Chemical Equilibrium

• Chemical equilibrium describes a system where the conversion of reactants to products and products to reactants occurs at equal rates
• There is no net change in concentrations of reactants or products once equilibrium is established

Understanding Reaction Rates and Equilibrium

• Initially, the forward reaction rate is high due to high reactant concentrations
• The reverse reaction rate starts at zero if only reactants are present
• Over time, reactant concentrations decrease, reducing the forward rate
• Simultaneously, product concentrations increase, raising the reverse rate
• Equilibrium is reached when forward and reverse reaction rates equalize

Dynamic Equilibrium Characteristics

• Dynamic equilibrium seems static at a macroscopic level, with constant concentrations, pressure, volume, and color
• At the molecular level, reactions continue in both directions
• Reactants convert to products as products convert back to reactants
• Dynamic equilibrium maintains equal forward and reverse reaction rates, preventing net change

Law of Chemical Equilibrium

• Reactions in closed systems at constant temperatures spontaneously move to dynamic equilibrium
• The reaction quotient (Q) measures amounts of reactants and products at any given time
• At equilibrium, Q equals the equilibrium constant (K)
• K is constant for a specific reaction at a specific temperature, independent of initial concentrations

Thermodynamic Reaction Quotient (Q)

• Q is based on the activities (effective concentrations) of reactants and products
• Activity is relative to a standard state and is unitless
• For most uses, activity equals pressure in bars (for gases), 1 (for pure liquids and solids), or molar concentration (for solutions)
• General form: Q = (activities of products) / (activities of reactants), with each activity raised to the power of its coefficient in the balanced equation

Activities of Substances

• Gases: Activity, a = (γP) / P°, where P° =1 bar and γ is the activity coefficient (assumed to be 1 for ideal behavior)
• Pure Liquids/Solids: Activity, a ≈ 1 because they are nearly incompressible with small molar volumes
• Aqueous Solutions: Activity, a = γ[X]/C°, where C° = 1 mol/L and γ is close to 1 in diluted solutions

Key Idea Summary

• Activities are relative to a standard state and presented as unitless quantities
• For accurate activity values for high pressure or high concentration situation where inter molecular forces dominate
• For standard applications, activity equals numerical pressure (in bars) for gases, one for liquids/solids, or concentration (M) aqueous solutions

Relating K and Reaction Direction

• If Q < K, the reaction goes forward to reach equilibrium, increasing product production until Q = K
• If Q > K, the reaction goes in reverse, converting products back to reactants until Q = K
• If Q = K, the reaction is at equilibrium, with no net change

Gibbs Energy and Equilibrium

• Gibbs Energy (ΔG) indicates reaction spontaneity
• At equilibrium, ΔG = 0 (minimum Gibbs energy)
• The relationship ΔG = ΔG° + RTlnQ links Gibbs energy, standard Gibbs energy, and Q
• At equilibrium, ΔG = 0 so ΔG° = -RTlnK, where and K = Qeq

Calculating Equilibrium Constants

• Equilibrium constants (K) can be derived from standard Gibbs free energy (ΔG°) values
• Equilibrium constant values can be calculated from standard thermodynamic data
• The formula lnK = -ΔG°/RT is used to obtain K from tabulated thermodynamic data