Kinetics mainly works off of collision theory. This is the theory that, for a reaction to occur, the particles

must collide with Energy (E) ≥ the Activation Energy (Ea), and in the correct orientation to one another in

order to trigger a reaction. If either one of these conditions isn’t met, the reaction doesn’t occur and the

collision is not successful.

The activation energy (Ea) is the minimum energy required to trigger a reaction when 2 reactant particles

collide with one another.

Changing conditions and the effect on rate;

Increasing temperature increases the rate. Reactant particles have more Ek, so move faster &

therefore collide more often. As particles have more Ek, these collisions have E ≥ Ea more often,

triggering more frequent successful collisions

Increasing pressure increases rate. More particles in the same volume, so more frequent collisions

between reactant particles, and therefore more frequent successful collisions triggering a reaction

Increasing the surface area of a solid increases rate. Less particles ‘locked away’ in the centre of the

substance, so more solid particles available to collide, so more frequent successful collisions triggering a reaction

Catalysts increases the rate by providing an alternative chemical pathway with a lower Ea, increasing the number of successful collisions which trigger a reaction. Catalysts are not changed or used

up in the reaction

Maxwell-Boltzmann Distributions;

This shows the relative proportions of particulate energy. There are a few key features;

Note, during changing conditions, the area under the graph must remain constant as this represents the

total number of particles

X axis = particulate energy

Y axis = Number of particles

The graph should start at the origin as no particles have 0 energy

Peak = most common particulate energy

Mean particulate energy lies just to the right of the peak

Asymptote at high energy, as a few particles have a very high particulate energy

Activation energy is shown by a vertical line near the asymptote. All particles to the right of this line

have the Ea required to react in a successful collision

Changing Conditions;

Temperature can affect the distribution of the particulate energy in the diagram;

Increasing Temperature;

Peak of curve drops down and to the right. Crosses original curve and then stays ABOVE it. Greater proportion of particles have E ≥ Ea, so more frequent successful collisions and the rate increases

Decreasing Temperature;

Peak of curve moves up and to the left. Crosses original curve and then stays BELOW it. Smaller proportion

of particles have E ≥ Ea, so less frequent successful collisions and the rate decreases

Effect of Catalyst;

A catalyst does not actually effect the shape of the distribution; it just lowers the activation energy (new

vertical line to the left of the original). This results in a greater proportion of particles having E ≥ Ea and so

the rate of reaction increases

Note,

Rate of Reaction;

The rate is described as the change in concentration per unit time. It is usually measured in mol/dm3

/s. However, if

we cannot measure the concentration of reactants (for example if a precipitate is produced) we use 1/time as the

rate. To calculate the rate from a graph, draw a tangent and calculate the gradient at the specified time (initial rate is

when t = 0)

Rate Equation;

The rate equation can only be determined by experiments. Only the species that effect the rate of a

reaction appear in a rate equation. Any species in a rate equation which is not in the overall equation (e.g.

H

+

) is a catalyst. If we look at a reaction;

A(aq) + B(aq) -> C

We can deduce the rate equation in the general form ‘rate = k[A]x

[B]y

’ Where the [ ] represent the

concentrations of the reactants, the x & y represent the orders of these reactants in the rate equation (the

power to which these reactants appear in the rate equation) and k represents the rate constant which is

calculated by experiments.

You calculate the orders of reactants by varying the concentrations of each reactant. Usually, [A] will change whilst

[B] is stationary and then you have to use the order of A and the change in rate to deduce the order of B. Each

individual reactant can have one of three orders;

0

th Order -> change in concentration doesn’t affect rate

1

st Order -> change in concentration is directly proportional to change in rate

2

nd Order -> change in concentration is the change in rate but squared

If you add up the all the orders you can find the overall order of the reaction. Note, any substances of 0th order

doesn’t appear in the rate equation. To find the units of k rearrange to make k the subject and simplify the units.

Once we have measured the change in concentration (aqueous substances) of a reactant, or the change in volume of

a gaseous product at set intervals (every 10 seconds) we can plot a graph of concentration against time. Then

calculating the gradient (rate) at set points we can plot a graph of rate against concentration. From these graphs we

can find out the order of reaction;

Half-life in this sense refers to the time taken for the concentration of a given reactant to half. We can also calculate

the rate at this point by drawing a tangent and calculating the gradient.

Rate Determining Step (RDS);

This refers to the slowest step in a reactions mechanism, and is the step that governs the overall rate of reaction. The RDS only contains species which appear in the rate equation, and these species must appear

the same number of times in the RDS as their order in the rate equation

Rate = k[NO2][F2] This equation can have 2 possible RDS;

NO2 + F2 -> NO2F + F RDS

NO2 + F -> NO2F Fast reaction

The first equation contains the species in the rate equation the number of times of their order, so this is the RDS

Arrhenius Equation;

This equation is used to relate the rate constant (k) to the temperature of a reaction. By natural logging both sides,

we can plot a graph of lnk against 1/t. The gradient is –Ea/R but as R is constant, multiplying the gradient of the line

by –R will give the value of Ea in joules. The y axis intercept is lnA.

You can also use e^( –Ea/RT) to calculate the number of successful collisions. This will give the number of successful

collisions as a small number, which you should round off to a whole number and write as a fraction (e.g. 10/75000

successful collisions)