12_Unit_3_Equilibrium_ppt_2

Chemical Equilibrium

Reversible Reactions

• Reactions that do not go to completion.

• Mixture contains significant amounts of both products and reactants.

• Known as reversible reactions; the resulting mixture is in a state of equilibrium.

• The symbol (↔) signifies equilibrium between reactants and products.

Equilibrium Definitions

• Chemical Equilibrium: Forward and reverse reactions occur at the same rate.

• Reaction Rate: Number of atoms, ions, or molecules reacting per unit time to form products.

• Endergonic (Endothermic) Reaction: Non-spontaneous reaction.

• Exergonic (Exothermic) Reaction: Spontaneous reaction that releases free energy.

Additional Equilibrium Definitions

• Entropy (S): Measure of disorder; more disorder → greater reaction possibility.

• Enthalpy (H): Total energy stored in a substance; change in H (ΔH) is known as heat of reaction.

• Activation Energy: Minimum energy needed for colliding particles to react.

Catalyst & Reaction Dynamics

• Catalyst: Increases reaction rate without being consumed; quantity remains constant after.

• LeChatelier’s Principle: System at equilibrium will shift to relieve applied stress and re-establish equilibrium.

• Collision Theory: Particles form bonds upon collision, essential for reaction if they have sufficient kinetic energy.

• Law of Disorder: Matter moves towards maximum disorder; e.g., gases mixing in the atmosphere.

Reversible Reactions and Equilibrium

• Consider reaction: 2 NO(g, brown) ⇌ N2O4(g, colorless).

• In a sealed container, brown nitrogen dioxide gas will reach equilibrium over time.

Equilibrium Rules

1. Chemical equilibrium is constant at the macroscopic level but dynamic at the molecular level.

2. Opposing reaction rates are equal at equilibrium.

3. Equilibrium position is influenced by concentration, temperature, and pressure.

Heterogeneous Equilibrium

• Applies to reactions with reactants and products in different phases.

• Example: CaCO3 (s) ⇌ CaO (s) + CO2 (g)

• Kc = [CaO][CO2] / [CaCO3]

• Concentrations of solids/liquids not included in Kc expression.

Writing Equilibrium Constant Expressions

• Reacting species in condensed phase: expressed in molarity (M); for gases: M or atm.

• Pure solids and liquids excluded from Kc expressions.

• Kc is dimensionless; must specify balanced equation and temperature when quoted.

Example Equilibrium Constant Expressions

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

• Kc = [CH3COO-][H3O+] / [CH3COOH] (H2O is constant).

Le Châtelier’s Principle Applications

• System at equilibrium is subject to changes: concentration, volume, pressure, temperature, and catalysts will affect the equilibrium position.

• For concentration change, adjusting amounts leads to product or reactant shifts to restore equilibrium.

Stress Responses According to Le Châtelier’s Principle

• Increasing concentration of products shifts left.

• Decreasing concentration of products shifts right.

• Decreasing concentration of reactants shifts left; increasing shifts right.

Effects of Pressure and Volume

• For gases, increasing pressure shifts to the side with fewer molecules.

• Decreasing pressure shifts to the side with more molecules.

• Volume changes also affect the direction of equilibrium shifts.

Temperature Effects on Equilibrium

• Only temperature affects the equilibrium constant (K).

• For exothermic reactions:

  • Increase temperature → K decreases.

  • Decrease temperature → K increases.

• For endothermic reactions:

  • Increase temperature → K increases.

  • Decrease temperature → K decreases.

Adding or Removing Chemicals

• Adding reactants or removing products shifts towards products.

• Removing reactants or adding products shifts towards reactants.

Influence of Catalysts on Equilibrium

• Addition of a catalyst does not change K or shift equilibrium position but allows the system to reach equilibrium sooner.

Equilibrium Constant Fundamentals

• Used to indicate the equilibrium position; derived from the concentrations of reactants/products at equilibrium.

• Generic expression: K = [C]^{c}[D]^{d} / [A]^{a}[B]^{b}

• Concentrations used for calculating the equilibrium constant must be specified with a balanced equation and its temperature.

Sample Calculation of Equilibrium Constants

• Reaction: CO + Cl2 ⇌ COCl2

• Calculating Kc with given concentrations:

  • [CO] = 0.012 M, [Cl2] = 0.054 M, [COCl2] = 0.14 M

  • Kc = [COCl2] / ([CO][Cl2]) = 0.14 / (0.012 × 0.054) = 220.

Predicting Reaction Direction from Non-Equilibrium Mixtures

• Calculate Reaction Quotient (Q) using initial concentrations in the Kc expression.

• Comparisons:

  • If Q > K: shift left (favoring reactants).

  • If Q = K: system is at equilibrium.

  • If Q < K: shift right (favoring products).

Calculating Equilibrium Concentrations

1. Balance chemical equation.

2. Determine Q to predict shift direction.

3. Define changes needed to reach equilibrium with 'x'.

4. Write equilibrium constant expression and substitute defined expressions.

5. Solve for x and find final concentrations.