3.2 Physical Chemistry

• Activation Energy?

  minimum energy required to start a reaction by the breaking of bonds
• Exothermic Reactions:

  heat is given out to the surroundings, so temperature of surroundings rises therefore system loses energy so enthalpy is negative

  
• Endothermic Reactions:

  heat is taken in from surroundings so temperature of surroundings fall, but the system gains temperature, therefore enthalpy is positive

  
• Standard Conditions:

  temperature - 298K

  pressure - 100kPa or 1atm

  concentration - 1moldm-3
• Standard States:

  state of substance in standard conditions

  water is liquid

  hydrogen is gas

  sodium is solid
• Enthalpy of reaction?


  enthalpy change that accompanies reaction in the molar quantities expressed in a chemical equation under standard conditions, all reactants and products in their standard states
• Enthalpy of formation?


  the enthalpy change when one mole of a compound is produced from its elements in their standard states all under standard conditions
• Enthalpy of combustion?


  the enthalpy change when one mole of a substance is completely burned up in oxygen under standard conditions and all reactants and products in their standard states
• Enthalpy of neutralisation?

  enthalpy change when one mole of water is formed from a neutralisation reaction between an acid and a base, all reactants and products in their standard states
• enthalpy change = energy/number of moles
• energy (J)= mass*shc*temperature change
• enthalpy sign?

  temp rise means negative enthalpy

  temp fall means positive enthalpy
• Average bond enthalpy?


  the average energy required to break one mole of a given bond in the molecules in the gaseous state
• Endothermic?


  when more energy is needed to break the bonds then the energy released making bonds
• Exothermic?


  when less energy is needed to break bonds than is released making bonds
• Enthalpy change = sum of bonds broken - sum of bonds formed
• Hess's Law:

  states that if a reaction can take place by more than one route the enthalpy change of a reaction is independent of the route taken, provided the conditions at the start and end are the same
• Hess's Application

  
• Reasons for using indirect route:

  -activation energy too high

  -rate is too slow
• Effect of concentration/pressure on rate:

  increasing it increases the rate of reaction, as their is more particles per unit volume so they are more likely to collide with each other and there are more collisions per unit time
• Initial rate calculation?


  tangent at t=0 and find the gradient
• Rate at specific time calculation?


  tangent at t=time and gradient
• Catalyst?


  -substance which speeds up a reaction without being consumed

  -it provides an alternative route with a lower activation energy
• Homogeneous catalyst?


  same physical state as reactants
• Heterogeneous catalyst?


  different physical state from reactants
• Economic importance of catalysts?


  -lower temperatures and pressures are required

  -increased atom economy so fewer pollutants

  -benefits to the environment of improved sustainability weighed against toxicity of some catalysts
• Boltzman distribution graphs features:

  -no particles have 0 energy, so graph starts at (0,0)

  -area under the curve = number of particles
• Boltzman temperature change:

  -higher temperature, faster reaction

  -higher temperatures shifts graph to the right, this shows more particles have energy above activation energy
• Boltzman catalyst change:

  -increases rate of reaction

  -alternative route with lower activation energy

  -activation energy on graph moves left
• Techniques to investigate rate of reaction:

  -loss in mass, must be done on a balance

  -production of gas, gas syringe or measuring cylinder

  -change in colour or precipitate forms
• Features of a Dynamic equilibrium?


  -rate of forward reaction=rate of reverse reaction

  -concentration of reactants and products are the same unless conditions changed

  -only occurs in a closed system
• Le chateliers principle?


  if a system is subjected to a change it will move in the direction which opposes that change
• Concentration change?


  -concentration of a reactant is increased, the equilibrium will shift to the right to oppose it (forward)

  -concentration of a reactant is decreased, the equilibrium will shift to the left to oppose it (backward)
• Pressure change?


  -total pressure is increased, the equilibrium moves in the direction that has fewer moles of gas

  -total pressure is decreased, the equilibrium moves in the direction that has more moles of gas
• Temperature change?


  -temperature is increased the equilibrium will move in the endothermic direction(+)

  -temperature is decreased, the equilibrium will move in the exothermic direction(-)
• Catalyst, le chateliers?


  -increases the rate of forward and reverse by the same amount

  -no effect on the equilibrium position
• Haber Process:

  -N2 + 3H2 --> 2NH3

  -pressure: 200 atm, increases yield and rate, leads to energy costs that are increased and safety issues

  -temperature: 400-450'c , high temp increases rate, lower yield, energy costs are increased and environmental issues

  -catalyst: iron, increases rate and no effect on equilibrium
• Kc?

  $Kc=$$[products]/[reactants]$

  -if multiple moles, put it to the power of the compounds concentration
• Large Kc?


  equilibrium is to the right
• Small Kc?


  equilibrium is to the left