3. ACID BASE TITRATIONS

Acid Base Titrations – Introduction and Concepts

Acid-Base Titration is also called a neutralization titration. It is based on the chemical reaction between an acid and a base to form salt and water. This type of titration helps to determine the unknown concentration of an acid or a base.

Acidimetry and Alkalimetry

  • Acidimetry: The process in which a standard acid solution is used to find out the strength of a base or basic salt (like Na2CO3 or Na2B4O7).
  • Alkalimetry: The process in which a standard base or alkali solution is used to determine the concentration of an acid or acidic salt.

Both of these titrations depend on the neutralization reaction between acid and base.

Advantages of Acid-Base Titrations

  • The reaction between acid and base is very fast and almost instantaneous.
  • The reaction is simple and occurs without side reactions.
  • The reaction goes to completion (it is stoichiometric and definite).
  • Results are accurate and easy to reproduce.

Theories of Acids and Bases

There are several theories that explain how acids and bases behave. The most commonly discussed ones are:

  1. Arrhenius Theory
  2. Bronsted–Lowry Theory
  3. Lewis Theory
  4. Usanovich Theory
  5. Lux-Flood Theory

1. Arrhenius Theory (H+ and OH– Concept)

According to Arrhenius, acids are substances that produce hydrogen ions (H+) in water, while bases produce hydroxide ions (OH–).

Example:

HCl → H+ + Cl–
NaOH → Na+ + OH–

When an acid reacts with a base, they neutralize each other to form salt and water.

Example: HCl + NaOH → NaCl + H2O

Limitations:

  • Applies only to aqueous solutions.
  • Cannot explain acidity or basicity in non-water solvents.
  • Fails to define acids and bases without H+ or OH– ions (e.g., SO2, CO2).

2. Bronsted–Lowry Theory (Proton Donor and Acceptor)

Proposed in 1923, this theory states:

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

Example:

NH3 + H2O ⇌ NH4+ + OH–

Here, ammonia acts as a base (accepts H+), and water acts as an acid (donates H+).

Advantages:

  • Explains acid-base reactions in both aqueous and non-aqueous solvents.
  • Introduces the concept of conjugate acid-base pairs.

Limitation: Cannot explain acids and bases that do not involve proton transfer.

3. Lewis Theory (Electron Donor and Acceptor)

G.N. Lewis (1923) proposed that:

  • Acids are electron pair acceptors.
  • Bases are electron pair donors.

Example:

BF3 + NH3 → F3B←NH3

Here, BF3 (acid) accepts a lone pair from NH3 (base).

This theory explains acid-base reactions without the need for protons and applies to both aqueous and non-aqueous systems.

4. Usanovich Theory (Cation and Anion Donor/Acceptor)

According to Usanovich (1934):

  • Acid: A species that gives cations or accepts anions/electrons.
  • Base: A species that gives anions or accepts cations/electrons.

This theory includes oxidation-reduction reactions as a part of acid-base behavior.

5. Lux-Flood Theory (Oxide-Ion Donor/Acceptor)

This theory, proposed by Lux (1929) and Flood (1947), is based on oxide ions.

  • Acids: Species that accept oxide ions (O2–).
  • Bases: Species that donate oxide ions.

Example: SiO2 (acid) + CaO (base) → CaSiO3

Law of Mass Action

According to Goldberg and Wage (1867), the rate of a chemical reaction depends on the active masses (concentrations) of the reactants. This concept is the foundation of the equilibrium constant (K) used in titration calculations.

Role of Solvents in Titration

The solvent plays a key role in acid-base titration by influencing ionization and reaction rates.

  • Neutral Solvents: Acetonitrile, alcohols, benzene.
  • Acidic Solvents: Acetic acid, formic acid, propionic acid.

Water is the most common solvent in aqueous titrations because it can act as both an acid and a base (amphiprotic).

Buffer Solutions

Buffer Solutions resist changes in pH when small amounts of acid or base are added.

  • Acidic Buffer: Formed from a weak acid and its salt (e.g., acetic acid + sodium acetate).
  • Basic Buffer: Formed from a weak base and its salt (e.g., ammonia + ammonium chloride).

Example: Blood acts as a natural buffer maintaining pH between 7.35 and 7.45 using the bicarbonate system.

Henderson–Hasselbalch Equation

The Henderson–Hasselbalch equation is used to calculate the pH of buffer solutions:

pH = pKa + log([A–] / [HA])

  • If pH = pKa → acid and salt concentrations are equal.
  • If pH < pKa → solution is more acidic.
  • If pH > pKa → solution is more basic.

Classification of Acid–Base Titrations

  1. Strong Acid vs Strong Base (e.g., HCl vs NaOH)
  2. Weak Acid vs Strong Base (e.g., CH3COOH vs NaOH)
  3. Strong Acid vs Weak Base (e.g., HCl vs NH4OH)
  4. Weak Acid vs Weak Base (e.g., CH3COOH vs NH4OH)

Each type produces a different titration curve and requires a suitable indicator based on the pH range at the equivalence point.

Theories of Indicators

Indicators are substances that change color near the end point of a titration. They are usually weak acids or bases that exist in two colored forms depending on pH.

1. Ostwald Theory

Explains color change based on ionization of the indicator:

HIn ⇌ H+ + In–

The undissociated (HIn) and dissociated (In–) forms have different colors. The observed color depends on their ratio in solution.

2. Resonance Theory

According to this theory, color change occurs due to structural or electronic rearrangements in the indicator molecules between acidic and basic forms. The change in resonance structure alters light absorption and color.

Indicators for Acid–Base Titrations

  • Strong Acid vs Strong Base: Methyl Orange, Phenolphthalein, Bromothymol Blue
  • Weak Acid vs Strong Base: Phenolphthalein, Thymol Blue
  • Strong Acid vs Weak Base: Methyl Orange, Bromocresol Green
  • Weak Acid vs Weak Base: Mixed Indicators

Mixed Indicators

When both the acid and base are weak, the pH range near the end point is narrow. In such cases, a combination of two indicators (mixed indicator) gives a sharper color change.

Examples:

  • Neutral Red + Methylene Blue → Violet Blue to Green (around pH 7)
  • Methyl Green + Phenolphthalein → Gray to Pale Blue (pH 8.4–8.8)
  • Thymol Blue + Cresol Red → Yellow to Violet (pH 8.3)

Detailed Notes:

For PDF style full-color notes, open the complete study material below:

Share your love