Buffer Solutions

Definition and Importance

Buffer solutions are mixtures that resist changes in pH when small amounts of acid or base are added. They are essential in analytical chemistry for maintaining stable pH conditions during reactions and measurements.

• Buffer: A solution containing a weak acid and its conjugate base (or a weak base and its conjugate acid).

• Application: Used in titrations, biological systems, and chemical analysis.

Identifying Buffer Solutions

To determine if a solution is a buffer, check if it contains a weak acid/base and its salt.

• Example: 0.10 M NaCN forms a buffer (contains weak acid HCN and its salt NaCN).

• Example: 0.10 M NaCl does not form a buffer (contains only a strong electrolyte).

Calculating pH of a Buffer: Henderson-Hasselbalch Equation

The pH of a buffer can be calculated using the Henderson-Hasselbalch equation:

Example Calculation: For a buffer made by mixing 13.02 g of sodium acetate (NaC2H3O2) and 15.40 g of acetic acid (HC2H3O2) in 500 mL water:

• Moles of NaC2H3O2: mol

• Moles of HC2H3O2: mol

• Concentrations: M, M

Selecting a Buffer System

Choose a buffer whose is close to the desired pH.

• Example: For pH 5.02, use acetic acid ().

Buffer Capacity and Preparation

Buffer capacity depends on the concentrations of acid and base. The closer the ratio is to 1, the greater the buffer capacity.

Solubility and Precipitation Equilibria

Solubility Product Constant (Ksp)

The solubility product constant () describes the equilibrium between a solid salt and its ions in solution.

• General form: For :

• Example:

Calculating Ksp from Solubility

Given the solubility (S) of a salt, calculate :

• Example 1: AgCl solubility at 25°C is 0.00019 g/100 mL.

• Convert to molarity: M

• Example 2: Ag3PO4 solubility at 25°C is 0.20 mg/100 mL.

• Convert to molarity: M

Calculating Solubility from Ksp

• For salts of type MX:

• For salts of type MX2:

• For salts of type MX3:

Predicting Precipitation

To predict if a precipitate will form, compare the ion product (PI) to :

• PI < Ksp: No precipitate (unsaturated solution)

• PI = Ksp: Saturated solution (at equilibrium)

• PI > Ksp: Precipitate forms

Factors Affecting Solubility

• Common Ion Effect: Addition of a common ion decreases solubility.

• Complex Ion Formation: Increases solubility by removing ions from equilibrium.

• pH Effect: Addition of acid can increase solubility of salts containing basic anions.

• Temperature: Solubility generally increases with temperature for most salts.

Solubility in the Presence of Acids

For salts containing basic anions, solubility increases in acidic solutions due to reaction with H+:

• Example: AgCN in HNO3 solution:

• Combined:

Common Ion Effect Example

• AgCl:

• Adding Cl- (e.g., from NaCl) reduces AgCl solubility.

• New solubility:

Non-Common Ion Effect

Solubility of salts can increase in the presence of non-common ions due to ionic strength effects (ion-diverse effect).

Tables

Acid Strength Comparison Table

Buffer Selection Table

Solubility Product Table

Effect of Ionic Strength on Solubility

Summary

• Buffer solutions are crucial for maintaining pH stability in analytical procedures.

• Solubility equilibria and calculations allow prediction and control of precipitation reactions.

• Factors such as common ion effect, pH, and ionic strength influence solubility and must be considered in analytical chemistry.

Additional info: Some explanations and table entries were expanded for clarity and completeness based on standard analytical chemistry knowledge.