Chemical Equilibrium
Chemical Equilibrium
 In a chemical system, when the concentrations of both
reactants and products reach a steady state, they are
said to be in a dynamic equilibrium.
 This occurs when both the forward rate and the
reverse rate are equal.
 Reactants are still making products, but products are
also making reactants.
 Static equilibrium is when no change is occurring –
like in a tug-o-war.
Chemical Equilibrium
 Consider the equilibrium
reaction of:
N2O4(g)  2 NO2(g)
 When equilibrium is
established (no matter
the starting amounts!),
the forward and reverse
rates are equal.
Chemical Equilibrium
 This DOES NOT
imply that the
concentrations are
equal!
 Rather only the
amounts remain
constant.
 Use of a double
arrow () indicates
an equilibrium.
Chemical Equilibrium
 Analogy with traffic flow on a
bridge.
 More traffic coming in to the
city than leaving.
 But, if the rates coming in are
constant and the rates going
out are constant, then it is a
dynamic equilibrium.
Equilibrium Constant, K
 For the N2O4 reaction, the forward rate is:
Ratef = kf [N2O4]1
 The reverse rate is:
Rater = kr [NO2]2
 Setting these equal to each other:
kf [N2O4]1 = kr [NO2]2
Equilibrium Constant, K
 Re-arranging to:
k f  NO2 

k r  N 2 O4 1
2
 Where the ratio of kr / kf = K, the equilibrium
constant.
K
 This concept of the equilibrium is also referred to as
the law of mass action.
 In general, for a reaction of the type:
aA + bB  cC + dD
 The equilibrium constant will be:
 C  D 
a
b
 A   B
c
K=
d
K
 Equilibrium amounts can be in moles per liter
(Molarity) or as partial pressures, if gases are present.
 The equilibrium constant, K, then is labeled as Kc or
Kp.
 As we do more types of equilibria, you will also see Ka,
Kb, Ksp, etc.
 K is always unitless, regardless of any reaction.
K
 How do we prove that the law of mass action works?
 Experiments, experiments, experiments.
K
 Experiments 3 and 4
show that equilibrium
can be achieved from
either direction.
 Experiment 3 starts with
all products and
Experiment 4 starts with
all reactants.
K
 Are Kc and Kp the same?
 Recall that PV = nRT from Chapter 10.
 Molarity = n / V, so…
 P = M(RT)
 Resulting in:
 Kp = Kc(RT)Dn
 Where Dn =
K
 The magnitude of K can be
used to predict whether the
reactants or products are
favored.
 When K << 1, then the
_____________ are
predominant.
 When K >> 1, then the
_____________ are
predominant.
K
 What happens to K if a reaction is reversed?
 N2O4  2 NO2 ; Kc = 0.212
 2 NO2  N2O4 ; Kc = ???
 What happens to K if a reaction is multiplied by a
factor?
 2 N2O4  4 NO2 ; Kc = ???
K
 What happens when two reactions are added together




like in Hess’ Law?
2 NOBr  2 NO + Br2 ; Kc = 0.014
Br2 + Cl2  2 BrCl ; Kc = 7.2
___________________________________
2 NOBr + Cl2  2 NO + 2 BrCl ; Kc = ???
K
 When all substances in an equilibrium are of the same
phase, then that equilibrium is homogeneous.
 When two different phases are present, the
equilibrium is said to be heterogeneous.
 The amounts of solids and/or liquids in the presence
of gases or aqueous compounds are in large excess.
 Thus, their concentrations do not change.
K
Consider the equilibrium for:
CaCO3(s)  CaO(s) + CO2(g)
Evaluating K
 If given equilibrium amounts of all reactants and
products, then a value for Kc can be found.
 Or – can find an unknown concentration if Kc is
known.
 Can also use basic stoichiometric relationships in
some problems.
 These require an ICE table.
Evaluating K
 ICE stands for Initial, Change, and Equilibrium.
 Set-up as a table under the reaction.
 Initial = starting amounts in Molarity.
 Change = includes a sign plus a variable and reflects
mole-to-mole amounts.
 Equilibrium = combines the I and C lines.
 Suppose a 4.0L flask has 0.20 moles of N2O4 initially
added to it, the ICE table set-up would be:
Evaluating K
N2O4  2 NO2
N2O4
Initial
Change
Equilibrium
NO2
Evaluating K
 If the equilibrium amounts of the reactants and
products are unknown, then an equation with a
variable must be set-up and solved.
 Must be given the value of K.
 Two common problems are the perfect square problem
and the quadratic problem.
Evaluating K
 Problems with partial pressures are based on Dalton’s
Law of Partial Pressures.
 Ptotal = Pa + Pb + Pc + …
 Thus, if the intital total pressure is known, then one
gas has a pressure of “x” and the other is this pressure
minus “x”.
Reaction Quotient
 Given amounts of all reactants and products, which
direction will the reaction move?
 The Q expression is identical to K.
 Can think of Q as a continuum.
Q = 0,
all reactants
Q = ,
all products
Reaction Quotient
 When Q < K, the system has too many reactants and
not enough products.
Reaction Quotient
 When Q > K, the system has too many products and
not enough reactants.
Reaction Quotient
 When Q = K, the system is in equilibrium.
LeChatelier’s Principle
 If a system in equilibrium is disturbed by a change in
temperature, pressure, or the concentration of one of
the components, the system will shift its equilibrium
position to counteract the disturbance.
 Certain reactions, like the production of NH3 require
the manipulation of an equilibrium to maximize the
product.
LeChatelier’s Principle
 Haber process – Fritz
Haber looked at P, T
effects on making
NH3.
 Reaction favors lower
T and higher P.
 Why?
LeChatelier’s Principle
 Changes in reactant or
product concentrations
will force the equilibrium
to re-balance just like a
two pan balance.
 Can add or remove
components in the
equilibrium system.
LeChatelier’s Principle
 N2 + 3 H2  2 NH3
 Add some H2 will result
in:
 Removing NH3 as it is
made will:
LeChatelier’s Principle
 Pressure changes
 As Boyle’s Law predicts, as P increases, V decreases.
 For a system to reduce its volume, it must reduce the
moles of gas present.
 For reactions with different moles of products and moles
of reactants, increasing the pressure will favor the side
with fewer total moles of gases.
 Of course, the opposite is true if the pressure is
decreased.
 Does this agree with Haber results?
LeChatelier’s Principle
 Changing the pressure at constant temperature does
NOT change the value of K.
 As the total pressure is increased, the relative partial
pressures increase, but not equally.
 Thus, the K value remains the same.
LeChatelier’s Principle
 Temperature changes
 Depends on whether the reaction is endo- or exothermic.
 Endothermic: Reactants + Heat  Products

As T increases, reaction shifts to right and K will ___________
 Exothermic: Reactants  Products + Heat

As T increases, reaction shifts to left and K will ____________
Is the reaction of:
N2O4(g)  2 NO2(g),
Exothermic or
Endothermic?
LeChatelier’s Principle
 If a catalyst is added to an
equilibrium system, then
it will speed up the
reaction in both
directions.
 The equilibrium
concentration, though, is
unchanged.
Ammonia Production
 N2(g) + 3H2(g)  2NH3(g); DH = -92 kJ
 Favors high or low temperatures?
 Favors high or low pressures?
 Removal of NH3 as it is made
 Catalyst – iron oxides
 Over 100 million tons produced per year using 1-2% of
the annual world’s energy supply.
 Plants and some bacteria can perform the same reaction
at much lower temperatures and pressures.
Ammonia Production