Topic 7: Equilibrium SL Le Chatelier

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Topic 7 Equilibrium SLLe Chateliers Principle

• 7.2.3 Apply Le Chateliers 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.

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LeChateliers Principle

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

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Le Chatelier Translated

• When you take something away from a system at
  equilibrium, the system shifts in such a way as
  to replace what youve taken away.
• When you add something to a system at
  equilibrium, the system shifts in such a way as
  to use up what youve added.

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Le Chatelier Example 1
A closed container of ice and water at
equilibrium. The temperature is raised.
Ice Energy lt-- gt Water
The equilibrium of the system shifts to the
_______ to use up the added energy.
right
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Le Chatelier Example 2
A closed container of N2O4 and NO2 at
equilibrium. NO2 is added to the container.
N2O4 (g) Energy lt - - gt 2 NO2 (g)
The equilibrium of the system shifts to the
_______ to use up the added NO2.
left
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Le 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
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Le Chatelier Example 4
A closed container of N2O4 and NO2 at
equilibrium. The pressure is increased.
N2O4 (g) Energy lt - - gt 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
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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.

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Example

• If I increase the pressure, where is the shift?
• (right)
• If I decrease the pressure, where is the shift?
  (left)
• 2SO2 O2 lt--gt 2SO3
• (3moles) (2moles)

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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

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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.

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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 lt -- gt 2HCl Na2SO4
• Cooling the above reaction causes the system to
  shift to the left less reactants, so need to
  make up more

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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 lt--gt 2SO3 energy
• Cooling the above reaction causes the system to
  shift to the right, to make up for the lost heat.

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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.

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Haber Process

• N2(g) 3H2(g) lt - - gt 2NH3(g) ?H -92 kJ/mol
• Mixtures 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.

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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.

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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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http//www.absorblearning.com/media/item.action?qu
ick128

• Animation of Haber process

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Ammonias 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)

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Contact Process

• Production of sulfuric acid by the oxidation of
  sulfur.
• Sulfur is burnt in air to form sulfur dioxide
• S(s) O2(g) lt - - gt SO2(g)
• Sulfur dioxide is mixed with air and passed over
  vanadium(V)oxide catalyst to produce sulfur
  trioxide.
• 2SO2(g) O2(g) lt - - gt 2SO3(g) ?H -196 kJ/mol
• Sulfur trioxide is reacted with water to produce
  sulfuric acid.
• SO3(g) H2O(l) ? H2SO4(l)

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More money, more SO3

• 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.

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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

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http//www.absorblearning.com/media/item.action?qu
ick12b

• Contact process animation

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Uses of H2SO4

• Fertilizers (converting insoluble phosphate rock
  into soluble phosphates)
• Polymers
• Detergents
• Paints
• Pigments
• Petrochemical industry
• Processing of metals
• Electrolyte in car batteries

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• Le Chateliers principle is a memory aid, it
  doesnt 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/essentialche
  mistry/flash/lechv17.swf

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• Haber process notes
• http//www.chemguide.co.uk/physical/equilibria/hab
  er.html