process is Kw.

• Formula:

• At 25°C,

• As temperature increases, increases (endothermic reaction), making water more acidic.

Equation:

Weak Acids and Bases: Equilibrium Constants

Dissociation Constant (Ka) and pKa

Weak acids and bases only partially dissociate, forming an equilibrium mixture. The equilibrium dissociation constant (Ka) quantifies this behavior.

• General equation:

• pKa formula:

• These relationships are used to calculate the pH of weak acids and bases, depending on concentrations and equilibrium constants.

Titration Curves

Acid-Base Titration and Neutralisation Points

Titration curves show how pH changes as acid and base react. The neutralisation point (or equivalence point) is identified by a large vertical section in the curve.

• Strong acid + strong base: Neutralisation at pH 7.

• Strong acid + weak base: Neutralisation at pH < 7 (more acidic).

• Weak acid + strong base: Neutralisation at pH > 7 (more basic).

• Weak acid + weak base: Neutralisation point is hard to determine, usually near pH 7.

Example: The titration curve for a strong acid and strong base shows a sharp rise at the equivalence point.

Indicators in Acid-Base Reactions

Types and Colour Changes

Indicators are used to detect pH changes during titrations. The two most common indicators are methyl orange and phenolphthalein.

Selection of the correct indicator depends on the expected pH at the equivalence point.

Buffer Solutions

Buffer Action and Composition

Buffer solutions resist changes in pH when small amounts of acid or base are added. Acidic buffers contain a weak acid and its salt; basic buffers contain a weak base and its salt.

• Example: Ethanoic acid and sodium ethanoate form an acidic buffer.

• Definition: A solution able to resist changes in pH when small volumes of acid or base are added.

Calculations Involving Buffers

• Acid + Base: Find moles of each species, calculate concentrations at equilibrium, use total volume, and use Ka to find [H+] and pH.

• Acid + Salt: Find moles of salt, use Ka to find pH.

Effect of Adding Small Volumes

• Adding acid increases [H+], making the solution slightly more acidic.

• Adding base increases [OH-], making the solution slightly more basic.

Uses of Buffers

Buffers are important in biological systems to keep pH stable, which is essential for enzyme activity and metabolic reactions.

• Example: Blood contains buffer systems to maintain pH for proper physiological function.

Additional info: Buffer solutions are widely used in laboratory and industrial processes where pH control is critical.