Acid-Base Reactions, Buffers, Indicators, and Titrations

Acid-Base Reactions

Overview of Acid-Base Reactions

Acid-base reactions are fundamental chemical processes where an acid reacts with a base to produce water and a salt. These reactions are also known as neutralization reactions because the acidic and basic properties are neutralized.

• Acid: A substance that donates a proton (H+).

• Base: A substance that accepts a proton or provides hydroxide ions (OH-).

• Salt: An ionic compound formed from the cation of the base and the anion of the acid.

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

General Equation:

$\mathrm{Acid} + \mathrm{Base} \rightarrow \mathrm{Salt} + \mathrm{Water}$

Types of Acid-Base Reactions

• Strong Acid + Strong Base: Both fully ionize/dissociate; reaction goes to completion. Example: HCl(aq) + NaOH(aq) → H2O(l) + NaCl(aq)

• Strong Acid + Weak Base: Acid fully ionizes; base may not fully dissociate, especially if insoluble. Reaction goes to completion. Example: 3HBr(aq) + Al(OH)3(s) → AlBr3(aq) + 3H2O(l)

• Weak Acid + Strong Base: Acid only partially ionizes; base fully dissociates. Reaction goes to completion. Example: HC2H3O2(aq) + NaOH(aq) → NaC2H3O2(aq) + H2O(l)

• Weak Acid + Weak Base: Both only partially ionize/dissociate; reaction does not go to completion and is uncommon.

Buffers

Definition and Function of Buffers

Buffer: A solution that resists changes in pH when moderate amounts of acid or base are added. Buffers are essential for stabilizing pH in biological and chemical systems.

• Made from a weak acid or base and its salt.

• Used to maintain pH in acid-base reactions.

• Example: NH3(aq) and NH4Cl(aq) form a buffer system.

• Example: HC2H3O2(aq) and NaC2H3O2(aq) form a buffer system.

Buffer Reactions:

• Adding strong acid: NH3(aq) + H+(aq) → NH4+(aq)

• Adding strong base: NH4+(aq) + OH-(aq) → NH3(aq) + H2O(l)

• Adding strong acid: C2H3O2-(aq) + H+(aq) → HC2H3O2(aq)

• Adding strong base: HC2H3O2(aq) + OH-(aq) → C2H3O2-(aq) + H2O(l)

Real World Example: Blood acts as a buffer to maintain physiological pH.

Indicators

Definition and Use of Indicators

Indicators: Substances (acids or bases) that change color as they gain or lose hydrogen ions. They are used to visually determine the pH of a solution.

• Each indicator has a specific pH range for color change.

• Example: Phenol red is yellow at pH 7 and turns red at pH 8.

• The universal indicator covers nearly the entire pH scale.

Titrations

Definition and Process of Titration

Titration: A quantitative technique used to determine the molarity of an acid or base by reacting it with a solution of known concentration.

• Controlled addition of a solution of known concentration to a solution of unknown concentration.

• Measurement of the volume required for complete reaction allows calculation of the unknown molarity.

General Titration Equation:

$\mathrm{M}_1 \times \mathrm{V}_1 = \mathrm{M}_2 \times \mathrm{V}_2$

Where M = molarity, V = volume, and subscripts 1 and 2 refer to acid and base, respectively.

Application: Used in laboratory settings to analyze the concentration of acids or bases in solutions.

Summary Table: Types of Acid-Base Reactions

Acid Type	Base Type	Reaction Completion	Example
Strong	Strong	Complete	HCl + NaOH → H2O + NaCl
Strong	Weak	Complete	3HBr + Al(OH)3 → AlBr3 + 3H2O
Weak	Strong	Complete	HC2H3O2 + NaOH → NaC2H3O2 + H2O
Weak	Weak	Incomplete	Uncommon