Kinetics Flashcards
collision theory
In order for a reaction to occur the reactant particles must collide with each other with the activation energy
rate of reaction definition and equation and units
how fast reactants turn into products
rate of reaction (moldm^-3s^-1) =(amount of reactant used or product formed)/time
activation energy definition
minimum amount of energy required by particles for a reaction to start
why might some collisions not lead to a reaction
particles have energy is not greater than the activation energy and they may be colliding in the wrong orientation
features of a Maxwell-Boltzmann distribution graph (7)
No particles with no energy
Area under line = total number of particles
Most probable (mode) energy - Emp
Mean energy is to the right of peak
Some have low energies due to colliding with walls of container and losing energy
Activation energy
No maximum energy - graph is asymptotic (tends to zero but never becomes zero)
what happens to a Maxwell Boltzmann distribution graph if concentration increases (2)
More particles so more with activation energy
Same overall shape but above the curve as there are more particles
what happens to a Maxwell Boltzmann distribution graph if temperature increases (2)
Shape is similar but is shifted to the right as there are more particles with energy >_ Ea
Total area under the curve is the same as no particles are added
factors affecting rate of reaction (3)
temperature
concentration/ pressure
catalyst
how does temperature affect rate of reaction
10C , double rate
increase, increase
Increases KE of particles so increases the frequency of collisions
Increases number of particles with energy Ea so higher frequency of successful collisions
how does concentration/pressure affect rate of reaction
double concentration/ pressure, double rate
increase, increase
More particles per unit volume so higher frequency of collisions
More particles to have energy >_ Ea
catalyst definition
A substance which increases the rate of reaction without being changed chemically at the end of the reaction by providing an alternative pathway for the reaction with a lower activation energy
(more particles will have the activation energy so more collisions are successful)
heterogenous catalyst: definition, example, problem
in a different phase (state) to the reactants
E.g. solid catalyst with gaseous or liquid reactants: Haber process uses an iron catalyst: N2(g) + 3H2(g) 2NH3(g)
Need replacing due to the build-up of impurities called poisoning
homogenous catalyst definition, example
• Homogeneous catalysts are in the same phase as the reactants
E.g liquid catalyst with liquid reactants: Dehydrate and alcohol uses H+ acid catalyst: C2H5OH(l) C2H4 + H2O
describe reaction over time
starts fast, slows down and stops
why does a reaction start fast
Concentration of reactants is highest so high frequency of collisions
