Chemical Equilibrium

Chemical Equilibrium

Ch. 15

Equilibrium A state where the reactants and products

remain constant over time. For some reactions, the equilibrium

position favors the products and the reaction appears to have gone to completion (amount of reactants is negligible)-”Equilibrium lies to the right” (in direction of the products)Example: 2 H2 + O2 2 H2O

Other reactions only occur to a small extent with the product virtually undectable-”Equilibrium lies to the left” (in direction of the reactants)Example: 2 CaO 2Ca + O2

Equilibrium is not static

Because the concentrations do not change, it appears that the reaction has stopped.

Instead, equilibrium is highly dynamic with the reactions continuing to occur in both directions at the same rate.

products

reactants

As the concentration of the reactants decreases, the forward reaction slows down.

As the concentration of the products increases, the rate of the reverse reaction increases

Eventually, the rates become equal.

Factors affecting equilibrium position:

Initial concentrations*** Energies of reactants and products “organization” of reactants and

products.

***The factor to be addressed in this unit.

Equilibrium Constants

Through experimentation and observation, the Law of Mass Action was proposed.

The law suggests that for a reaction of the following type:jA + kB lC + mD

the following equilibrium expression is used to represent the reaction.

K = [C]l [D]m

[A]j [B]k

K is the equilibrium constant

[ ] represent the concentrations at equilibrium

Coefficients become the exponents

Practice

Write the equilibrium expression for the following equation:

4 NH3 + 7 O2 < -- > 4 NO2 + 6 H2O

Determining the Equilibrium Constant

The equilibrium constant can be calculated at a given temperature if the equilibrium concentrations of the reaction components are known.

N2 + 3H2 < -- > 2 NH3

[NH3] = 3.1 x 10-2 M[N2] = 8.5 x 10-1 M[H2] = 3.1 x 10-3 M

Calculate the value for K

Other methods for Determining Equilibrium Constants

If the reaction is reversed, 2NH3 < --> N2 + 3H2

then, K/ = 1/K If the reaction is multiplied by a

factor, ½ N2 + 3/2 H2 < -- > NH3

then, K// = (K)1/2

Application from the lab At a given temperature:

1) K always has the same value regardless of the starting concentrations2) the equilibrium concentrations will not always be the same(The set of equilibrium concentrations is called the equilibrium position.)

There is only one equilibrium constant at a particular temperature, but there is an infinite number of equilibrium positions.

The specific equilibrium position depends on the starting concentrations, but the equilibrium constant does not.

See example 15.1, page 632

Heterogeneous Equilibrium

Many reactions involve reaction components in more than one phase.

Example: CaCO3 (s) < -- > CaO (s) +CO2 (g) Because the concentrations of pure solids

and liquids cannot change, they are not included in the equilibrium expression.

K = [CO2] Write the equilibrium expression for the

following equation:PCl5(s) < -- > PCl3(l) + Cl2(g)

Homework

The following problems may now be completed:

15.6, 15.13, and 15.15 on pages 659-660

Equilibrium Involving Pressure

Equilibrium involving gases can be described in terms of pressures (as well as concentrations)

PV = nRT or P = (n/V)RT where n/V is the concentration of the gas.

The equilibrium constant in terms of partial pressures in written as KP

Sample Problem

Given the following reaction:2NO(g) + Cl2 (g) < -- > 2NOCl(g)

PNOCl = 1.2 atm PNO = 0.050 atm PCl2 = 0.30 atm

Calculate the value of Kp

Relationship between K and KP

KP = K(RT)Δn

(see pages 639-647 view how this equation was derived)

R = 0.08206 T is the temperature in Kelvin Δn = the difference in the sum of the

coefficients for the products and the reactants.

Example: jA + kB < -- > lC + mD Δn = (l + m) – (j + k)

Sample Problem

Using the value for Kp obtained in the previous sample problem, calculate the value of K for the reaction at 250C2NO(g) + Cl2(g) < -- > 2NOCl(g)

Applications of the Equilibrium Constant

Knowing the equilibrium constant allows us to predict several important features of the reaction.1) the tendency of the reaction to occur (but not the speed)2) whether a given set of concentrations represent an equilibrium condition

3) the equilibrium position that will be achieved from a given set of initial concentrations.

The Extent of a Reaction The tendency for a reaction to occur is

indicated by the magnitude of the equilibrium constant.

A value of K much larger than 1 means that at equilibrium the reaction will consist mostly of product-the equilibrium lies to the right. (The reaction goes to completion)

A value of K much smaller than 1 means that at equilibrium the reaction will consist mostly of reactants-the equilibrium lies to the left. (The reaction does not occur to any given extent)

The size of K and the time required to reach equilibrium are not directly related.

Reaction Quotient When given a set of reaction components, it is

helpful to know if the mixture is at equilibrium or, if not, in what direction the system must shift to reach equilibrium.

To determine the direction of the move toward equilibrium, we use the reaction quotient (Q).

The reaction quotient is obtained by applying the law of mass action to the initial concentrations instead of equilibrium concentrations.

If Q = K, the system is at equilibrium; no shift will occur

If Q > K, the system shifts to the left. If Q < K, the system shifts to the right. Complete the sample problem on page 645.

Le Chatelier’s Principle

Le Chatelier’s Principle states that if a change is imposed on a system at equilibrium, the position of equilibrium will shift in a direction that tends to reduce that change.

Effect of a Change in Concentration

If a component is added to a reaction system at equilibrium (at constant T and P), the equilibrium position will shift in the direction that lowers the concentration of that component (away from).

If a component is removed from the system, the system will shift in the direction that increases the concentration of that component (towards)

Copy the following equation:

As4O6(s) + 6C(s) < -- > As4(g) + 6CO (g)

In which direction will the equilibrium position shift if CO is added?

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0% 1. Left2. Right3. No shift will occur

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In which direction will the equilibrium position shift if C is removed?

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In which direction will the equilibrium position shift if CO is removed?

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Effect of a Change in Pressure Adding an inert gas has no effect on the

equilibrium position (has no effect on the concentrations or partial pressures of the reactants or products).

Increasing the volume of the container decreases the pressure. The system responds by increasing the pressure through the production of more gaseous molecules (a shift toward the side with the greatest number of gas molecules).

Decreasing the volume of the container results in the opposite occurrence.

What shift in the equilibrium position will occur if the volume is reduced for the following process:

P4(s) + 6Cl2(g) < -- > 4PCl3(l)

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PCl3 (g) + Cl2 (g) < -- > PCl5 (g)

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PCl3(g) + 3NH3(g) < -- > P(NH2)3(g) + 3HCl(g)

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Effect of a Change in Temperature

Changing the temperature changes the value of K.

In exothermic reactions (heat is given off and is therefore written as a product), the equilibrium position will shift to the left if the temperature is increased and to the right if the temperature is decreased.

For endothermic reactions, the opposite will occur.

For the following reaction, predict how the value of K changes as the temperature is increased.

N2(g) + O2(g) < -- > 2NO(g) ΔH = 181 kJ

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0% 1. Increases2. Decreases3. Remains the same

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For the following reaction, predict how the value of K changes as the temperature in increased.2SO2(g) + O2(g) < -- > 2SO3(g) ΔH = -198 kJ

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0% 1. Increases2. Decreases3. Remains the same

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Copy the following equation

N2(g) + 3H2(g) < -- > 2NH3(g) + 92.94 kJ

How will the equilibrium position shift if the temperature increases?

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How will the equilibrium position shift if the volume increases?

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0% 1. Left2. Right 3. No shift will occur

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How will the equilibrium position shift if NH3 is removed?

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How will the equilibrium position shift if N2 is added?

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How will the equilibrium position shift if some Ar(g) is added?

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Homework

All homework problems that accompany Chapter 15 should be completed for class tomorrow.

Pg. 660 # 15.27, 15.33, 15.35, 15.45, and 15.55

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Chemical Equilibrium