Acid-Base Reactions

Brønsted-Lowry Theory

The Brønsted-Lowry theory defines acids and bases based on their ability to donate or accept protons (H+). An acid is a chemical species that can release a proton, while a base is a species that can accept a proton.

• Acid: AH → H+ + A-

• Base: H+ + A- → AH

• Example: NH4+ → NH3 + H+

• Example: C6H5OH → C6H5O- + H+

• Example: CH3CO2H → CH3CO2- + H+

• Example: HCl → H+ + Cl-

Bond Polarity and Acid Strength

The strength of an acid is influenced by the polarity of the H–A bond and the electronegativity of atom A. Highly electronegative atoms and large electronegativity differences create polar bonds, which generally lead to stronger acids.

• H–C: Weakly polar, weak acid

• H–A (A = highly electronegative): Stronger acid

Strong and Weak Acids

Acids are classified as strong or weak based on their degree of dissociation in water. Strong acids dissociate completely, while weak acids only partially dissociate.

• Strong Acids: HCl, HBr, HI, HNO3, HClO4

• Weak Acids: HF, CH3COOH, H3PO4, H2CO3, H2SO3, HClO, H2S, HCOOH, C6H5OH

Acid Classification Table

The following table summarizes the classification of common acids:

Acid Dissociation and Chemical Equilibrium

Acid Dissociation in Water

When an acid is dissolved in water, it establishes a chemical equilibrium between the acid and water. The acid dissociates to form its conjugate base and hydronium ion (H3O+).

• General Reaction: AH + H2O ⇌ H3O+ + A-

• Monoacid Example: C6H5OH + H2O ⇌ C6H5O- + H3O+

• Polyacid Example: H2SO4 + H2O → H3O+ + HSO4- HSO4- + H2O → H3O+ + SO42-

Electrolyte Strength

• Strong acids and bases: Dissociate completely in water, forming strong electrolytes.

• Weak acids and bases: Partially dissociate, forming weak electrolytes.

Amphoteric and Ampholytic Species

Definition and Examples

A species is amphoteric or ampholytic if it can act as both an acid and a base. Water (H2O) is a classic example, as it can donate or accept a proton.

• Water: H2O/HO- (acid), H3O+/H2O (base)

• Hydrogen sulfide: H2S/HS- (acid), HS-/S2- (base)

Autoprotolysis of Water

Self-Ionization Reaction

Water can react with itself to form hydronium and hydroxide ions. This process is called autoprotolysis.

• Reaction: 2H2O(l) ⇌ H3O+(aq) + OH-(aq)

Lewis Theory of Acids and Bases

Electron Pair Concept

The Lewis theory defines acids and bases based on their ability to accept or donate electron pairs. A Lewis acid accepts an electron pair, while a Lewis base donates an electron pair.

• Lewis Acids: AlF3, BH3, AlCl3 (electron pair acceptors)

• Lewis Bases: Chloride ion, ammonia (electron pair donors)

Acidity and Basicity Constants

Acidity Constant (Ka)

The acidity constant (Ka) characterizes the strength of an acid in water. It is defined as:

• pKa:

• Strong acids: Large Ka, small pKa

Basicity Constant (Kb)

• pKb:

Solvent Constant (Kw)

• At 25°C:

Relationship Between Ka, Kb, and Kw

pKa Scale and Acid/Base Strength

pKa Values in Water

The pKa scale in water ranges from 0 to 14, allowing classification of acid-base pairs. The strongest acid in water is H3O+ (pKa = 0), and the strongest base is OH- (pKa = 14).

Definition of pH

pH Calculation

The pH of a solution is defined as the negative logarithm of the hydronium ion concentration:

• For neutral solutions: pH = 7

• For strong acids: pH approaches 0

• For strong bases: pH approaches 14

Acid-Base Reactions in Water

General Reaction

An acid-base reaction involves the exchange of a proton between two acid-base pairs. The reaction typically forms the weakest acid and base present.

• Example: HCl(aq) + NH3(aq) → NH4+(aq) + Cl-(aq)

• Example: NaOH(aq) + HNO3(aq) → NaNO3(aq) + H2O(l)

• Example: CH3COOH(aq) + H2O(l) ⇌ CH3COO-(aq) + H3O+(aq)

Summary Table: Strong vs. Weak Acids and Bases

Exercises and Applications

Sample Exercise

• Identify acids and bases according to Brønsted-Lowry definition and write their reactions with water.

• Justify the basic character of PO43- and (H3C)3N in water.

• Write the species present when a mole of NaOH, HClO4, or CH3COOH dissolves in water.

Conclusion

Acid-base chemistry is fundamental to understanding chemical reactions in aqueous solutions. The classification of acids and bases, their dissociation, and the calculation of pH are essential concepts in general chemistry.