University:
York University
Course:
CHEM 1001 | Chemical Dynamics
Academic year:

2022
Views:

237

Pages:

3
Author:

Niklas

Chapter 10: Equilibrium Exothermic vs. Exoergic Exothermic reactions have negative ΔH and can deliver heat. Exoergic reactions have negative ΔG and can deliver useful work. Standard Free Energy of Formation Change in free energy for the reaction in which 1 mole of a compound in its standard state is made from the elements in their standard states. ΔG° values are strongly temperature-dependent, so we are careful to include T=298.15 K in the definition of ΔG°. ΔG° = ΣΔG°f(products) - ΣΔG°f(reactants) ΔG° for elements in elemental form is 0. Pressure ΔG = ΔH - TΔS ΔG = ΔG° + RT ln(Q) Reactant Quotient (Q) Q = [products] / [reactants] When Q < K, the reaction shifts to the right (ΔG < 0). When Q > K, the reaction shifts to the left (ΔG > 0). When Q = K, the reaction is in equilibrium (ΔG = 0). Equilibrium Constant (K) K = exp(-ΔG°/RT) K > 1 favors products at equilibrium (ΔG° < 0). K < 1 favors reactants at equilibrium (ΔG° > 0). Chapter 10: Equilibrium Activities Gases: Partial pressure / 1 atm Solutes: Molarity / 1 M Pure liquids: 1.0 Pure solids: 1.0 Only include gases in Qatk calculations Le Châtelier's Principle If you apply pressure to the reaction, the reaction will shift to relieve the pressure. If more heat is added, the reaction shifts to the right. ICE Table A table used to organize the initial concentrations, changes, and equilibrium concentrations of reactants and products. Example: N₂ + 3H₂ ↔ 2NH₃ Equilibrium Constant (K) K = [products] / [reactants] Temperature dependence of K: ln(K) = -ΔH°/R * (1/T - 1/T°) Diagram: A graph showing the relationship between ln(K) and 1/T. The slope of the line is equal to -ΔH°/R. Chapter 10: Equilibrium / Chapter 6 van't Hoff Equation ln(K₂) - ln(K₁) = ΔH°/R * (1/T₁ - 1/T₂) Assumes ΔH° and ΔS° don't change with temperature. Works best for small temperature changes. Units for K and Q Always cancel. Reference States Gases: 1 atm Solutes: 1 M Pure solids or liquids: 1 Ideal Gases and Ideal Solutions Write K and Q in terms of pressure or concentration with no units. For pure solids or liquids, replace concentration with 1. When you don't know pressure, use concentrations for gas phase. Equilibrium Constant (K) K = exp(-ΔG°/RT) For heterogeneous equilibria: If ALL species are gases, then ΔG = RT ln(K) and K = exp(-ΔG°/RT). If ALL species are in solution, then ΔG = RT ln(K) and K = exp(-ΔG°/RT). If mixed gases and solutions, then ΔG = RT ln() and K = exp(-ΔG°/RT). Le Châtelier's Principle Decrease temperature: Shifts to exothermic side. Increase temperature: Shifts to endothermic side. Remove SO₃ at constant temperature: Shifts to right. Increase volume at constant temperature: Shifts to side with more moles of gas. Add Ar(g) at constant temperature: No effect. Add a catalyst at constant temperature: No effect.