Topic 7: Equilibrium SLLe Chatelier’s Principle

7.2.3 Apply Le Chatelier’s principle to predict 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 process

Suitable examples include the Haber and Contact processes.

LeChatelier’s PrincipleLeChatelier’s Principle

• When a system at equilibrium is placed When a system at equilibrium is placed under stress, the system will undergo a under stress, the system will undergo a change in such a way as to relieve that change in such a way as to relieve that stress.stress.

Le Chatelier Translated:

• When you take something away from a When you take something away from a system at equilibrium, the system system at equilibrium, the system shifts in such a way as to replace what shifts in such a way as to replace what you’ve taken away.you’ve taken away.

• When you add something to a system When you add something to a system at equilibrium, the system shifts in at equilibrium, the system shifts in such a way as to use up what you’ve such a way as to use up what you’ve added.added.

Le Chatelier Example #1Le Chatelier Example #1

A closed container of ice and water at equilibrium. The temperature is raised.

Ice + Energy <-- > Water

The equilibrium of the system shifts to the _______ to use up the added energy.

right

Le Chatelier Example #2Le Chatelier Example #2

A closed container of N2O4 and NO2 at equilibrium. NO2 is added to the container.

N2O4 (g) + Energy < - - > 2 NO2 (g)

The equilibrium of the system shifts to the _______ to use up the added NO2.

left

Le Chatelier Example #3Le Chatelier Example #3

A closed container of water and its vapor at equilibrium. Vapor is removed from the system.

water + Energy vapor

The equilibrium of the system shifts to the _______ to replace the vapor.

right

Le Chatelier Example #4Le Chatelier Example #4

A closed container of N2O4 and NO2 at equilibrium. The pressure is increased.

N2O4 (g) + Energy < - - > 2 NO2 (g)

The equilibrium of the system shifts to the _______ to lower the pressure, because there are fewer moles of gas on that side of the equation.

left

Pressure Changes to system

If the volume decreases, the concentration increases, and there will be a shift to the side with the less amount of moles.

If the volume increases, the concentration decreases, and there will be a shift to the side with the more amount of moles.

Example;

If I increase the pressure, where is the shift?(right)

If I decrease the pressure, where is the shift? (left)

2SO2 + O2 <--> 2SO3

(3moles) (2moles)

Effect of Concentration

1. If you add more reactant, it shifts to the right increasing the formation of product, using up the reactants.

2. If you add product, it shifts to the left

3. If you remove product, it shifts to the right, increasing the formation of product.

4. If you remove reactant, it shifts to the left

Effect of temperature

• Energy is treated as a reactant if endothermic equation, and as a product if exothermic equation.

• If cooling a system, then it shifts so more heat is produced.

• If heating a system, then it shifts so extra heat is used up.

Example for temp. changes for Endothermic Reaction

Heating the below reaction causes the system to shift to the right = more products, because you treat energy like a reactant.

2NaCl +H2SO4 + energy < -- > 2HCl + Na2SO4

Cooling the above reaction causes the system to shift to the left = less reactants, so need to make up more

Effect of temp change on exothermic reactions

• Heating the below reaction causes the system to shift to the left, to use up the extra heat.

2SO2 + O2 <--> 2SO3 + energy

• Cooling the above reaction causes the system to shift to the right, to make up for the lost heat.

The effect of a catalyst on equilibrium

• Adding a catalyst speeds up a reaction by providing an alternative mechanism with a lower activation energy, thus speeding up both the forward and backward reaction rate.

• It shortens the time needed to attain equilibrium concentrations

• It has no effect on the position of equilibrium, however equilibrium will be attained more quickly.

Haber Process

• N2(g) + 3H2(g) < - - > 2NH3(g) ΔH= -92 kJ/mol

– Mixture’s volume is compressed and passed over a heated iron catalyst.

– Conditions for his equilibrium is critical. • High pressure is favourable due to 4 moles on left

and 2 moles on right. Increased pressure causes a shift to the left, favouring product formation.

• This is expensive to due and most production plants will resist compressing gases in terms of operating costs. Compromise will be met.

Compromise

• This is an exothermic reaction, so low temperatures would be favourable to produce product.– Low temps mean low reaction rates, so we

may get a higher yield but it will take a long time to get it. Not good for business.

– A compromise temp, as well as the use of a catalyst will aid in speeding up the reaction to a more acceptable standard.

Typical conditions

• Pressure between 20-100 MPa (200-1000 atm)• Temperatures around 700 K• The reaction is not allowed to reach equilibrium,

because reaction rate decreases as we approach equilibrium, and typically only 20% of N2 and H2 is converted.

• The gases are cooled and NH3 is condensed and removed, leaving unused N2 and H2 available for further production.

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Animation of Haber process

Ammonia’s Uses

• Manufacture of fertilizers (ammonia salts and urea)

• Manufacturing nitrogen used in polymers for the fabrication of nylon

• Used in the production of explosives (TNT, dynamite)

Contact Process• Production of sulfuric acid by the oxidation of

sulfur.

1. Sulfur is burnt in air to form sulfur dioxide• S(s) + O2(g) < - - > SO2(g)

2. Sulfur dioxide is mixed with air and passed over vanadium(V)oxide catalyst to produce sulfur trioxide.

• 2SO2(g) + O2(g) < - - > 2SO3(g) ΔH= -196 kJ/mol

3. Sulfur trioxide is reacted with water to produce sulfuric acid.

• SO3(g) + H2O(l) H2SO4(l)

• High pressure would favour the formation of SO3 in the 2nd step, however its too expensive.

• Reactants are compressed to 2 atm to achieve the desired flow rate in the reactor.

• Pure O2 would drive the equilibrium to the right, however its an unnecessary expense.

• Low temperatures, because its exothermic, would be best, but it slows the rate too much.

More money, more SO3…

Compromised conditions

• Temp between 700-800 K (fast initial reaction rate)

• The use of a finely divided V2O5 catalyst

• Oxidation is done in converters at lower temperatures (slows reaction rate)– Overall conversion is 90% to SO3

http://www.absorblearning.com/media/item.action?quick=12b

Contact process animation

Uses of H2SO4

• Fertilizers (converting insoluble phosphate rock into soluble phosphates)

• Polymers• Detergents• Paints• Pigments• Petrochemical industry• Processing of metals• Electrolyte in car batteries

• Le Chatelier’s principle is a memory aid, it doesn’t explain why these changes occur.

• Listen carefully and read over text pages to help you develop further understanding of explanation.

• http://www.mhhe.com/physsci/chemistry/essentialchemistry/flash/lechv17.swf

• Haber process notes

• http://www.chemguide.co.uk/physical/equilibria/haber.html