Chemical Equilibrium

39
Unit 2: Chemical Equilibrium and Ionic Equilibrium

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Notes on Chemical equilibrium

Transcript of Chemical Equilibrium

Page 1: Chemical Equilibrium

Unit 2: Chemical Equilibrium andIonic Equilibrium

Page 2: Chemical 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)

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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

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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

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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

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• 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

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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

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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

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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

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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

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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)

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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

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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)

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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

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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.

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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

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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?

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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

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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?

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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

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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

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Source: McMurray & J., Fay, R. (1998). Chemistry 2nd Edition, New Jersey- Prentice-hall Inc. p.520

Page 25: Chemical Equilibrium

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

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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

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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

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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

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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

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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

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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

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Industrial Application of ChemicalEquilibrium Industrial preparation of ammonia (Haber Process)

Industrial preparation of hydrogen iodide

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Haber Process

Source: http://www.chemguide.co.uk/physical/equilibria/haberflow.gif

N2(g) + 3H2(g) 2NH3(g)

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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

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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

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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

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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?