1. Chapter 15 Chemical Equilibrium
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Chapter 15

2. The Equilibrium Constant
For a general reaction
the equilibrium constant expression for everything in solution is
where Keq is the equilibrium constant.
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Chapter 15

3. The Equilibrium Constant
Gas Phase Reaction
For a general reaction in the gas phase
the equilibrium constant expression is
where Keq is the equilibrium constant.
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Chapter 15

4. The Equilibrium Constant
General Notes
Keq is based on the molarities (or partial pressures) of reactants and products at equilibrium. (Note the upper case K!)
We generally omit the units of the equilibrium constant.
The equilibrium constant expression boils down to “products over reactants”.
The same equilibrium is established no matter how the reaction is begun.
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5. The Same Equilibrium
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6. Significance of K
K expresses the relative amounts of products to reactants.
A larger K means more products are present at equilibrium (and vice versa!)
If K >> 1, then products dominate at equilibrium and equilibrium lies to the right.
If K << 1, then reactants dominate at equilibrium and the equilibrium lies to the left.
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Chapter 15

7. Significance of K
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Chapter 15

8. The Equilibrium Constant
The Direction of the Chemical Equation
An equilibrium can be approached from any direction.
Example:
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9. The Equilibrium Constant
Other Ways to Manipulate Chemical Equations and Keq Values
The equilibrium constant for a reaction which is the sum of other reactions is the product of the equilibrium constants for the individual reactions.
2NOBr(g) = 2NO(g) + Br2(g) Keq = 0.42
Br2(g) + Cl2(g) = 2 BrCl Keq = 7.2
2NOBr(g) + Cl2(g) = 2NO(g) BrCl(g) Keq =
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10. Heterogeneous Equilibria
Experimentally, the amount of CO2 does not depend on the
amounts of CaO and CaCO3. Why?
The concentration of a solid or pure liquid is its density (g/mL) divided by its molar mass (g/mol). Both the density and the molar mass are constants.
Thus, the concentrations of solids and pure liquids are constant.
We ignore the concentrations of pure liquids and pure solids in equilibrium constant expressions.
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12. Calculating Equilibrium Concentrations
Proceed as follows:
Tabulate initial and equilibrium concentrations (or partial pressures).
If an initial and equilibrium concentration is given for a species, calculate the change in concentration.
Use stoichiometry on the change in concentration line only to calculate the changes in concentration of all species.
Deduce the equilibrium concentrations of all species.
Usually, the initial concentration of products is zero.
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13. Predicting the Direction of Reaction
We define Q, the reaction quotient, as follows:
Q = K only at equilibrium.
If Q > K then the reaction must proceed in reverse to reach equilibrium.
If Q < K then the forward reaction must occur to reach equilibrium.
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15. Le Châtelier’s Principle
As the pressure increases, the amount of ammonia present at equilibrium increases.
As the temperature decreases, the amount of ammonia at equilibrium increases.
Can this be predicted?
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16. Le Châtelier’s Principle
If a system at equilibrium is disturbed, the system will move in such a way as to counteract the disturbance.
An equilibrium reaction adjusts the “position of equilibrium” to relieve any stress that is applied.
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17. Le Châtelier’s Principle
Reactant and Product Concentrations
If H2 is added while the system is at equilibrium, the system must respond to counteract the added H2
The system reacts to consume the H2 and produce NH3 until equilibrium is once again established.
So, [H2] and [N2] will decrease and [NH3] increases.
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19. Le Châtelier’s Principle
Reactant and Product Concentrations
Adding a reactant or product shifts the equilibrium away from the increase.
Removing a reactant or product shifts the equilibrium towards the decrease.
Usually we desire to maximize the amount of the desired product of a chemical reaction. Two strategies that are often used are (1) use excess of one of the reactants and (2) remove one of the products as it forms.
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20. Le Châtelier’s Principle
Volume and Pressure Changes
As volume is decreased pressure increases.
If pressure is increased the system will shift to counteract the increase. That is, the equilibrium shifts to remove gases and decrease pressure.
An increase in pressure favors the direction that has fewer moles of gas.
In a reaction with the same number of product and reactant moles of gas, pressure has no effect.
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21. Le Châtelier’s Principle
Volume and Pressure Changes
An increase in pressure (by decreasing the volume) causes the system to not be at equilibrium because the concentration of both gases is increased.
The system shifts to reduce the number of moles of gas.
Thus, it shifts to the right (the forward reaction is favored).
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22. Le Châtelier’s Principle
Effect of Temperature Changes
The equilibrium constant is temperature dependent.
For an endothermic reaction, H > 0 and heat can be considered as a reactant.
For an exothermic reaction, H < 0 and heat can be considered as a product.
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23. Le Châtelier’s Principle
Effect of Temperature Changes
Adding heat (i.e. heating the vessel) favors away from the increase:
if H > 0, adding heat favors the forward reaction,
if H < 0, adding heat favors the reverse reaction.
Removing heat (i.e. cooling the vessel), favors towards the decrease:
if H > 0, cooling favors the reverse reaction,
if H < 0, cooling favors the forward reaction.
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24. Le Châtelier’s Principle
The Effect of Catalysis
A catalyst lowers the activation energy barrier for both the forward and the reverse reaction.
Therefore, a catalyst will decrease the time taken to reach equilibrium.
A catalyst does not effect the composition of the equilibrium mixture.
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Chapter 15