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)