7.1 Chemical equilibrium
Chemical reactions can reach a state of balance in which
reactants are being converted to products at the same rate
as the products are converted back again to reactants. This
is called Dynamic equilibrium.
Two physical processes to illustrate dynamic equilibrium
Carbon dioxide in a sealed fizzy drink bottle
Carbon dioxide in air can dissolve in water. At the same
time, dissolved carbon dioxide can leave a solution. At
equilibrium, molecules are entering (and leaving) each
phase at the same rate, illustrated below:
Gas phase
air
water
Liquid phase
The change is reversible as shown as:
CO2 (g)
CO2 (aq)
Water in equilibrium
Water molecules in a drink can be in the liquid or gas
phase. As they move about, some water molecules are
evaporating at the same rate that other water molecules are
condensing. This is represented as:
H2O(l)
H2O(g)
Chemical equilibrium
Carbon dioxide can react with water to form
hydrogencarbonate ions:
CO2(aq) +
H2O(l)
forward reaction
HCO3- (aq) + H+(aq)
reverse reaction
Only a small fraction of carbon dioxide molecules react
with water. The reaction is reversible AND reaches a
dynamic equilibrium because HCO3- ions react with H+
ions to reform CO2 and H2O (the reverse reaction) as
quickly as the forward reaction happens.
Explaining chemical equilibrium
Concentration
If a mixture of nitrogen dioxide and carbon monoxide is heated in
a sealed container, the following reversible reaction occurs:
NO + CO2
NO2 + CO
Reaction
Equilibrium
NO2 (g) + CO(g)
NO(g) + CO2 (g)
At the start of the reaction, there is only nitrogen dioxide
and carbon monoxide (in high concentrations). As the
products are made (they start in very low concentrations),
some of the nitrogen monoxide and carbon dioxide
molecules are converted back to the reactants. The rate of
the forward reaction decreases as the rate of the reverse
reaction increases. Eventually both forward and reverse
reactions will happen at the same rate, so the overall
concentrations of reactants and products remains constant.
Position of equilibrium
The reaction conditions, such as temperature and pressure
will affect the concentrations of reactants and products at
equilibrium.
If the conditions change, then the position of equilibrium
shifts. The new conditions can result in different
concentrations of reactants and products, even through the
rate of the forwards and reverse reactions stays the same at
equilibrium.
If the forwards reaction is almost complete when
equilibrium is established, then we say that the position of
equilibrium has:
 Shifted to the right
 Shifted to favour the forwards reaction
 Shifted towards the products
If few reactants have been converted to products by the
time equilibrium has been established, then we say that the
position of equilibrium has been shifted:
 To the left
 In favour of the reactants
 In favour of the reverse reaction
Whether we start with reactants or products, we will
always end up with the same concentrations at equilibrium.
Consider the reaction:
H2 (g) + I2 (g)
2HI (g)
We could start with 1 mole of reactants OR 2 moles of
product….and always end up with identical amounts of HI
and I2 and H2.
H2, I2
time
HI
Concentration
Concentration
HI
H2, I2
time
The position of equilibrium is shifted by:
 The concentrations of reacting substances in solution
 The pressure of reacting gases
 The temperature
Henri Le Chatelier (1888) proposed some rules called Le
Chatelier’s principle:
If a system is in equilibrium, and a change is made in any
of the conditions, then the system responds to counteract
the change as much as possible.
A catalyst does NOT shift the position of equilibrium…it
just helps to reach equilibrium more quickly.
Changing the Concentration
The following reversible reaction reaches equilibrium:
Fe3+ (aq) + SCN- (aq)
[Fe(SCN)]2+ (aq)
Yellow
colourless
deep red
The red intensity can be used to measure how much
[Fe(SCN)]2+ is present.
Equilibrium shifts to the left  solution gets lighter due to
an decreased concentration of [Fe(SCN)]2+ ions.
Equilibrium shifts to the right  solution gets darker due
to an increased concentration of [Fe(SCN)]2+ ions.
Experiment: Add ammonium chloride to decrease the
concentration of Fe3+ (aq) ions (4Cl- + Fe3+  [FeCl4]-).
This causes the red colour to become paler, because
equilibrium has shifted to the left.
Rules for deciding how changes in concentration affect
equilibrium (based from experimental observations):
 Increasing the concentration of reactant causes the
equilibrium to move to the product side.
 Increasing the concentration of products causes the
equilibrium to move to the reactant side.
 Decreasing the concentration of reactants causes the
equilibrium to move to the reactant side.
 Decreasing the concentration of products causes the
equilibrium to move to the product side.
(In reactions in which the product is a gas, removing the
gas will help shift the reaction to the product side.)
Any change in concentration will shift the position of
equilibrium to counteract that change.
Changing the pressure
Gas-phase equilibria rules are:
 Increasing the pressure moves the equilibrium to the
side of the equation with fewer gas molecules as this
tends to reduce the pressure.
 Decreasing the pressure moves the equilibrium to the
side of the equation with more gas molecules as this
tends to increase the pressure.
Any change in pressure will shift the position of
equilibrium to counteract that change.
For the following two reactions (2 steps for manufacturing
methanol from methane), say how altering the pressure
will change the position of equilibrium;
CH4 (g) + H2O (g)
CO (g) + 3H2 (g)
Increasing the pressure….shifts left
Decreasing the pressure….shift right
CO (g) + 2H2 (g)
CH3OH (g)
Increasing the pressure….
Decreasing the pressure…
What conditions are needed in each step to maximize the
yield of methanol?
Step 1:
Step 2:
Changing the temperature
According to Le Chatelier’s principle:
 Heating a reversible reaction at equilibrium shifts the
reaction in the direction of the endothermic reaction.
 Cooling a reversible reaction at equilibrium shifts the
reaction in the direction of the exothermic reaction.
These rules were based on observations from:
2NO2 (g)
brown
Exothermic
N2O4 (g)
Endothermic
colourless
ΔH values when given are
for the forward reaction.
Nitrogen dioxide is a dark brown gas which exists in
equilibrium with its dimer, dinitrogen tetroxide.
The forwards reaction (forming dinitrogen tetroxide) is
exothermic and releases thermal energy to the
surroundings.
The reverse reaction (forming nitrogen dioxide) is
endothermic and thermal energy is taken in from the
surroundings.
Predict the observations that were first made when a sealed
container of the equilibrium mixture was:
Cooled on ice….
Heated in boiling water…
Linking chemical equilibria
The product of one equilibrium reaction may be the
reactant in another equilibrium reaction. For example:
CO2 (g)
CO2 (aq)
CO2 (aq) + H2O (l)
HCO3 (aq) + H+ (aq)
Use Le Chatelier’s principle to predict what will happen if
the concentration of gaseous CO2 is increased……
Chemical equilibria and steady state systems
Equilibrium is only established in a CLOSED SYSTEM
(sealed from the surroundings).
Thermal decomposition of calcium carbonate can reach
equilibrium in a closed system:
CaCO3 (s)
CaO (s) + CO2 (g)
In the open, CO2 is lost, so equilibrium is constantly
shifted to the ___________ to replace the lost CO2 until all
the CaCO3 is converted to CaO.
Open systems can reach a steady state, eg when methane
burns, the methane is used up as quickly as it is
replaced….so concentrations of reactants and products
remain constant.
Ozone production and destruction in the atmosphere is also
in a steady state:
O + O2  O3
O3 + hv  O2 + O
O + O3  O2 + O2
ozone production
ozone destruction
None of these reactions reaches equilibrium; the series of
reactions has reached a steady state.
 Do chemical ideas 7.1 p.170, questions 1-6.