chemical equilibria and le chatelier

• all reversible reactions reach a dynamic equilibrium state
• dynamic equilibrium means

  1. forward and backward reactions are occurring at equal rates simultaneously
  2. concentrations of reactants and products stay constant
• le chatelier's principle states if an external condition is changed the equilibrium will shift to oppose the change
• if temperature is increased the equilibrium will shift to oppose this and move in the endothermic direction to try to reduce the temperature by absorbing heat
• if temperature is decreased, the equilibrium will shift to oppose this and move in the exothermic direction to try to increase the temperature by giving out heat
• low temperatures may give a higher yield of product but will also result in slow rates of reaction
• often a compromise temperature is used that gives a reasonable yield and rate
• increasing pressure will cause the equilibrium to shift towards the side with fewer moles of gas to oppose the change and thereby reduce the pressure
• decreasing pressure will cause the equilibrium to shift towards the side with more moles of gas to oppose the change and thereby increase the pressure
• if the number of moles of gas is the same on both sides of the equation then changing pressure will have no effect on the position of equilibrium
• increasing pressure may give a higher yield of product and will produce a faster rate
• industrially high pressures are expensive to produce (high electrical energy costs for pumping gases to make high pressure) and equipment is expensive (to contain high pressures)
• increasing the concentration of a reactant causes equilibrium to shift to the right to reduce the effect of the increase
• increasing the concentration of a product or decreasing the concentration of a reactant causes equilibrium to shift to the left to reduce the effect
• a catalyst does not affect the position of equilibrium but speeds up the rate of reaction
• Haber process is an exothermic reaction

  N2 + 3H2 ⇌ 2NH3
• Haber process conditions

  T=450 degrees Celsius

  P=200-1000atm

  catalyst=iron
• haber process conditions explained

  low temperature gives good yield but slow rate so we use a compromise temperature

  high pressure gives a good yield and high rate but too high would lead to high energy costs for pumps to produce the pressure
• production of methanol from CO

  CO + 2H2 ⇌ CH3OH

  reaction is exothermic
• production of methanol from CO conditions

  T= 400 degrees celsius

  P= 50atm

  catalyst= chromium and zinc oxides
• production of methanol from CO conditions explained

  low temperatures gives good yield but slow rate so a compromise temperature is used

  high pressure gives good yield and good rate, too high pressures lead to high energy costs for pumps to produce pressure
• contact process

  stage 1: S + O2 ⇌ SO2

  stage 2: SO2 + 1/2O2 ⇌ SO3

  exothermic reaction
• contact process conditions

  T= 450 degrees celsius

  P= 1 or 2 atm

  Catalyst = V2O5
• Contact process conditions explained

  low temperature gives good yield but slow rate so compromise moderate temperature is used

  high pressure only gives slightly better yield and high rate: too high a pressure would lead to too high energy costs for pumps to produce pressure
• hydration of ethene to produce ethanol

  CH2=CH2 + H2O ⇌ CH2CH2OH

  exothermic reaction
• hydration of ethene to produce ethanol conditions

  T= 300 degrees Celsius

  P= 70 atm

  Catalyst= concentrated H3PO4
• hydration of ethene to produce ethanol conditions explained

  low temperature gives good yield but slow rate so compromise temperature is used

  high pressure gives good yield and good rate but too high leads to high energy costs for pumps to produce pressure and unwanted polymerisation of ethene to poly(ethene)
• recycling unreacted reactants back into the reactor can improve the overall yields of all these processes
• carbon neutral refers to an activity that has no net annual carbon (greenhouse gas) emissions to atmosphere
• methanol and ethanol can be used as fuels

  if carbon monoxide used in the production was extracted from atmosphere and energy required was not made by combustion of fossil fuels, it can be classed as carbon neutral
• Kc= equilibrium constant

  aA + bB ⇌ cC + dD

  [C]c [D]d / [A]a [B]b
• unit of Kc (concentration mol dm-3) depends on the equation
• moles of reactant at equilibrium = initial moles - moles reacted
• moles of product at equilibrium = initial moles + moles formed
• The larger the Kc the greater the amount of products
• If Kc is small we say the equilibrium favours the reactants
• Kc only changes with temperature

  changes in pressure, concentration or the presence of a catalyst has no effect