chemical equilibrium
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Notes on Chemical equilibriumTRANSCRIPT
Unit 2: Chemical Equilibrium andIonic Equilibrium
Reversible Reactions There exist a group of reactions in which the direction of chemical change can be easily
reversed by changing the conditions under which the reaction is taking place
E.g. When hydrated copper(II)sulphate is heated the blue colour of the crystals change to
the white appearance of the anhydrous salt
CuSO4.5H2O (s) → CuSO4(s) + 5H2O(g)
hydrated salt (blue) anhydrous salt (white)
However anhydrous copper(II)sulphate may be changed to the blue hydrated form simply
by adding water
CuSO4(s) + 5H2O(g) → CuSO4.5H2O (s)
There exist a group of reactions in which the direction of chemical change can be easily
reversed by changing the conditions under which the reaction is taking place
E.g. When hydrated copper(II)sulphate is heated the blue colour of the crystals change to
the white appearance of the anhydrous salt
CuSO4.5H2O (s) → CuSO4(s) + 5H2O(g)
hydrated salt (blue) anhydrous salt (white)
However anhydrous copper(II)sulphate may be changed to the blue hydrated form simply
by adding water
CuSO4(s) + 5H2O(g) → CuSO4.5H2O (s)
Reversible Reactions
Because the reaction can easily be reversed it is known as a reversible reaction. The
equation for the reversible reaction is:
CuSO4.5H2O (s) CuSO4(s) + 5H2O(g)
The indicating that the reaction may proceed in one direction or the other
according to the conditions under which it is carried out
Reversible reactions attain equilibrium during the course of the reaction
Because the reaction can easily be reversed it is known as a reversible reaction. The
equation for the reversible reaction is:
CuSO4.5H2O (s) CuSO4(s) + 5H2O(g)
The indicating that the reaction may proceed in one direction or the other
according to the conditions under which it is carried out
Reversible reactions attain equilibrium during the course of the reaction
Equilibrium• What is equilibrium?
For equilibrium to exist there needs to be ‘sameness’ and ‘changelessness’, e.g. if a cup of hot
water is placed on a table then the temperature of the hot water will be different from the
surroundings (disequilibrium). However if the cup of hot water is left for awhile then the
temperature of the water will eventually be the same as the surroundings (sameness) and
it no longer changes (changelessness)
• There exists different types of equilibrium, i.e. static and dynamic
• Static equilibrium occurs when a object is at rest and is in a state of equilibrium, for e.g. two persons
of the same weight on a see-saw
• Dynamic equilibrium involves two opposing processes occurring at the same rate, for e.g. walking on a
treadmill
• Chemical equilibrium is a kind of dynamic equilibrium
• What is equilibrium?
For equilibrium to exist there needs to be ‘sameness’ and ‘changelessness’, e.g. if a cup of hot
water is placed on a table then the temperature of the hot water will be different from the
surroundings (disequilibrium). However if the cup of hot water is left for awhile then the
temperature of the water will eventually be the same as the surroundings (sameness) and
it no longer changes (changelessness)
• There exists different types of equilibrium, i.e. static and dynamic
• Static equilibrium occurs when a object is at rest and is in a state of equilibrium, for e.g. two persons
of the same weight on a see-saw
• Dynamic equilibrium involves two opposing processes occurring at the same rate, for e.g. walking on a
treadmill
• Chemical equilibrium is a kind of dynamic equilibrium
Dynamic Equilibrium
• When a chemical reaction takes place in a container which prevents the entry or escape
of any of the substances involved in the reaction, the quantities of these components
change as some are consumed and others are formed
• Eventually this change will come to an end, after which the composition will remain
unchanged as long as the system remains undisturbed
• The system is then said to be in its equilibrium state or at equilibrium
• A chemical reaction is in equilibrium when there is no tendency for the quantities of
reactants and products to change
• When a chemical reaction takes place in a container which prevents the entry or escape
of any of the substances involved in the reaction, the quantities of these components
change as some are consumed and others are formed
• Eventually this change will come to an end, after which the composition will remain
unchanged as long as the system remains undisturbed
• The system is then said to be in its equilibrium state or at equilibrium
• A chemical reaction is in equilibrium when there is no tendency for the quantities of
reactants and products to change
• The direction in which we write a chemical reaction (and thus which components are
considered reactants and products) is subjective
• E.g. H2 + I2 → 2 HI (synthesis of hydrogen iodide)
and 2 HI → H2 + I2 (dissociation of hydrogen iodide)
The equations represent the same chemical reaction system in which the roles of the
components are reversed, and both yield the same mixture of components when the
change is completed
It makes no difference whether we start with two moles of HI or one mole each of H2 and
I2; once the reaction has run to completion, the quantities of these two components will
be the same
Once this equilibrium composition has been attained, no further change in the quantities
of the components will occur as long as the system remains undisturbed
• The direction in which we write a chemical reaction (and thus which components are
considered reactants and products) is subjective
• E.g. H2 + I2 → 2 HI (synthesis of hydrogen iodide)
and 2 HI → H2 + I2 (dissociation of hydrogen iodide)
The equations represent the same chemical reaction system in which the roles of the
components are reversed, and both yield the same mixture of components when the
change is completed
It makes no difference whether we start with two moles of HI or one mole each of H2 and
I2; once the reaction has run to completion, the quantities of these two components will
be the same
Once this equilibrium composition has been attained, no further change in the quantities
of the components will occur as long as the system remains undisturbed
Dynamic Equilibrium
Chemical equilibrium can be represented graphically
A substance having zero initial concentration is the product of the reaction
Chemical equilibrium can be represented graphically
A substance having zero initial concentration is the product of the reaction
Dynamic Equilibrium Whether we start with an equi-molar mixture of H2 and I2 or a pure sample of hydrogen
iodide, the composition after equilibrium is attained will be the same
The equilibrium composition is independent of the direction of the reaction
Whether we start with an equi-molar mixture of H2 and I2 or a pure sample of hydrogen
iodide, the composition after equilibrium is attained will be the same
The equilibrium composition is independent of the direction of the reaction
Dynamic Equilibrium Consider the equation:
aA + bB cC + dD
If we combine the two reactants A and B, the forward reaction starts immediately as a
result the forward reaction is favoured or by convention the equilibrium shifts to the
right
As the products C and D begin to build up, then the reverse process gets underway i.e.
the reversed reaction is favoured or the equilibrium shifts to the left
Consider the equation:aA + bB cC + dD
If we combine the two reactants A and B, the forward reaction starts immediately as a
result the forward reaction is favoured or by convention the equilibrium shifts to the
right
As the products C and D begin to build up, then the reverse process gets underway i.e.
the reversed reaction is favoured or the equilibrium shifts to the left
Dynamic Equilibrium The rate at which the equation proceed can be written as
rate of forward reaction = kf [A]a [B]b
rate of reverse reaction = kr [C]c [D]d
The proportionality constants kf and kr are called rate constants and the quantities in
square brackets represent concentrations
As the reaction proceeds, the rate of the forward reaction diminishes while that of the
reverse reaction increases
Eventually the two processes are proceeding at the same rate, and the reaction is at
equilibrium, i.e. the
rate of forward reaction = rate of reverse reaction
kf [A]a [B]b = kr [C]c [D]d
The rate at which the equation proceed can be written as
rate of forward reaction = kf [A]a [B]b
rate of reverse reaction = kr [C]c [D]d
The proportionality constants kf and kr are called rate constants and the quantities in
square brackets represent concentrations
As the reaction proceeds, the rate of the forward reaction diminishes while that of the
reverse reaction increases
Eventually the two processes are proceeding at the same rate, and the reaction is at
equilibrium, i.e. the
rate of forward reaction = rate of reverse reaction
kf [A]a [B]b = kr [C]c [D]d
Equilibrium Constants
The rate of the reversible reaction is affected by the concentration, temperature and
pressure (gaseous systems)
The gas equilibrium constant (K) can be expressed using the molar concentration of the
components present in the reaction (Kc) or from the partial pressure of the components
present in the system (Kp)
The rate of the reversible reaction is affected by the concentration, temperature and
pressure (gaseous systems)
The gas equilibrium constant (K) can be expressed using the molar concentration of the
components present in the reaction (Kc) or from the partial pressure of the components
present in the system (Kp)
In equilibrium equations, even though the arrows point both ways ( ) we usually
associate the left as reactants and the right as products
The products are written at theTOP of the fraction (the numerator)
The reactants are written at the BOTTOM of the fraction (the denominator)
The concentrations of the products and reactants are always raised to the power of
their coefficient in the balanced chemical equation
If any of the reactants or products are solids or liquids, their concentrations are equal to
one because they are pure substances
In equilibrium equations, even though the arrows point both ways ( ) we usually
associate the left as reactants and the right as products
The products are written at theTOP of the fraction (the numerator)
The reactants are written at the BOTTOM of the fraction (the denominator)
The concentrations of the products and reactants are always raised to the power of
their coefficient in the balanced chemical equation
If any of the reactants or products are solids or liquids, their concentrations are equal to
one because they are pure substances
Calculations
The following concentrations were measured for an equilibrium mixture at 500 K:
[N2] = 3.0 x 10-2 M; [H2] = 3.7 x 10-2 M;
[NH3] = 1.6 x 10-2 M.
Calculate the equilibrium constant at 500 K for the reaction
N2(g) + 3H2(g) 2NH3(g)
Kc = [NH3]2
[N2][H2]3
(1.62 x 10-2)2
(3.0 x 10-2)(3.7 x 10-2)3
Kc = 1.7 x 10-2
Kc’ = [N2][H2]3 (3.0 X 10-2)(3.7 X 10-2)3 = 5.9 X 10-3
[NH3]2 (1.62 X 10-2)2
OR
Kc’ = 1/ 1.7 X 10-2 = 5.9 X 10-3
Kc = [NH3]2
[N2][H2]3
(1.62 x 10-2)2
(3.0 x 10-2)(3.7 x 10-2)3
Kc = 1.7 x 10-2
Kc’ = [N2][H2]3 (3.0 X 10-2)(3.7 X 10-2)3 = 5.9 X 10-3
[NH3]2 (1.62 X 10-2)2
OR
Kc’ = 1/ 1.7 X 10-2 = 5.9 X 10-3
The oxidation of sulphur dioxide to give sulphur trioxide is a important step in the
industrial process for the synthesis of sulphuric acid
2SO2(g) + O2(g) 2 SO3(g)
The following equilibrium concentrations were measured at 800 K: [SO2] = 3.0 x 10-3 M;
[O2] = 3.5 x 10-3 M; [SO3] = 5.0 x 10-2 M.
Calculate the equilibrium constant at 800 K for the reaction.
Kc = [SO3]2
[SO2]2[O2]
Kc = (5.0 x 10-2)2 Kc = 7.936 x 104
(3.0 x 10-3)2 (3.5 x 10-3)
Kc’ = [SO2]2[O2] = (3.0 x 10-3)2 (3.5 x 10-3) = 1.26 x 10-5
[SO3]2 (5.0 x 10-2)2
1. An equilibrium mixture of gaseous O2 , NO, and NO2 at 500K contains 1.0 x 10-3 M
O2 and 5.0 x 10-2 M NO2. At this temperature, the equilibrium constant Kc for the
reaction
2 NO (g) + O2(g) 2NO2(g)
is 6.9 x 105 . What is the concentration of NO?
Kc = [NO2]2
[NO]2[O2]
[NO] = [NO]2
√ [O2]Kc
[NO] = √ (5.0 x 10-2)2 / (1.0 x 10-3) (6.9 x 105)
[NO] = √3.6 x 10-6
[NO] = ± 1.9 x 10-3 M
Kc = [NO2]2
[NO]2[O2]
[NO] = [NO]2
√ [O2]Kc
[NO] = √ (5.0 x 10-2)2 / (1.0 x 10-3) (6.9 x 105)
[NO] = √3.6 x 10-6
[NO] = ± 1.9 x 10-3 M
2. N2O4 (l) is an important component of rocket fuel. At 25 ºC N2O4 is a colorless gas
that partially dissociates into NO2. Equilibrium is established in the reaction
N2O4 (g) 2NO2 (g) at 25 ºC
Given: 3.00 L container, 7.64 g N2O4 and 1.56 g NO2
What is the Kc for this reaction?
2. N2O4 (l) is an important component of rocket fuel. At 25 ºC N2O4 is a colorless gas
that partially dissociates into NO2. Equilibrium is established in the reaction
N2O4 (g) 2NO2 (g) at 25 ºC
Given: 3.00 L container, 7.64 g N2O4 and 1.56 g NO2
What is the Kc for this reaction?
HINT
Step 1: Convert grams to moles
Step 2: Convert moles to Molarity (moles/L)
Step 3: Write the Equilibrium constant for Kc
Kc = 4.61 x 10-3
HINT
Step 1: Convert grams to moles
Step 2: Convert moles to Molarity (moles/L)
Step 3: Write the Equilibrium constant for Kc
Kc = 4.61 x 10-3
3. The H2/CO ratio in mixtures of carbon monoxide and hydrogen (called synthesis gas) is
increased by the water gas shift reaction CO(g) + H2O(g) CO2(g) + H2 (g),
which has an equilibrium constant Kc = 4.24 at 800K. Calculate the equilibrium
concentrations of CO2, H2, CO and H2O at 800K if only CO and H2O are present
initially at concentrations of 0.150M.
4. Calculate the equilibrium concentrations of N2O4 and NO2 at 25ºC in a vessel that
contains an initial N2O4 concentration of 0.0500M. The equilibrium constant for the
reaction N2O4(g) 2NO2(g) is 4.64 x 10-3 at 25ºC
3. The H2/CO ratio in mixtures of carbon monoxide and hydrogen (called synthesis gas) is
increased by the water gas shift reaction CO(g) + H2O(g) CO2(g) + H2 (g),
which has an equilibrium constant Kc = 4.24 at 800K. Calculate the equilibrium
concentrations of CO2, H2, CO and H2O at 800K if only CO and H2O are present
initially at concentrations of 0.150M.
4. Calculate the equilibrium concentrations of N2O4 and NO2 at 25ºC in a vessel that
contains an initial N2O4 concentration of 0.0500M. The equilibrium constant for the
reaction N2O4(g) 2NO2(g) is 4.64 x 10-3 at 25ºC
Source: McMurray & J., Fay, R. (1998). Chemistry 2nd Edition, New Jersey- Prentice-hall Inc. p.520
To calculate Kp using KC, the following equation is used
Kp = Kc (RT)∆n
∆n = (Total moles of gas on the products side) -
(Total moles of gas on the reactants side)
Therefore for the equation:
aA + bB cC + dD
∆n = (d + c) - (a + b) [The coefficients]
R = 0.0821 L.atm mol-1K-1 from the ideal gas law constant (PV = nRT)
T = temperature in Kelvin
To calculate Kp using KC, the following equation is used
Kp = Kc (RT)∆n
∆n = (Total moles of gas on the products side) -
(Total moles of gas on the reactants side)
Therefore for the equation:
aA + bB cC + dD
∆n = (d + c) - (a + b) [The coefficients]
R = 0.0821 L.atm mol-1K-1 from the ideal gas law constant (PV = nRT)
T = temperature in Kelvin
Factors that Alter the Composition of anEquilibrium Mixture
The principal goal of a chemical reaction is to effect the maximum conversion of reactants to
products
For reversible reactions this may prove difficult as the reaction can take place in two
directions
However in utilizing Le Chatelier’s principle one is able to alter conditions in order to get the
maximum concentration of a desired substance
Several factors that can be exploited to alter the composition of an equilibrium mixture are
(a) temperature
(b) pressure and volume
(c) concentration
(d) catalyst
The principal goal of a chemical reaction is to effect the maximum conversion of reactants to
products
For reversible reactions this may prove difficult as the reaction can take place in two
directions
However in utilizing Le Chatelier’s principle one is able to alter conditions in order to get the
maximum concentration of a desired substance
Several factors that can be exploited to alter the composition of an equilibrium mixture are
(a) temperature
(b) pressure and volume
(c) concentration
(d) catalyst
Le Chatelier’s Principle Le Chatelier’s principle states that if a chemical system is in equilibrium and one of the
factors involved in the equilibrium is altered, the equilibrium will shift so as to tend to
annul the effect of the change
Effect of Temperature Change
If the temperature is increased, then according to Le Chatelier’s principle, the reaction
that is favoured is the reaction that lowers the temperature i.e. endothermic reaction
If the temperature is decreased, the according to Le Chatelier’s principle the reaction that
increases the temperature is favoured i.e. exothermic reaction
Effect of Pressure The effect of pressure is much more noticeable in reactions occuring in the gaseous state
If the pressure is increased then the reaction that is favoured is the one that lowers the
pressure, that is the reaction that has fewer number of molecules present in the reaction
mixture
Types of reaction Effect of increase in totalpressure
Effect of decrease intotal pressure
Effect of increase in totalpressure
Effect of decrease intotal pressure
Increase the number ofmolecules left to right2O3(g) 3O2(g)
Equilibrium shifts to the leftproducing more O3
Equilibrium shifts to theright producing more O2
Decrease in the number ofmolecules from left o rightN2(g) + 3H2(g) 2NH3(g)
Equilibrium shifts to theright producing more NH3
Equilibrium shifts to the leftproducing more N2 and H2
No change in the number ofmolecules left to rightH2(g) + I2(g) 2 HI
No effect equilibriummaintained
No effect equilibriummaintained
Effect of Concentration
If the concentration of the reactant or product decreases then the reaction that is
favoured is the one that replenishes the decreased concentration.
In addition if the concentration of the reactant or product increases then the reaction that
is favoured is the one that uses up the reactant or product
Effect of Catalyst A catalyst does not affect the equilibrium position, it only enables you get to the point of
equilibrium faster
If a reaction mixture is not at equilibrium, at catalyst accelerates the rate at which the
equilibrium is reached, but it does not affect the composition of the equilibrium mixture
A catalyst speeds up both the forward and reverse reactions by lowering the activation
energy
A catalyst does not affect the equilibrium position, it only enables you get to the point of
equilibrium faster
If a reaction mixture is not at equilibrium, at catalyst accelerates the rate at which the
equilibrium is reached, but it does not affect the composition of the equilibrium mixture
A catalyst speeds up both the forward and reverse reactions by lowering the activation
energy
Industrial Application of ChemicalEquilibrium Industrial preparation of ammonia (Haber Process)
Industrial preparation of hydrogen iodide
Haber Process
Source: http://www.chemguide.co.uk/physical/equilibria/haberflow.gif
N2(g) + 3H2(g) 2NH3(g)
N2(g) + 3H2(g) 2NH3(g) exothermic
(a) Pressure
Ammonia is produced from its element with a reduction of volume
If the pressure of the system is increased (keeping the temp. and conc. constant),
then Le Chatelier’s principle dictates that the system will oppose that change, i.e.
reduce the pressure
To achieve this the reaction proceeds in the direction the produces the lesser volume
An increase in pressure favours the production of ammonia
(a) Pressure
Ammonia is produced from its element with a reduction of volume
If the pressure of the system is increased (keeping the temp. and conc. constant),
then Le Chatelier’s principle dictates that the system will oppose that change, i.e.
reduce the pressure
To achieve this the reaction proceeds in the direction the produces the lesser volume
An increase in pressure favours the production of ammonia
N2(g) + 3H2(g) 2NH3(g) exothermic
(b) Temperature
The production of ammonia is exothermic, therefore according to Le Chatelier’s
principle a decrease in temperature favours the production of ammonia
As a result a catalyst of finely divided reduced iron is introduced to speed up the reaction
rate inspite the low temperature
N2(g) + 3H2(g) 2NH3(g) exothermic
(c) Concentration
An increase in the concentration of either N2 or H2 would result in the increase
production of ammonia
However in practice there is no advantage of increasing either reactants. This is because
the gases are mixed in the theoretical proportion of N2: H2, 1:3 by volume
N2(g) + 3H2(g) 2NH3(g) exothermic
(c) Concentration
An increase in the concentration of either N2 or H2 would result in the increase
production of ammonia
However in practice there is no advantage of increasing either reactants. This is because
the gases are mixed in the theoretical proportion of N2: H2, 1:3 by volume
H2(g) + I2(g) 2HI(g) exothermic
(a) Decrease in Temperature
(b) Increase in concentration of reactants
(c) Change in pressure does not affect the equilibrium…Why?