Reaction rates & equilibrium

• Collision theory --> Particles must collide with sufficient activation energy
• To increase rate of reaction:

  • increase frequency of collisions
  • increase frequency of successful collisions
• Increase rates or reaction:

  • increase temperature
  • increase concentration
  • increase pressure
  • increase surface area
  • Use a catalyst
• Rate = concentration / time taken
• • The rate of reaction is fastest at the start of the reaction, as each reactant is at its highest concentration
  • The rate of reaction slows down as the reaction proceeds, because the reactants are being used up and their concentrations decreases
  • Once one of the reactants has been completely used up, the concentrations stop changing and the rate of reaction is zero
• Catalyst --> speeds up a reaction but doesn't get used up
• • A catalyst provides an alternative route with a lower activation energy
• Importance of catalysts:

  • Catalysts reduce energy demand
  • Increased sustainability
  • Less emissions
• Homogeneous --> in the same phase (state)
• Experiments to measure rates of reaction:

  • Measure cloudiness - disappearing cross
  • volume of gas - gas syringe
  • Mass loss - balance
  • pH - pH meter
• Heterogeneous catalysts --> Has a different physical state from the reactants. Heterogeneous catalysts are usually solids in contact with gaseous reactants or reactants in solution.
• • Reactant molecules are adsorbed (weakly bonded) onto the surface of the catalyst, where the reaction takes place. After reaction, the product molecules leave the surface of the catalyst by desorption.
• Features of Boltzmann distribution:

  • No molecules have zero energy - the curve starts at the origin.
  • The area under the curve is equal to the total number of molecules.
  • There is no maximum energy for a molecule - the curve does not meet the x-axis at the high energy. The curve would need to reach infinite energy to meet the x-axis
• At higher temperatures:

  • More molecules have an energy greater than or equal to the activation energy
  • Therefore a greater proportion of collisions will lead to a reaction, increasing the rate of reaction
  • Collisions will also be more frequent as the molecules are moving faster, but the increased energy of the molecules is much more important than the increased frequency of collisions
• Dynamic equilibrium:

  • Closed system - no material can escape/ get in
  • The forward and reverse reactions occur at the same rate
  • The concentrations remain constant
• Le Chatelier's principle:

  • If a change is made to a system in dynamic equilibrium the equilibrium opposes the change
• Change in temperature --> Raising the temperature favours the endothermic reaction. In the Haber process, the equilibrium will shift to the left.
• Change in pressure --> Raising the pressure favours the side with the least moles of gas. In the Haber process the equilibrium will shift to the right.
• Using a catalyst --> No effect on position. Equilibrium is reached faster.
• Change in concentration --> The equilibrium responds by reducing the concentration of the introduced reactant until a new equilibrium is established.
• The equilibrium constant:

  • Kc = [products] / [reactants]
  • High value of Kc means high concentrations of products
• Effect on temperature of Kc --> Kc would decrease with an increase in temperature. Kc only changes with temperature