Equilibrium and Le Chatelier's Principle: General Chemistry Study Notes

Equilibrium and Le Chatelier's Principle

Introduction to Chemical Equilibrium

Chemical equilibrium occurs when the rates of the forward and reverse reactions are equal, resulting in constant concentrations of reactants and products. The position of equilibrium can be affected by changes in temperature, pressure, concentration, and the addition of catalysts.

• Dynamic Equilibrium: Both forward and reverse reactions continue to occur at equal rates.

• Equilibrium Constant (Kc or Kp): Quantifies the ratio of product to reactant concentrations at equilibrium.

• Le Chatelier's Principle: If a system at equilibrium is disturbed, it will shift to counteract the disturbance and restore equilibrium.

Le Chatelier's Principle: Effects of Changes on Equilibrium

Consider the reaction: CH4(g) + 2 H2S(g) → CS2(g) + 4 H2(g) (endothermic)

• Increase Temperature: Shifts equilibrium toward products (right) for endothermic reactions.

• Add H2S Gas: Shifts equilibrium toward products (right) due to increased reactant concentration.

• Remove CS2: Shifts equilibrium toward products (right) to replace removed product.

• Increase Pressure: Shifts equilibrium toward the side with fewer moles of gas. Count moles on each side to determine direction.

• Add a Catalyst: No effect on equilibrium position; only increases rate at which equilibrium is reached.

Example: If temperature is increased for an endothermic reaction, the system absorbs heat and shifts toward products.

Calculating Equilibrium Concentrations

For the equilibrium: NH4HS(s) ↔ NH3(g) + H2S(g) at 24°C, Kc = 1.58 × 10-2 mol/L. Given initial amounts of NH3 and H2S, and excess solid NH4HS, the direction of reaction can be predicted by comparing the reaction quotient (Q) to Kc.

• Reaction Quotient (Q): $Q = [NH_3][H_2S]$

• If Q < Kc, reaction shifts right (toward products).

• If Q > Kc, reaction shifts left (toward reactants).

Example: Calculate Q using initial concentrations and compare to Kc to predict the shift.

Equilibrium Constant Calculations

For the reaction: CaCO3(s) ↔ CaO(s) + CO2(g), only CO2 gas appears in the equilibrium expression because solids are omitted.

• Equilibrium Expression: $K_p = P_{CO_2}$

• Given $P_{CO_2} = 0.220$ atm at 800 K, $K_p = 0.220$ atm.

Example: If the equilibrium pressure of CO2 is measured, it equals Kp for this reaction.

Stoichiometry and Equilibrium: Mass Calculations

For the equilibrium: 3Fe(s) + 4H2O(g) ↔ Fe3O4(s) + 4H2(g), given Kc and initial amounts, equilibrium concentrations and masses can be calculated using ICE tables (Initial, Change, Equilibrium).

• Set up ICE Table: Track changes in concentrations as reaction proceeds to equilibrium.

• Equilibrium Constant Expression: $K_c = \frac{[H_2]^4}{[H_2O]^4}$ (solids omitted)

• Solve for equilibrium concentrations and convert to mass using molar mass.

Example: Calculate the mass of Fe3O4 formed at equilibrium.

Solubility Equilibria and Common Ion Effect

Consider the equilibrium: FeS(s) ↔ Fe2+(aq) + S2-(aq). The solubility of FeS is affected by the addition of various ions due to the common ion effect and precipitation reactions.

• Common Ion Effect: Adding an ion already present in equilibrium shifts the reaction to reduce solubility.

• Precipitation: Adding ions that form insoluble compounds with Fe2+ or S2- can decrease their concentrations.

Stress	Shift	[Fe2+]	[S2-]
Add water	no shift	no change	no change
Add HCl(aq)	left	decrease	decrease
Add Na2S(aq)	left	decrease	decrease
Add FeS(s)	no shift	no change	no change
Add AgNO3(aq)	right	increase	decrease
Add FeSO4(aq)	left	decrease	decrease

Additional info: Table entries inferred based on standard solubility and common ion effect principles.

Acid-Base Equilibria

Equilibria involving weak acids and bases, such as HS- ↔ H+ + S2-, are governed by their dissociation constants (Ka or Kb).

• Equilibrium Expression: $K_a = \frac{[H^+][S^{2-}]}{[HS^-]}$

• Shifts in equilibrium occur with changes in concentration, addition of strong acids/bases, or removal of products.

Determining Equilibrium Constants Experimentally

Equilibrium constants can be determined by measuring concentrations of reactants and products at equilibrium. For example, the decomposition of hydrogen iodide:

• Reaction: 2HI(g) ↔ H2(g) + I2(g)

• Equilibrium concentrations are measured by titration or other analytical methods.

• Equilibrium Constant Expression: $K_c = \frac{[H_2][I_2]}{[HI]^2}$

Example: Use titration data to calculate equilibrium concentrations and determine Kc.

Summary Table: Key Equilibrium Concepts

Concept	Description	Example Equation
Le Chatelier's Principle	System shifts to counteract changes	$N_2(g) + 3H_2(g) \rightleftharpoons 2NH_3(g)$
Equilibrium Constant (K)	Ratio of product to reactant concentrations	$K_c = \frac{[NH_3]^2}{[N_2][H_2]^3}$
Common Ion Effect	Solubility decreases with added common ion	$AgCl(s) \rightleftharpoons Ag^+(aq) + Cl^-(aq)$

Additional info: Some explanations and table entries expanded for clarity and completeness.