measures the speed that reactants are consumed/ products are formed in a reaction
affected by concentration, pressure, temperature, catalysts & surface area
(/state of subdivision)
collision theory- for a reaction to occur: o reacting particles/atoms/molecules/ions must collide o collision energy must be greater or equal to activation energy o reacting particles must collide with suitable orientation
Potential Energy Profile
as particles approach each other, repulsive forces between electron clouds make them slow down and lose kinetic energy (which converts to potential energy)
if they collide they form a transition state/ activated complex o which is the highest potential energy state where new bonds are forming & old bonds are breaking o this is an unstable arrangement which decomposes quickly to form either original reactants or new products
activation energy: minimum collision energy required to form transition state & start reaction
Factors affecting reaction rate
Temperature

higher temp.  greater kinetic energy  more successful collisions  increased reaction rate
Concentration
higher concentration of reacting particles  more collisions  increased reaction rate
Pressure
higher pressure  greater concentration of reacting gas molecules  more collisions  increased reaction rate
Catalyst
provides reaction pathway with lower activation energy  more successful collisions  increased reaction rate
State of sub-division/ surface area
increased surface area  more particles exposed  more collisions  increased reaction rate
Reversible reactions and chemical equilibrium
many reactions are reversible
in a closed system, reversible reactions will reach equilibrium
two equal but opposing reactions proceeding at same rate
double arrows used to show reversible reactions
forward reaction and reverse reaction
at equilibrium, concentration of reactants and products remain constant
colour, pressure & temperature also remain constant
reactions are still occurring at equilibrium (dynamic)
Equilibrium composition: Reactants vs. products
relative concentrations of reactants compared to products is different for different systems
equilibrium constant K gives numerical value relating concentration of all species in a system at equilibrium aA + bB  cC + dD
K= [C] c [D] d
[A] a [B] b
K= equilibrium constant
only gases and aqueous species appear in K o solids and liquids have a fixed concentration so they are not included
K has a constant value for conditions of concentration and pressure but does change with temperature
Size of K shows composition of reaction mixture at equilibrium o Large value (above 10 4 ): equilibrium favours products o Small value (below 10 -4 ): equilibrium favours reactants o Value close to 1: significant concentrations of both reactants and products are present at equilibrium
Shifting the equilibrium position
system able to return to equilibrium after imposed changes: o concentration o total pressure  volume

o temperature
Le Chatelier’s Principle: if a system is at equilibrium and a change in conditions is imposed on the system then the system will re-establish a new equilibrium in such a way as to partially counteract the imposed change
Altering the concentration of one species
if the concentration of any species in an equilibrium system is changed then by Le Chatelier’s Principle a new equilibrium will form which partially counteracts the change
the system will consume some of the added reagent or replace some of the removed reagent
while these changes are happening the system is temporarily out of equilibrium & the rate of forward & reverse reactions wont be equal
when it returns to equilibrium the concentration/amounts of all reagents will have changed (to minimise original change)
increased conc. of products- equilibrium shifts left
decreased conc. of products- equilibrium shifts right
increased conc. of reactants- equilibrium shifts right
decreased conc. of reactants- equilibrium shifts left
Altering the pressure of an equilibrium system
equilibrium system involving gases must be in sealed container to achieve equilibrium
pressure can be changed by altering container volume
increased volume = decreased pressure
decreased volume = increased pressure
increasing pressure will favour side with fewer moles of gas
decreasing pressure favours side with greater moles of gas
if both sides have equal moles of gas then changing pressure has no effect on equilibrium position
adding inert gas has no effect on equilibrium position
Altering the temperature of an equilibrium system
only factor that can change value of K
if temp. increases then endothermic reaction will be favoured o if forward reaction is endo- K increases o if it is exo- K decreases
if temp. decreases then exothermic reaction will be favoured o if forward reaction is exo- K increases o if it is endo- K decreases
exothermic reactions have enthalpy written as
N
2
+ 3H
2
 2NH
3
+ 92kJ OR
N
2
+ 3H
2
 2NH
3
 H= -92kJ
endothermic:
C + H
2
O + 131kJ  CO + H
2
OR
C + H
2
O  CO + H
2
 H= + 131kJ
Other alterations
catalyst has no effect on equilibrium position but speeds up rate of attainment of equilibrium (will reach equilibrium sooner) o increases rate of both forward & reverse reactions equally o doesn’t change K

o changes activation energy for forward & reverse reaction
addition/removal of solids or liquids in the reaction doesn’t alter equilibrium concentration or amount of other species
adding water will not effect K but may affect equilibrium position
Summary o affects concentration of chemical species involved in equilibrium
Change Shift equilibrium Change equilibrium
Concentration
Pressure
Volume
Temperature
Yes
Yes
Yes
Yes
No constant (K)
No
No
No
Yes
No Catalyst