Ch. 14: Chemical Equilibrium
Dr. Namphol Sinkaset
Chem 201: General Chemistry II
I. Chapter Outline
I.
II.
III.
IV.
V.
VI.
Introduction
The Equilibrium Constant (K)
Values of Equilibrium Constants
The Reaction Quotient (Q)
Equilibrium Problems
Le Châtelier’s Principle
I. Introduction
• Equilibrium will be the focus for the next
several chapters.
• Most reactions are reversible, meaning
they can proceed in both forward and
reverse directions.
• This means that as products build up,
they will react and reform reactants.
• At equilibrium, the forward and
backward reaction rates are equal.
I. Example Equilibrium
II. Equilibrium Concentrations
• Equilibrium does not mean that
concentrations are all equal!!
• However, we can quantify concentrations
at equilibrium.
• Every equilibrium has its own equilibrium
constant.
II. The Equilibrium Constant
• equilibrium constant: the ratio at
equilibrium of the [ ]’s of products raised
to their stoichiometric coefficients
divided by the [ ]’s of reactants raised to
their stoichiometric coefficients.
• The relationship between a balanced
equation and equilibrium constant
expression is the law of mass action.
II. The Equilibrium Constant
• For a general equilibrium aA + bB  cC + dD,
the equilibrium expression is:
II. Sample “Problem”
• Write the equilibrium constant
expression for the reaction:
2H2(g) + O2(g)  2H2O(g).
II. Physical Meaning of K
• Large values of K mean that the
equilibrium favors products, i.e. there
are high [ ]’s of products and low [ ]’s of
reactants at equilibrium.
• Small values of K mean that the
equilibrium favors reactants, i.e. there
are low [ ]’s of products and high [ ]’s of
reactants at equilibrium.
II. Rules for Manipulating K
• If the equation is reversed, the equilibrium
constant is inverted.
II. Rules for Manipulating K
• If the equation is multiplied by a factor, the
equilibrium constant is raised to the same
factor.
II. Rules for Manipulating K
• When chemical
equations are
added, their
equilibrium
constants are
multiplied together
to get the overall
equilibrium
constant.
II. Sample Problem
• Predict the equilibrium constant for the
first reaction given the equilibrium
constants for the second and third
reactions.
CO2(g) + 3H2(g)  CH3OH(g) + H2O(g)
CO(g) + H2O(g)  CO2(g) + H2(g)
CO(g) + 2H2(g)  CH3OH(g)
K1 = ?
K2 = 1.0 x 105
K3 = 1.4 x 107
II. K in Terms of Pressure
• Up to this point, we’ve been using
concentration exclusively in the
equilibrium expressions.
• Partial pressures are proportional to
concentration via PV = nRT.
• Thus, for gas reactions, partial
pressures can be used in place of
concentrations.
II. Two Different K’s
• For the reaction 2SO3(g)  2SO2(g) + O2(g),
we can write two equilibrium expressions.
II. Relationship Between
Concentration and Pressure
• To be able to
convert between Kc
and Kp, we need a
relationship between
concentration and
pressure.
II. Converting Between Kc and Kp
II. Converting Between Kc and Kp
• The Δn is the change in the number of moles
of gas when going from reactants to products.
• When does Kp equal Kc?
II. Sample Problem
• Methanol can be synthesized via the
reaction CO(g) + 2H2(g)  CH3OH(g). If
Kp of this reaction equals 3.8 x 10-2 at
200 °C, what’s the value of Kc?
II. Heterogeneous Equilibria
• If an equilibrium contains pure solids or pure
liquids, they are not included in the
equilibrium constant expression.
III. Values of K
• Values of K are most easily calculated by
allowing a system to come to equilibrium
and measuring [ ]’s of the components.
• For the equilibrium H2(g) + I2(g)  2HI(g),
let’s say equilibrium [ ]’s at 445 °C were
found to be 0.11 M, 0.11 M, and 0.78 M
for molecular hydrogen, molecular iodine,
and hydrogen iodide, respectively.
III. Kc for a H2/I2 Mixture
• Note that units are
not included when
calculating K’s.
• Thus, equilibrium
constants are
unitless.
III. Equilibrium [ ]’s Vs. K
• For any reaction, the equilibrium [ ]’s will
depend on the initial [ ]’s of reactants or
products.
• However, no matter how you set up the
reaction, the value of the equilibrium
constant will be the same if the
temperature is the same.
III. Equilibrium [ ]’s Vs. K
IV. The Reaction Quotient
• What happens when we mix reactants
together and wait?
• Can we predict what will happen when
we have a mixture of reactants and
products?
• The reaction quotient, Qc or Qp, is used
to predict in which direction an
equilibrium will move.
IV. Formula for Qc or Qp
• You already know the formula because it’s
the same as for Kc or Kp!!
• The difference is, we don’t know if the
reaction is at equilibrium, thus, we cannot set
the ratio equal to K!
• For the reaction aA + bB  cC + dD:
IV. Using Q
• The value of Q relative to K tells you
whether the reaction will form more
products or more reactants to reach
equilibrium.
 Q < K means reaction forms products.
 Q > K means reaction form reactants.
 Q = K means reaction is at equilibrium.
IV. Sample Problem
• Consider the reaction N2O4(g)  2NO2(g)
with Kc = 5.85 x 10-3. If a reaction mixture
contains [NO2] = 0.0255 M and [N2O4] =
0.0331 M, which way will the reaction
proceed?