Chapter 16: PRINCIPLES OF REACTIVITY: CHEMICAL EQUILIBRIA
Chapter 16 Problem Set
Pages 778-783
9, 11, 13, 15, 19, 21, 27, 30, 31, 39
In principle all chemical reactions are reversible. Most reactions continue until they reach an
equilibrium, a balance between the forward and reverse reaction.
Chemical systems can be forced to move to new equilibrium position by outside influences.
16.1 THE NATURE OF THE EQUILIBRIUM STATE
Chemical equilibria are dynamic
The forward and reverse reactions occur at the same rate
No net change in concentration.
[R]
[P]
16.2 THE EQUILIBRIUM CONSTANT
How can we describe the equilibrium position of a chemical reaction?
At equilibrium if [P] > [R] equilibrium lies on the product side
if [R] > [P] equilibrium lies on the reactant side
The relative concentration can be represented by the equilibrium constant, K
Equilibrium constant expression
K
[P]
[R]
The nature and properties of the equilibrium state will be the same, no matter the direction of
approach.
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Consider the formation of HI from H2, and I2, in a 1.00 L vessel at 425 °C.
H2(g) + I2(g)  2 HI(g)
The chemical reaction is:
Initial and Equilibrium Concentrations
H2(g)
I2(g)
HI(g)
Initial concentration, mol/L
0.0175
0.0175
 concentration
-0.0138
-0.0138
+0.0276
Equilibrium Concentration, mol/L
0.0037
0.0037
0.0276
0
Putting these equilibrium values in the expression
K
[HI] 2
[H 2 ][I 2 ]
K
[.0276] 2
[.0037] [.0037]
K = 56
Repeat the experiment beginning only with HI
Initial and Equilibrium Concentrations
Initial concentration, mol/L
 concentration
Equilibrium Concentration, mol/L
H2(g)
I2(g)
HI(g)
0
0
0.0350
+0.00369
+0.00369
-0.00738
0.00369
0.00369
0.02762
Putting these equilibrium values in the expression
K
[HI] 2
K = 56
[H 2 ][I 2 ]
The ratio is always the same for all the experiments at 425°C, whether the reaction is approached as a
forward or reverse reaction and regardless of the initial concentration.
2
For the general reaction
aA + bB  cC + dD
The equilibrium expression is
k
[C]c [D] d
[A] a [B] b
Important features of this equation are
●
Product concentrations appear in the numerator.
●
Reactant concentrations appear in the denominator.
●
Each concentration is raised to the power of its stoichiometric coefficient in the balanced equation.
●
The value of the constant K depends on the particular reaction and on the temperature.
Writing Equilibrium Constant Expressions
Reactions involving solids and/or water
Solid reactants and products are not included in the equilibrium expression.
For example the oxidation of sulfur
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S8(s) + O2(g)  SO2(g)
would be expected to have the equilibrium constant expression K 
[SO 2 ]
[S8 ]1/8[O 2 ]
However it is an experimental fact that the concentration of O2 and SO2 are not changed by the amount
of sulfur present as long as there is some sulfur present.
Compared to gases, concentration of solids is very high, thus concentration of solid is not significantly
changed and is not included in the equilibrium constant expression.
The equilibrium constant expression is K 
[SO 2 ]
[O 2 ]
Because the water concentration is very high in dilute solutions, the molar concentration of water is
essentially unchanged. Therefore the molar concentration of water is not included in the equilibrium
expression.
For example the reaction of NH3(g) + H2O(l)  NH4+(aq) + OH-(aq)
Has the equilibrium constant expression K 
3
[NH 4 ] [OH - ]
[NH 3 ]
Expressing Concentrations: Kc and Kp
Concentration in equilibrium expressions usually given in M, so K may be written as Kc
Equilibrium constants involving gases can be written in terms of partial pressure and designated as Kp
See pg 752 and Closer Look pg 753
EXERCISE 16.1 Page 752 Writing Equilibrium Constant Expressions
Use (aq) and (g)
Don’t use (s) and (l)
Manipulating Equilibrium Expressions
In general, when the stoichiometric coefficients of a balanced equation are multiplied by some factor,
the equilibrium constant for the new equation (k2) is the old equilibrium constant (K1) raised to the
power of the multiplied factor.
For example ½ H2(g) + ½ I2(g)  HI(g) using same concentrations as previous examples.
Equilibrium constants for a reaction and its reverse are the reciprocals of one another.
It is often useful to add two equations together to obtain the equation for a net process.
In general, when two or more equations are added to produce a net equation, the equilibrium constant
for the net equation is the product of the equilibrium constants for the added equations.
EXAMPLE 16.1 Page 755 Manipulating Equilibrium Constant Expressions
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EXERCISE 16.2 Page756 Manipulating Equilibrium Constant Expressions
Problem-Solving Tips and Ideas Page 756
1.
2.
3.
4.
5.
EXERCISE 16.3 Page 756 Manipulating Equilibrium Constant Expressions
The Meaning of the Equilibrium Constant
The value of the equilibrium constant indicates whether a reaction is product or reactant favored. In
addition it can be used to calculate how much product will be present at equilibrium.
A large value of K means that reactants are converted largely to products at eq.
Products are strongly favored over reactants. See pg 757
K>>1: Reaction is product-favored; equilibrium concentrations of products are greater than
equilibrium concentrations of reactants. “Reaction has gone to completion”
The reverse also holds true.
K<<1: Reaction is reactant-favored; equilibrium concentrations of reactants are greater than than
equilibrium concentrations of products.
Equilibrium constant values for a few reactions are given in Table 16.1. These reactions occur to
widely varying extents, as shown by wide range of values of K.
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EXERCISE 16.4 Page 757 The Equilibrium Constant and Extent of Reaction
16.3 THE REACTION QUOTIENT
The reaction quotient has the same form as the equilibrium constant but Q differs from K in that the
concentrations in the expression are not necessarily equilibrium concentrations.
The following general statements can be made
●
By calculating the reaction quotient we can predict the direction in which a reaction will occur to a
greater extent.
●
At a particular temp the value of K for a particular reaction is constant. Ratio of products to
reactants is constant.
●
Concentrations can vary as long as ratio remains the same.
●
The ratio of the concentrations is called the reaction quotient, Q.
●
If Q = k then system is in equilibrium.
●
If Q > K system contains too much product and to reach equilibrium, some product must reform
reactant.
●
If Q < K, then there is not enough product and the reactant must proceed further to the product
side.
An equilibrium plot pg 759 fig 16.6
[product]
slope corresponds to ratio
Q>K
Q<K
[reactant]
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EXAMPLE 16.2 Page 761 The Reaction Quotient
EXERCISE 16.5 Page 761 The Reaction Quotient
16.4
CALCULATING AN EQUILIBRIUM CONSTANT
When the values of the concentrations of all of the reactants and products are known at equilibrium
calculating an equilibrium constant simply involves substituting the data into the equilibrium
expression.
More often initial conc are known, 1 of the eq conc is measured, and the other must be calc.
General procedure is to set up an ICE box
2 SO2 + O2  2 SO3 eq express is K =
EXAMPLE 16.3 Page 762 Calculating an Equilibrium Constant
EXAMPLE 16.4 Page 763 Calculating an Equilibrium Constant
EXAMPLE 16.5 Page 764 Calculating an Equilibrium Constant
EXAMPLE 16.6 Page 764 Calculating an Equilibrium Constant
EXERCISE 16.6 Page 765 Calculating an Equilibrium Constant
EXERCISE 16.7 Page 765 Calculating an Equilibrium Constant
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16.5 USING EQUILIBRIUM CONSTANTS IN CALCULATIONS
Another type of situation arises when you know the value of K as well as the initial number of moles,
concentrations, or partial pressures of reactants and products and need to find the quantities present at
equilibrium.
Once again .
1. Quantities of reactants and products must be in conc or press (not as amts)
2. It is sometimes possible to assume that a change in a conc of reactant is so small that it can be
ignored. Rule of thumb is conc of reactant can be ignored if it is less than 5% of the smallest
quantity initially present.
EXAMPLE 16.7 Page 766 Calculating a Concentration from an Equilibrium Constant
EXAMPLE 16.8 Page 768 Calculating a Concentration from an Equilibrium Constant
EXERCISE 16.8 Page 769 Calculating a Concentration from an Equilibrium Constant
EXERCISE 16.9 Page 769 Calculating a Concentration from an Equilibrium Constant
16.6 DISTURBING A CHEMICAL EQUILIBRIUM; LeChatelier's Principle
There are 3 common ways to disturb an eq.
1. change temp
2. change conc of R or P
3. change V or P
Le Chatelier's principle: A change in any of the factors that determine the equilibrium conditions of a
system will cause the system to change in such a manner as to relieve the stress.
Effect of Temperature Changes on Equilibria
Table 16.2 Page
770 Effects of Disturbances on Equilibrium
Disturbance
Addition of R
Addition of P
Decrease in V
Increase in P
Increase in V
Decrease in P
Rise in T
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Change
Some additional R is
consumed
Some additional P is
consumed
P decreases
P increases
Heat energy is
Effect on Eq
Shift 
Effect on K
None
Shift 
None
Shift toward fewer
gas particles
Shift toward more gas
particles
Shift in endothermic
None
None
Change
Drop in T
consumed
Heat energy is
released
direction
Shift in exothermic
direction
Change
EXERCISE 16.10 Page 772 LeChatelier's Principle
EXAMPLE 16.9 Page 773 The Effect of Concentration Changes on an Equilibrium
EXERCISE 16.11 Page 775 The Effect of Concentration Changes on an Equilibrium
EXERCISE 16.12 Page 776 Effect of Concentration and Volume Changes on Equilibrium
16.7 IS THERE LIFE AFTER EQUILIBRIUM?
A reaction that is spontaneous will go until the system reaches equilibrium (a battery that has run down
has reached equilibrium. Many important reactions never reach equilibrium – for example a product
may be continually removed from the system. These reactions are said to run under nonequlibrium
conditions.
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