Buffers and Buffer Action

Definition and Components of a Buffer

A buffer solution is a solution that resists significant changes in pH when small amounts of acid or base are added. Buffers are essential in many chemical and biological systems to maintain a stable pH environment.

• Components: A buffer typically consists of a weak acid and its conjugate base, or a weak base and its conjugate acid.

• Example: An acetic acid (CH3COOH) and sodium acetate (CH3COONa) mixture forms a buffer.

Action of a Buffer

Buffers function by neutralizing added acids (H+) or bases (OH-) through equilibrium reactions involving the weak acid/base pair.

• When a small amount of strong acid is added, the conjugate base reacts with the added H+ to form the weak acid, minimizing pH change.

• When a small amount of strong base is added, the weak acid donates a proton to neutralize the OH-, forming water and the conjugate base.

• Figure 18.3 (Action of a Buffer): Illustrates how the buffer components interact to resist pH changes. (Additional info: Figure typically shows the equilibrium shift upon acid/base addition.)

Calculating the pH of a Buffer Solution

The pH of a buffer can be calculated using either an ICE table (Initial, Change, Equilibrium) or the Henderson-Hasselbalch equation.

• ICE Table Method: Set up the equilibrium for the weak acid/base and solve for [H+].

• Henderson-Hasselbalch Equation:

• Where [A-] is the concentration of the conjugate base and [HA] is the concentration of the weak acid.

• Example: Calculate the pH of a buffer containing 0.10 M acetic acid and 0.10 M sodium acetate. Given :

Buffer Capacity and Preparation

Buffer capacity is the amount of acid or base a buffer can neutralize before the pH changes significantly. It depends on the concentrations of the buffer components.

• To prepare a buffer with a desired pH, select a weak acid with a pKa close to the target pH and adjust the ratio of acid to conjugate base.

Acid-Base Titrations

Overview of Acid-Base Titrations

Acid-base titrations are used to determine the concentration of an acid or base by reacting it with a titrant of known concentration. The progress of the titration is monitored by measuring pH as titrant is added.

• Key Features: Initial pH, buffer region, equivalence point, and post-equivalence region.

• Equivalence Point: The point at which stoichiometrically equivalent amounts of acid and base have reacted.

• pH = pKa Point: For weak acid-strong base titrations, this occurs when half the acid has been neutralized.

Types of Titrations and Their Curves

• Strong Acid–Strong Base: Sharp pH change at equivalence; initial pH is low, final pH is high.

• Weak Acid–Strong Base: Buffer region before equivalence; pH at equivalence is above 7.

• Polyprotic Acid: Multiple equivalence points, each corresponding to the loss of a proton.

Calculating pH During Titration

• Calculate pH at various stages: before titrant addition, within buffer region, at equivalence, and after equivalence.

• Use stoichiometry and equilibrium calculations as appropriate for each region.

Comparison of Titration Curves

Solubility and Solubility Product (Ksp)

Relating Solubility to Ksp

The solubility product constant (Ksp) quantifies the equilibrium between a solid ionic compound and its dissolved ions in a saturated solution.

• Molar Solubility: The number of moles of solute that dissolve per liter of solution.

• Calculating Ksp from Solubility: For a salt AB that dissociates as AB(s) ⇌ A+(aq) + B-(aq):

• Example: If the molar solubility of AgCl is s, then .

Common Ion Effect on Solubility

The presence of a common ion decreases the solubility of an ionic compound due to Le Châtelier's principle.

• Example: Adding NaCl to a saturated AgCl solution decreases AgCl solubility because Cl- is a common ion.

Effect of pH on Solubility

Solubility of some salts increases or decreases with pH, especially if the anion is basic and can react with H+.

• Example: The solubility of CaF2 increases in acidic solution because F- reacts with H+ to form HF.

Selective Precipitation and Complex Ion Formation

Selective Precipitation

Selective precipitation is used to separate ions in a mixture by adding a reagent that precipitates one ion while leaving others in solution.

• Compare the ion product (Q) to Ksp to predict precipitation:

• If Q > Ksp, precipitation occurs; if Q < Ksp, no precipitation.

• Calculate the minimum reagent concentration needed for selective precipitation.

• Determine concentrations of ions remaining after precipitation.

Complex Ion Formation and Solubility

Formation of complex ions can increase the solubility of certain salts by removing ions from the equilibrium.

• Example: AgCl dissolves in ammonia due to formation of [Ag(NH3)2]+ complex ion.

Summary Table: Key Concepts

Additional info: These notes synthesize the learning outcomes and practice recommendations from the provided material, expanding on buffer systems, titration analysis, solubility equilibria, and selective precipitation, as covered in a standard general chemistry curriculum.