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IB Topic 7: Equilibrium
7.1: Dynamic equilibrium
7.1.1 Outline the characteristics of chemical and physical
systems in a state of equilibrium.
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7.1.1 Outline the characteristics of chemical
and physical systems in a state of equilibrium.
The reactions we have studied so far have gone to
completion. In other words the reaction proceeds until one
or more of the reactants runs out.
Mg + 2HCl  MgCl2 + H2
There are reactions that are reversible. In other words, the
reactions occur simultaneously in both directions
2SO2(g) + O2(g)  2SO3(g)
2SO3(g)  2SO2(g) + O2(g)
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7.1.1 Outline the characteristics of chemical
and physical systems in a state of equilibrium.
Reversible Reactions:
Reactions occurring simultaneously in both directions
2SO2(g) + O2(g)
2SO3(g)
Reaction 1: Sulfur dioxide reacts with oxygen to produce
sulfur trioxide. SO2(g) & O2(g) are reactants, SO3(g) is
the product.
Reaction 2: Sulfur trioxide decomposes to sulfur dioxide &
oxygen. SO3(g) is the reactant and SO2(g) & O2(g) are
products.
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7.1.1 Outline the characteristics of chemical
and physical systems in a state of equilibrium.
Chemical Equilibrium
• A state in which the rate
of the forward reaction
equals the rate of the
reverse reaction.
• Once equilibrium is
reached, the
concentrations of the
reactants and the
concentrations of the
products do not change.
• The reaction continues
but no change in
concentrations
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7.1.1 Outline the characteristics of chemical
and physical systems in a state of equilibrium.
2HI
H2 + I2
• Reaction 1 (Graphs 1 & 2)
Starting with a concentration
of 2.0 HI and 0 H2 or I2. HI
begins to decompose,
forming H2 & I2. As H2 & I2
form, they begin to react
forming HI. Eventually the
rates become equal so the
amount of HI reacted = the
amount of HI produced. The
reaction continues but no
change in concentrations.
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7.1.1 Outline the characteristics of chemical
and physical systems in a state of equilibrium.
2HI
H2 + I2
• Reaction 2 (Graphs 3 & 4)
Explain what is happening.
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7.1.1 Outline the characteristics of chemical
and physical systems in a state of equilibrium.
Physical System at Equilibrium
Liquid water evaporates to form
water vapor. At a given
temperature in a closed
system, water will evaporate
until the vapor reaches a
certain pressure. When that
occurs, equilibrium is
reached. Water still
evaporates but at the same
rate as water condensing.
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IB Topic 7: Equilibrium
7.2: The position of equilibrium
7.2.1 Deduce the equilibrium constant expression (Kc)
from the equation for a homogeneous reaction.
7.2.2 Deduce the extent of a reaction from the magnitude
of the equilibrium constant.
7.2.3 Apply Le Chatelier’s principle to predict the
qualitative effects of changes of temperature, pressure
and concentration on the position of equilibrium and on
the value of the equilibrium constant.
7.2.4 State and explain the effect of a catalyst on an
equilibrium reaction.
7.2.5 Apply the concepts of kinetics and equilibrium to
industrial processes.
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7.2.1 Deduce the equilibrium constant
expression (Kc) from the equation for a
homogeneous reaction.
Equilibrium Constant (Kc)
[ ] means concentration expressed in mol dm-3
When a system reaches equilibrium, the [reactants] stays
the same and the [products] stays the same. There is a
mathematical relationship between the [rcts] and [prod].
aA + bB
cC + dD
Kc = [C]c x [D]d
[A]a x [B]b
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7.2.1 Deduce the equilibrium constant
expression (Kc) from the equation for a
homogeneous reaction.
Write the equilibrium constant expression for the following:
Contact Process (manufacture of sulfuric acid)
2SO2(g) + O2(g)
2SO3(g)
Kc =
[SO3]2
[SO2]2 x [O2]
A homogeneous reaction is one in which all the reactants
and products are in the same phase. We can write an
equilibrium expression if the substances are all gases, all
liquids or all in aqueous solution.
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7.2.1 Deduce the equilibrium constant
expression (Kc) from the equation for a
homogeneous reaction.
Write the equilibrium constant expression for the following:
Haber Process (manufacture of ammonia)
3H2(g) + N2(g)
2NH3(g)
The dissociation of hydrogen iodide
2HI(g)
H2(g) + I2(g)
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7.2.2 Deduce the extent of a reaction
from the magnitude of the equilibrium
constant.
The equilibrium constant is a measure of the amount of
products at equilibrium compared with the amount of
reactants.
a) More products than reactants at equilibrium. The
reaction goes almost to completion.
Kc >>1
Reactants
Products
H2(g) + I2(g)
2HI(g) Kc = 794
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7.2.2 Deduce the extent of a reaction
from the magnitude of the equilibrium
constant.
The equilibrium constant is a measure of the amount of
products at equilibrium compared with the amount of
reactants.
b) More reactants than products at equilibrium. The
reaction hardly proceeds.
Kc <<1
Reactants
N2(g) + O2(g)
Products
2NO(g) Kc = 1 x 10-30
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7.2.2 Deduce the extent of a reaction
from the magnitude of the equilibrium
constant.
The equilibrium constant is a measure of the amount of
products at equilibrium compared with the amount of
reactants.
c) Reactants and products present in somewhat equal
amounts.
Kc ≈1
Reactants
C2H6O + C2H4O2
Products
C4H8O2 + H2O Kc = 4
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7.2.3 Apply Le Chatelier’s principle to predict the
qualitative effects of changes of temperature, pressure
and concentration on the position of equilibrium and on
the value of the equilibrium constant.
Henri-Louis Le Chatelier
(1850-1936)
French industrial chemist
If a system at equilibrium is
disturbed by a change in
temperature, pressure, or the
concentration of one of the
components, the system will
shift its equilibrium position so
as to counteract the effect of
the disturbance.
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7.2.3 Apply Le Chatelier’s principle to predict the
qualitative effects of changes of temperature, pressure
and concentration on the position of equilibrium and on
the value of the equilibrium constant.
N2(g) + 3H2(g)
2NH3(g) ΔHo = -92 kJ
Effects of Concentration Change
• If a chemical system is at equilibrium and we add a
substance (either a reactant or a product) the reaction
will shift to reestablish equilibrium by consuming part of
the added substance. Removal of a substance will result
in a shift that forms more of the substance.
• The value of Kc does not change (think paper clips)
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7.2.3 Apply Le Chatelier’s principle to predict the
qualitative effects of changes of temperature, pressure
and concentration on the position of equilibrium and on
the value of the equilibrium constant.
•
•
•
•
•
N2(g) + 3H2(g)
2NH3(g)
Added H2
ΔHo = -92.0 kJ
Reaction will shift to use
up some of the added H2.
Forward reaction
temporarily speeds up.
N2 used up as it reacts
with some of the extra H2.
More NH3 is being
produced.
Eventually a new
equilibrium is reached.
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7.2.3 Apply Le Chatelier’s principle to predict the
qualitative effects of changes of temperature, pressure
and concentration on the position of equilibrium and on
the value of the equilibrium constant.
N2(g) + 3H2(g)
•
•
•
•
2NH3(g)
ΔHo = -92 kJ
What effect will removing NH3 have on the equilibrium?
System will shift to make more NH3 so it will temporarily
speed up to the right.
Some N2 & H2 will react to produce more NH3.
At the new equilibrium there will be less N2, less H2, and
less NH3 than the original equilibrium.
The value of Kc remains the same.
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7.2.3 Apply Le Chatelier’s principle to predict the
qualitative effects of changes of temperature, pressure
and concentration on the position of equilibrium and on
the value of the equilibrium constant.
N2(g) + 3H2(g)
2NH3(g)
ΔHo = -92 kJ
What effect will adding NH3 have on the equilibrium?
What effect will removing N2 have on the equilibrium?
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7.2.3 Apply Le Chatelier’s principle to predict the
qualitative effects of changes of temperature, pressure
and concentration on the position of equilibrium and on
the value of the equilibrium constant.
N2(g) + 3H2(g)
2NH3(g) ΔHo = -92 kJ
Effects of Pressure/Volume Change
• If a chemical system is at equilibrium and we increase
the pressure (reduce the volume), the reaction will shift
toward the side having the fewest moles of gas.
Decreasing the pressure (increasing the volume) causes
a shift in the direction that produces more gas
molecules.
• Only affects systems containing gas molecules.
• The value of Kc does not change.
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7.2.3 Apply Le Chatelier’s principle to predict the
qualitative effects of changes of temperature, pressure
and concentration on the position of equilibrium and on
the value of the equilibrium constant.
N2(g) + 3H2(g)
2NH3(g) ΔHo = -92 kJ
What effect does increasing the pressure have on the
equilibrium?
• Reaction will shift toward the side with the fewest gas
molecules.
• The left side has 4 gas molecules (1N2 & 3H2). The right
side has 2 gas molecules. Reaction will shift to the right.
• N2 & H2 will react and more NH3 will be produced.
• Kc does not change
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7.2.3 Apply Le Chatelier’s principle to predict the
qualitative effects of changes of temperature, pressure
and concentration on the position of equilibrium and on
the value of the equilibrium constant.
N2(g) + 3H2(g)
2NH3(g) ΔHo = -92 kJ
What effect does decreasing the pressure have on the
equilibrium?
If a reaction has equal numbers of gas molecules on the
left and on the right, changing the pressure has no
effect.
2HI(g)
H2(g) + I2(g)
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7.2.3 Apply Le Chatelier’s principle to predict the
qualitative effects of changes of temperature, pressure
and concentration on the position of equilibrium and on
the value of the equilibrium constant.
N2(g) + 3H2(g)
2NH3(g) ΔHo = -92 kJ
Effects of Temperature Change
• Increasing temperature causes the equilibrium position
to shift in the direction that absorbs heat (endothermic).
• Decreasing temperature causes the equilibrium position
to shift in the direction that produces heat (exothermic).
• The value of Kc will change with a change in temp. If
the reaction shifts right, the value of Kc increases. If the
reaction shifts left, the Kc value decreases.
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7.2.3 Apply Le Chatelier’s principle to predict the
qualitative effects of changes of temperature, pressure
and concentration on the position of equilibrium and on
the value of the equilibrium constant.
N2(g) + 3H2(g)
2NH3(g) ΔHo = -92 kJ
What effect does increasing the temperature have on the
equilibrium?
• Increasing the temperature causes the reaction to shift
to use up some of the added heat (endothermic rx).
• The reaction as written is exothermic so the endothermic
rx is from right to left. The rx will shift left.
• [N2] increases, [H2] increases, [NH3] decreases.
• Kc value will decrease
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7.2.3 Apply Le Chatelier’s principle to predict the
qualitative effects of changes of temperature, pressure
and concentration on the position of equilibrium and on
the value of the equilibrium constant.
N2(g) + 3H2(g)
2NH3(g) ΔHo = -92 kJ
What effect does decreasing the temperature have on the
equilibrium?
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7.2.3 Apply Le Chatelier’s principle to predict the
qualitative effects of changes of temperature, pressure
and concentration on the position of equilibrium and on
the value of the equilibrium constant.
N2O4(g)
2NO2(g)
ΔHo = 58.0 kJ
Both gases are present in a flask at equilibrium. N2O4 is a
colorless gas while NO2 is brown.
1) What color will the contents of the flask be if the
pressure is increased? Explain.
2) State and explain three (3) ways the amount of NO2
production can be increased.
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7.2.4 State and explain the effect of a catalyst
on an equilibrium reaction.
1) A catalyst lowers the
activation energy barrier for
both the forward and the
reverse reactions.
2) Therefore a catalyst increase
the rates of both reactions by
the same factor.
3) A catalyst increases the rate
at which equilibrium is
achieved, but does not
change the final composition
of the substances.
4) The Kc value is not affected
by the presence of a catalyst.
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7.2.5 Apply the concepts of kinetics and
equilibrium to industrial processes.
N2(g) + 3H2(g)
2NH3(g) ΔHo = -92 kJ
The Haber process to manufacture ammonia
• Ammonia is an important starting point for the
production of fertilizers, nitric acid, explosives and
polymers (nylon).
• Under what conditions can will an industrial chemist run
this reaction to increase the yield of ammonia?
• An optimum temperature must be found
• Read pg. 133-134
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7.2.5 Apply the concepts of kinetics and
equilibrium to industrial processes.
2SO2(g) + O2(g)
2SO3(g) ΔHo = -192 kJ
The contact process to manufacture sulfuric acid
•
•
•
•
SO3(g) + H2O(l)
H2SO4(l)
Sulfuric acid is used in many chemical processes
Under what conditions can an industrial chemist run this
reaction to increase the yield of sulfur trioxide?
An optimum temperature must be found.
Read pg. 134
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