6. PRECIPITATION TITRATIONS

Introduction

Precipitation titration is a type of volumetric analysis based on the formation of an insoluble solid (precipitate) during a chemical reaction. These titrations are used to determine the concentration of ions that form slightly soluble salts with the titrant.

A typical example involves the reaction of silver nitrate with halides (Cl⁻, Br⁻, I⁻), forming insoluble silver halides. Such titrations are called argentometric titrations.

Precipitation Reaction

When two aqueous ionic solutions are mixed, an insoluble compound may form and separate as a solid. The general requirements for a reaction to be useful in titrimetric analysis include:

  • The precipitate must be practically insoluble.
  • The reaction must be rapid and quantitative.
  • Adsorption or co-precipitation should not interfere.
  • The equivalence point should be clearly detectable.

Example: AgNO₃ + NaCl → AgCl↓ + NaNO₃

Here, silver chloride forms as a white precipitate, driving the reaction to completion.

Solubility Product (Ksp)

The solubility product (Ksp) expresses the equilibrium between a solid salt and its ions in a saturated solution. For a slightly soluble salt BA:

BA(s) ⇌ B⁺ + A⁻

At equilibrium, Ksp = [B⁺][A⁻]

If the product of ion concentrations exceeds Ksp, precipitation occurs. If it is lower, the solid dissolves until equilibrium is reached.

Factors Affecting Solubility

1. Effect of Acids

Solubility of salts increases in acidic solutions when the anion is the conjugate base of a weak acid. For example, carbonates, phosphates, and sulfides are more soluble in acidic conditions.

2. Effect of Temperature

Most inorganic salts are more soluble at higher temperatures. Precipitation in hot solutions aids purity by dissolving impurities and speeding filtration. However, salts like magnesium ammonium phosphate require cooling before filtration to avoid loss.

3. Effect of Solvent

Inorganic salts are generally more soluble in water than in organic solvents. Adding alcohols such as ethanol or methanol decreases solubility, aiding precipitation.

Examples:

  • Addition of ethanol reduces solubility of PbSO₄ for quantitative separation.
  • Calcium nitrate dissolves in alcohol while strontium nitrate precipitates.
  • Potassium can be separated from sodium using potassium hexachloroplatinate (K₂PtCl₆) in alcohol-water mixtures.

Argentometric Titrations

Argentometric titrations involve the use of silver nitrate (AgNO₃) as a titrant to precipitate halide and pseudohalide ions. The resulting precipitates include AgCl, AgBr, AgI, AgSCN, etc.

Requirements:

  • Precipitation must be stoichiometric and quantitative.
  • Reaction equilibrium must be rapid and well-defined.
  • Precipitate must be sparingly soluble.
  • Endpoint detection must be precise and visible.

Titration Curve (NaCl vs AgNO₃)

The typical titration curve for argentometric titration is sigmoid-shaped when plotting pAg (–log[Ag⁺]) against the volume of titrant added. It can be divided into three regions:

  • Pre-equivalence region: Excess Cl⁻, Ag⁺ concentration is low.
  • Equivalence point: Stoichiometric completion, sharp pAg change.
  • Post-equivalence region: Excess Ag⁺, pAg rises sharply.

Sharper endpoints are obtained with higher reagent concentrations.

Methods of Endpoint Detection

Several classical methods are used to identify endpoints in argentometric titrations:

  • Gay Lussac’s Method
  • Mohr’s Method
  • Volhard’s Method
  • Fajan’s Method

1. Gay Lussac’s Method

This is an indicator-free method that detects the endpoint visually by disappearance of turbidity. When no new precipitate forms upon adding titrant, the endpoint is reached.

Procedure:

  1. Dissolve 0.4 g of AgNO₃ in 100 ml water with a few drops of HNO₃ and a crystal of Ba(NO₃)₂.
  2. Titrate with standard 0.1 M NaCl solution while shaking after each addition.
  3. When no turbidity appears after a drop of NaCl, the endpoint is reached.

Advantage: Simple and does not require indicators.

2. Mohr’s Method

Mohr’s method determines chlorides and bromides using potassium chromate (K₂CrO₄) as an indicator in a neutral medium.

Reaction:

At equivalence, Ag⁺ ions react with chromate to form red-brown Ag₂CrO₄ precipitate.

Indicator Reaction:

2Ag⁺ + CrO₄²⁻ → Ag₂CrO₄ ↓ (reddish brown)

Restrictions:

  • Not applicable in basic solution (forms AgOH).
  • Cannot be used in presence of ammonia or reducing agents.
  • Interference from anions like PO₄³⁻ or S²⁻ must be avoided.

3. Volhard’s Method

In Volhard’s method, a back titration is performed using standard AgNO₃ followed by titration of the unreacted Ag⁺ with thiocyanate (SCN⁻) using ferric alum as indicator (red complex formation at endpoint).

Indicator Reaction:

Fe³⁺ + SCN⁻ → [Fe(SCN)]²⁺ (red color at endpoint)

Modified Volhard’s Method:

To prevent AgCl dissolution, chloroform or other wetting agents are added after excess AgNO₃ addition.

4. Fajan’s Method

Developed by K. Fajan, this method uses adsorption indicators that are absorbed on the surface of the precipitate, changing color near the equivalence point.

Mechanism:

  • Initially, AgCl particles adsorb Cl⁻ ions forming a negatively charged surface with Na⁺ as counter-ions.
  • After equivalence, excess Ag⁺ reverses the charge, and dye anions (like fluorescein) are adsorbed instead of NO₃⁻.
  • This adsorption produces a visible color change.

Indicators:

  • Fluorescein (for Cl⁻)
  • Eosin (for Br⁻, I⁻)

pH Range:

  • 6.5–10.3 for chloride titration
  • 2.0–10.3 for bromide and iodide titration

Indicator Mechanism: The adsorption of dye on AgCl surface at endpoint changes the color from yellow-green to pink/red, indicating completion.

Thiocyanatometric and Mercurometric Titrations

Thiocyanatometric titrations use AgNO₃ or Fe³⁺ ions to titrate thiocyanate-containing compounds, while mercurimetric titrations use Hg²⁺ ions forming insoluble mercury salts with halides or thiocyanates.

Detailed Notes:

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