CIE A Level Chemistry: Chemical Equilibria and Le Chatelier's Principle (Chapter 7.1)

Reversible reaction

Some reactions go to completion where the reactants are used up to form the products and the reaction stops when all of the reactants are used up. In reversible reactions the products can react to reform the original reactants.

Dynamic equilibrium

In a dynamic equilibrium the reactants and products are dynamic (they are constantly moving). The rate of the forward reaction is the same as the rate of the backward reaction in a closed system. The concentrations of the reactants and products are constant.

Closed system

A closed system is one in which none of the reactants or products escape from the reaction mixture.

Open system

In an open system, matter and energy can be lost to the surroundings.

Le Chatelier's Principle

Le Chatelier's principle says that if a change is made to a system at dynamic equilibrium, the position of the equilibrium moves to minimise this change.

Position of the equilibrium

The position of the equilibrium refers to the relative amounts of products and reactants in an equilibrium mixture.

Effects of concentration

When the concentration of reactants increases, the position of equilibrium shifts to the right. When the concentration of products increases, the position of equilibrium shifts to the left.

Increase in concentration of a reactant

Equilibrium shifts to the right.

Decrease in concentration of a reactant

Equilibrium shifts to the left.

Increase in concentration of a product

Equilibrium shifts to the left.

Decrease in concentration of a product

Equilibrium shifts to the right.

Effect of adding CH3COOC2H5

The position of the equilibrium moves to the left and more ethanoic acid and ethanol are formed.

Effect of removing C2H5OH

The position of the equilibrium moves to the left and more ethanoic acid and ethanol are formed.

Effect of adding water to equilibrium mixture

There is no effect as the water dilutes all the ions equally so there is no change in the ratio of reactants to products.

Changes in pressure

Changes in pressure only affect reactions where the reactants or products are gases.

Increase in pressure

Equilibrium shifts in the direction that produces a smaller number of molecules of gas to decrease the pressure again.

Decrease in pressure

Equilibrium shifts in the direction that produces a larger number of molecules of gas to increase the pressure again.

Equilibrium Constant (Kc)

Links concentrations of reactants and products at equilibrium.

Equilibrium Expression

Mathematical representation of Kc based on stoichiometry.

Endothermic Reaction

Absorbs heat, shifts right with increased temperature.

Exothermic Reaction

Releases heat, shifts left with increased temperature.

Catalyst

Increases reaction rate without affecting equilibrium position.

Partial Pressure

Pressure a gas would exert alone in a container.

Mole Fraction

Ratio of moles of a gas to total moles.

Total Pressure

Sum of partial pressures of all gases present.

Equilibrium Shift

Movement of equilibrium position in response to changes.

Gas Molecules

Affects pressure; fewer molecules decrease pressure.

Temperature Increase Effect

Shifts equilibrium in endothermic direction.

Temperature Decrease Effect

Shifts equilibrium in exothermic direction.

Equilibrium Concentration

Concentration of reactants/products when equilibrium is reached.

Kp

Equilibrium constant based on partial pressures of gases.

Solid in Equilibrium

Solids are ignored in equilibrium expressions.

ΔH

Change in enthalpy; indicates heat absorbed/released.

Reaction Rate

Speed at which reactants convert to products.

Stoichiometry

Relationship between quantities of reactants/products.

Equilibrium Position

State of a system where reactants/products are balanced.

Gas Reaction Example

2NO (g) ⇌ 2NO (g) + O (g) shows pressure effects.

AgCO Reaction

AgCO (s) ⇌ AgO (s) + CO (g) is endothermic.

Pressure Decrease Effect

Shifts equilibrium towards side with more gas molecules.

Equilibrium Constant Specificity

Kc only changes with temperature variations.

Equilibrium Expression for N2 + 3H2

Kc = [NH3]^2 / ([N2][H2]^3).

Equilibrium Expression for N2O4

Kc = [NO]^2 / [N2O4].

Equilibrium Expression for SO2

Kc = [SO3]^2 / ([SO2]^2[O2]).

Equilibrium Constant

A value that expresses the ratio of the concentrations of products to reactants at equilibrium.

Concentration

The amount of a substance per defined space, typically expressed in mol dm−3.

Kc

The equilibrium constant for reactions expressed in terms of concentration.

Volume

The amount of space that a substance occupies, measured in dm3.

Moles

A measure of the amount of substance, where 1 mole contains approximately 6.022 x 10^23 entities.

Significant Figures

Digits in a number that contribute to its precision.

ICE Table

A table used to track the initial, change, and equilibrium concentrations of reactants and products.

Balanced Chemical Equation

An equation that represents a chemical reaction with equal numbers of each type of atom on both sides.

Hydrolysis

A chemical reaction involving the breaking of a bond in a molecule using water.

Equilibrium Reaction

A reaction that can proceed in both the forward and reverse directions, reaching a state of balance.

Stoichiometric Equation

An equation that shows the proportion of reactants and products in a chemical reaction.

Units of K

The units of the equilibrium constant depend on the form of the equilibrium expression.

Equilibrium Constant Calculation

The process of determining the value of K by substituting equilibrium concentrations into the equilibrium expression.

Example Calculation

A worked example demonstrating how to calculate K using given concentrations and volumes.

Concentration Calculation Formula

Concentration (mol dm−3) = number of moles / volume (dm3).

Equilibrium Constant Expression for Reaction

Kc = [Products] / [Reactants], where concentrations are raised to the power of their coefficients.

Equilibrium Constant Value

A numerical value that indicates the extent to which a reaction proceeds to completion.

Example of K Calculation

Kc = 0.364 × 0.364 / (0.070 × 0.470) = 4.03.

Example of Kp Calculation

Kp = (8.0 x 10^6)^2 / (1.0 x 10^6)^2 × (7.0 x 10^6) = 9.1 x 10^-6.

Final Answer Significant Figures

The final answer should be given to the smallest number of significant figures used in the question.

Equilibrium Reaction Example

The equilibrium between hydrogen, iodine, and hydrogen bromide: H2(g) + I2(g) ⇌ 2HI(g).

Equilibrium

State where reactants and products are balanced.

Equilibrium Constant (K)

Ratio of product concentrations to reactant concentrations.

Le Chatelier's Principle

System adjusts to counteract changes in conditions.

Haber Process

Synthesis of ammonia from nitrogen and hydrogen.

Contact Process

Synthesis of sulfuric acid from sulfur dioxide and oxygen.

Compromise Pressure

Optimal pressure balancing yield and cost.

Compromise Temperature

Optimal temperature balancing yield and reaction rate.

Molar Ratio

Proportion of moles of reactants to products.

K Value Increase

Occurs with temperature rise in endothermic reactions.

K Value Decrease

Occurs with temperature rise in exothermic reactions.

Ammonia Yield Maximization

Achieved by increasing pressure and decreasing temperature.

Sulfuric Acid Yield Maximization

Achieved by increasing pressure and using optimal temperature.

Equilibrium Restoration

Achieved by adjusting concentrations or removing products.

Gas Molecule Count

Determines direction of equilibrium shift under pressure changes.

Iron Catalyst

Used in Haber process to speed up ammonia synthesis.

Vanadium(V) Oxide

Catalyst used in Contact process for sulfuric acid production.

Kinetic Energy

Energy of molecules affecting reaction rates.

Decomposition Reaction

Reaction where a compound breaks down into simpler substances.