Chemical Equilibrium

• What is equilibrium?

• Expressions for equilibrium constants, Keq;

• Calculating Keq using equilibrium concentrations;

• Factors that affect equilibrium;

• Le Chatelier’s Principle

What is Equilibrium?

This is not Equilibrium?

Chemical Equilibrium

• Consider the following reactions:

CaCO3(s) + CO2(aq) + H2O(l) Ca2+(aq) + 2HCO3

-(aq) ..(1)

and

Ca2+(aq) + 2HCO3-(aq) CaCO3(s) + CO2(aq) + H2O(l) ..(2)

Reaction (2) is the reverse of reaction (1).

At equilibrium the two opposing reactions occur at the

same rate.

Concentrations of chemical species do not change once

equilibrium is established.

Expression for Equilibrium Constant

Consider the following equilibrium system:

wA + xB ⇄ yC + zD

Keq =

• The numerical value of Keq is calculated using the

concentrations of reactants and products that exist at

equilibrium.

xw

z

[B][A]

[D][C]y

Expression and Value of Equilibrium Constant for a Reaction

• The expression for K depends on the equation

• The value of K applies to that equation; it does

not depend on how the reaction occurs;

• Concentrations used to calculate the value of K

are those measured at equilibrium.

Expressions for Equilibrium Constants

Examples:

N2(g) + 3H2(g) ⇄ 2NH3(g); Keq =

PCl5(g)⇄ PCl3(g) + Cl2(g); Keq =

CH4(g) + H2(g)⇄ CO(g) + 3H2(g);

Keq =

3

22

2

3

]][H[N

][NH

][PCl

]][Cl[PCl

5

23

O]][H[CH

][CO][H

24

3

2

Relationships between chemical equations and

the expressions of equilibrium constants

• The expression of equilibrium constant depends on how the

equilibrium equation is written. For example, for the

following equilibrium:

H2(g) + I2(g)⇄ 2 HI(g);

• For the reverse reaction:

2HI(g)⇄ H2(g) + I2(g);

]][I[H

[HI]

22

2

eq K

eq2

22eq 1/

[HI]

]][I[H ' KK

Homogeneous & Heterogeneous Equilibria

Homogeneous equilibria:

CH4(g) + H2O(g) ⇄ CO(g) + 3H2(g);

CO(g) + H2O(g) ⇄ CO2(g) + H2(g);

Heterogeneous equilibria:

CaCO3(s) ⇄ CaO(s) + CO2(g);

HF(aq) + H2O(l) ⇄ H3O+(aq) + F-(aq);

PbCl2(s) ⇄ Pb2+(aq) + 2 Cl-(aq);

Equilibrium Constant Expressions for

Heterogeneous System

Examples:

CaCO3(s)⇄ CaO(s) + CO2(g);

K = [CO2] or K = PCO2;

HF(aq) + H2O(l)⇄ H3O+(aq) + F-(aq);

[HF]

]F][O[H

-

3eq

K

PbCl2(s) ⇄ Pb2+

(aq) + 2Cl-(aq);

K = [Pb2+][Cl-]2

Le Chatelier’s Principle states that:

When a system at equilibrium is stressed, the equilibrium will shift in the direction that will relieve the stress.

What are “stresses” to an equilibrium?

• change in pressure

• change in concentration

• change in temperature

Changes in PRESSURE

• only affect gases

RULE: If the pressure on a system increases, the shift will be towards the side of the eqn. with the LOWER # of moles of gas

Changes in PRESSURE

•How do you figure out the number of moles of gas?

•Add up the coefficients in the balanced eqn.

EXAMPLE 1:

3 H2 (g) + N2 (g) 2 NH3 (g)

If P increases… shift to RIGHT side

because there are 4 moles of gas on left side, only 2 moles of gas on right side.

EXAMPLE 2:

H2 (g) + I2 (g) 2 HI (g)

If P increases… there will be NO SHIFT

because there are 2 moles of gas on the left side & 2 moles of gas on the right side.

Changes in CONCENTRATION

RULE:

If the [concentration ] of substance on one side of eqn. increases, equilibrium will shift towards the other side.

EXAMPLE 3:3 H2 (g) + N2 (g) 2 NH3 (g)

If [N2] increases…

shift towards RIGHT side.

If [NH3] increases…

shift towards LEFT side.

Changes in CONCENTRATION

RULE:

If the [concentration] of substance on one side of eqn. decreases, equilibrium will shift towards that side.

EXAMPLE 4:

3 H2 (g) + N2 (g) 2 NH3 (g)

If [H2] decreases…

shift towards LEFT side.

If [NH3] is removed…

shift towards RIGHT side.

Changes in CONCENTRATION

RULE:

If the [concentration] of substance on one side of eqn. increases, substances on same side of eqn. will decrease.

Substances on other side will increase.

In other words…

Same side of eqn. = opposite direction

Opposite side of eqn. = same direction

EXAMPLE 5:

4 HCl (g) + O2 (g) 2 H2O (g) + 2 Cl2 (g)

If [O2] decreases… shift towards LEFTside.

[HCl] increases (same side as O2, so opposite direction)

EXAMPLE 5:4 HCl (g) + O2 (g) 2 H2O (g) + 2 Cl2 (g)

If [O2] decreases…

[H2O] decreases (opposite side from O2, so same direction)

[Cl2] decreases

EXAMPLE 5: still!

4 HCl (g) + O2 (g) 2 H2O (g) + 2 Cl2 (g)

If [H2O] increases…

shift towards LEFT side

[HCl] increases

EXAMPLE 5:

4 HCl (g) + O2 (g) 2 H2O (g) + 2 Cl2 (g)

If [H2O] increases…

[O2] increases

[Cl2] decreases

Changes in TEMPERATURE

RULE:

The word “heat” or a # of J, kJ, or calshould be treated as another reactant or product.

Follow same rules as with concentration.

If heat is added to start the rxn. & the temp. increases…

Heat is located on the left side of the eqn.

It is an endothermic rxn.

The Keq value increases.

If heat is given off by the rxn. & the temp. increases…

Heat is located on the right side of the eqn.

It is an exothermic rxn.

The Keq value decreases.

EXAMPLE 6:

2 H2O (g) 2 H2 (g) + O2 (g) + 16 kcal

(exothermic)

If T increases…

shift towards LEFT side

Changes in Keq value

RULE:

Only changes in temperature affect the Keqvalue.

Changes in pressure and/or concentration do NOT affect the Keq value.