Dynamic Equilibrium

Dynamic equilibrium

Dynamic equilibrium is where the rate of the forward reaction equals the rate of the reverse reaction, and the concentrations of reactants and products remain constant.

Closed system

A closed system is a system where none of the reactants or products can enter or leave the reaction vessel.

Open system

An open system is a system where the reactants or products can both enter or leave the reaction vessel.

Equilibrium in closed system

Equilibrium can only be reached in a closed system.

Concentration of reactants

In a reversible reaction, the concentration of the reactants is highest at the start of the reaction.

Rate of reverse reaction

The rate of the reverse reaction will increase until the system reaches dynamic equilibrium.

Reversible reaction

A reversible reaction is a reaction that can occur in both directions: the forward reaction (which forms the products) and the reverse direction (which forms the reactants).

Representation of reversible reactions

Reversible reactions are represented using two opposing arrows, ⇌.

Energy change of reverse reaction

If the forward reaction of a reversible reaction is exothermic, the reverse reaction is endothermic.

Macroscopic properties in dynamic equilibrium

In dynamic equilibrium, there is no change in macroscopic properties such as color and density as they depend on concentration.

Concentrations in dynamic equilibrium

In dynamic equilibrium, the concentrations of reactants and products are constant, not necessarily equal.

Equilibrium in open flasks

Equilibrium can be reached in open flasks for reactions taking place entirely in solution.

Approaching dynamic equilibrium

Dynamic equilibrium can be approached starting with either reactants or products.

Key difference between closed and open systems

For reactions involving gases, equilibrium can only be reached in a closed system, not in an open system.

Equilibrium constant expression

The equilibrium constant expression is an equation that links the equilibrium constant, K, to the concentrations of reactants and products at equilibrium, taking into account the stoichiometry of the equation.

General form of equilibrium constant expression

The general form of the equilibrium constant expression is: K = [C]c [D]d / [A]a [B]b.

Solids in equilibrium constant expressions

Solids are ignored in equilibrium constant expressions.

Square brackets in equilibrium constant expression

The square brackets represent the equilibrium concentrations of the substances in mol dm-3.

Specificity of equilibrium constant

The equilibrium constant, K, is specific to a given equation.

Reversing reaction equation

When a reaction equation is reversed, the equilibrium constant becomes the reciprocal of the original K value.

Initial concentrations in equilibrium constant expression

The equilibrium constant expression uses equilibrium concentrations of reactants and products.

Equilibrium constant expression

Typically expressed in mol dm-3.

Equilibrium constant expression inclusion

Excludes solids and pure liquids.

Size of the equilibrium constant, K

Indicates how the equilibrium mixture is made up with respect to reactants and products.

K > 1

The equilibrium lies to the right hand side (products are favored).

K >> 1

The equilibrium lies far over to the right hand side and the reaction almost goes to completion.

K < 1

The concentration of reactants is greater than the concentration of products.

K = 1

There are significant amounts of both reactants and products, and equilibrium does not lie in favor of either.

K temperature dependence

K is constant at a specified temperature but can change with temperature.

Stronger acids and K

Stronger acids always have a higher value of K than weaker acids.

Relationship between K and K'

K' = 1/K or K' = K-1, where K' is the equilibrium constant for the reverse reaction.

Effect of concentration changes on K

The equilibrium constant, K, is not affected by changes in concentration of reactants or products at a given temperature.

Le Chatelier's principle

States that if a change is made to a system at dynamic equilibrium, the position of the equilibrium moves to minimise this change.

Increasing reactant concentration

Shifts the equilibrium to the right.

Decreasing product concentration

Shifts the equilibrium to the right to reduce the effect of the decrease.

Increase in pressure effect

Shifts the equilibrium in the direction that produces the smaller number of gas molecules.

Changes in pressure effect

Only affect reactions where the reactants or products are gases, and the number of moles of gaseous reactant are different to the number of moles of gaseous product.

Increase in temperature effect on endothermic reaction

Shifts the equilibrium in the endothermic direction, favoring the products.

Catalysts effect on equilibrium

Catalysts have no effect on the position of equilibrium or the value of K.

Catalysts effect on reaction speed

Catalysts only cause a reaction to reach its equilibrium faster, without affecting the position of equilibrium.

Le Chatelier's principle in heterogeneous equilibria

Can be applied in the same way as homogeneous equilibria, considering the effects of changes in concentration, pressure, and temperature.

Exothermic reaction and temperature

If the forward reaction is exothermic, the backward reaction is favoured by increasing the temperature.

Increasing pressure in a reaction

Increasing the pressure would favour the forward reaction.

Molecules on each side of the reaction

There are 4 molecules on the left and side and 2 molecules on the right hand side.

Effect of pressure on reaction direction

An increase in pressure will favour the side with the fewest number of molecules.

Effect of temperature on endothermic reaction yield

If the forward reaction is endothermic, a high temperature would increase the yield of products.

Reason for high temperature favoring endothermic reaction

A high temperature would favour the endothermic pathway to oppose the increase.

Catalyst effect on reaction rate

A catalyst speeds up the rate of both the forward and backward reactions.

Effect of increasing pressure on equilibrium

Increasing pressure shifts the equilibrium to favour the side with fewer gas molecules.

Equilibrium shift with decreasing reactant concentration

Equilibrium shifts to the left to reduce the effect of a decrease in reactant.

Equilibrium shift with decreasing pressure

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

Definition of reaction quotient, Q

The reaction quotient, Q, is calculated using the same equation as the equilibrium constant expression, but with non-equilibrium concentrations of reactants and products.

Expression for Q compared to K

The expression for Q is the same as K, but uses non-equilibrium concentrations.

What does Q = K indicate?

When Q = K, the reaction is at equilibrium.

Reaction direction when Q < K

If Q < K, the reaction will proceed to the right in favor of the products.

Reaction direction when Q > K

If Q > K, the reaction will proceed to the left in favor of the reactants.

Using Q to determine equilibrium

By calculating Q using concentration values and comparing it to K, we can determine if a reaction is at equilibrium (Q = K) or not (Q ≠ K).

Is Q a fixed value?

Q is not a fixed value and can be measured at any time, unlike K which is constant at a given temperature.

Information provided by Q

Q provides information about how far a reaction is from equilibrium and in which direction the reaction will proceed to reach equilibrium.

Calculating Q with concentrations

Q can be calculated using initial concentrations of reactants and products, or concentrations at any point during the reaction.

Q's relation to reaction progress

As a reaction progresses towards equilibrium, Q approaches the value of K.

Concentration units in equilibrium calculations

Sometimes amounts in moles are given, and concentrations need to be calculated using the volume of the reaction mixture.

Calculating concentration from moles and volume

Concentration (mol dm-3) = amount of substance (mol) / volume (dm3).

Approximation when K < 10-3

When K < 10-3, the initial concentration of reactants can be approximated as the equilibrium concentration of reactants.

Assumption for small K value approximation

The assumption is that the change from the initial amount of reactant to the equilibrium amount is close to zero when K is very small.

Use of approximation method for small K values

The approximation method for small K values can always be used regardless of the value of K.

Equilibrium constant (K)

K = fraction numerator open square brackets straight H subscript 2 straight O close square brackets open square brackets CH subscript 3 COOC subscript 2 straight H subscript 5 close square brackets over denominator open square brackets straight C subscript 2 straight H subscript 5 OH close square brackets open square brackets CH subscript 3 COOH close square brackets end fraction

Value of K for the reaction

K = fraction numerator open square brackets 0.5 close square brackets open square brackets 0.3 close square brackets over denominator open square brackets 0.6 close square brackets open square brackets 0.4 close square brackets end fraction = 0.625

Effect of ethyl ethanoate concentration increase

If the concentration of ethyl ethanoate increases by 0.25 mol dm-3, the concentration of ethanoic acid decreases by 0.25 mol dm-3.

Concentration of ethanol

Concentration = moles / volume in dm3; Concentration = 0.15 / 0.25 = 0.6 mol dm-3

Gibbs energy change equation

ΔG° = -RT ln K

Negative ΔG° and K relationship

A negative ΔG° always indicates that K > 1, meaning the products are favored at equilibrium.

ΔG° = 0 and K relationship

When ΔG° = 0, the equilibrium constant K = 1, meaning neither reactants nor products are favored at equilibrium.

Equilibrium constant and reaction kinetics

The equilibrium constant is independent of reaction kinetics and provides no information about individual rates of reaction.

Feasibility of a reaction

A negative ΔG° indicates that a reaction is feasible (spontaneous) under standard conditions.

ΔG° and equilibrium concentrations relationship

A negative ΔG° indicates that the equilibrium concentration of products is greater than the equilibrium concentration of reactants.

ΔG° and equilibrium constant relationship

As ΔG° becomes more negative, the value of the equilibrium constant increases.

Units for ΔG°

ΔG° is typically expressed in kJ mol-1 in the equation ΔG° = -RT ln K.

Rearranged equation for ln K

ln K = - fraction numerator straight capital delta straight G over denominator RT end fraction

Temperature dependence of ΔG° and K

The relationship between ΔG° and K is temperature-dependent, as shown by the presence of T in the equation ΔG° = -RT ln K.