The Collision Theory and Factors Affecting Reaction Rates
Risa Thevakumaran
The Collision Theory
-
A reaction only occurs between two entities if they collide at the correct relative orientation with a
fixed minimum energy.
Effective Collisions
-
For a collision between reactant particles to be effective (result in a reaction):
1. The orientation of the reactants (the collision geometry) must be favorable
2. The collision must occur with sufficient energy 1
1) The Correction Orientation of Reactants
- Reactant particles must collide relatively to each other in the proper orientation (i.e. having the
correction COLLISION GEOMETRY)
Ex.1: NO(g) + O3(g) → NO2(g) + O2(g)
-
Below there are 3 possibilities in which nitrogen monoxide and ozone can collide and only ONE
will yield a reaction:
-
In our effective collision, the angle at which the nitrogen in nitrogen monoxide collides with
ozone is the same angle presently formed in nitrogen dioxide (the lower oxygen atom relative to
the central nitrogen).
2
1
2
RETRIEVED from Textbook, pg.365
http://chemwiki.ucdavis.edu/@api/deki/files/15977/14.23.jpg
2) Sufficient Activation Energy
-
Activation Energy, Ea: the minimum collision energy required for the reaction to occur; reactant
particles must collide with this energy in order to yield a fraction
Only a small fraction of total collisions have energy greater than/equal to Activation Energy.
3
-
-
Maxwell- Boltzmann Distribution: a defined curve that results when plotting the number of
collisions between particles in a substance (at a given temp.) against the KINETIC ENERGY of
each collision.
It important to mention that Activation energy is INDEPENDENT of temperature.
Representing the Progress of a Chemical Reaction – Potential Energy Diagrams
4
A) Exothermic Reaction: Activation energy is lower compared to that of reverse reaction, products
are at a stable, low energy state than reactants; Energy is released.
B) Endothermic Reaction: Activation energy is higher compared to that of reverse reaction,
products are at an unstable, high energy state than reactants; more energy is absorbed.
3
http://home.freeuk.com/braybrook/6-1/How_far_how_fast/Catalysis_files/image004.jpg
http://staff.prairiesouth.ca/~chemistry/chem30/graphics/2_graphics/exo.gif
5
http://staff.prairiesouth.ca/~chemistry/chem30/graphics/2_graphics/endo.gif
4
5
-
From the transition state (i.e. the activated complex) where the reactants have partially
converted to their products, the reaction may proceed to form products or reform reactants.
The enthalpy change can be calculated by subtracting the activation energy required for the
reverse reaction, Ea(rev) from the activation energy required for the original reaction, Ea(fwd):
ΔH = Ea(fwd) – Ea(rev)
Activation Energy and Enthalpy
-
-
You cannot predict the activation energy of a reaction from its enthalpy change; the enthalpy
change is independent of a reaction pathway and only depends on difference in potential
energy of the reactants and products.
However, the activation energy of a reaction can be determined by analyzing the reaction rate
at various temperatures (at room temperature, those with low Ea will proceed at a faster rate
while those with high Ea will proceed at a slower rate).
Factors Affecting Reaction Rate
-
Any factor that increases the frequency of collisions between particles also increases the
reaction rate and vice versa; there are 6 factors introduced:
1) Nature of the Reactants
- Reaction rate depends on the nature of the reactants (characteristics that influence the
chemical reaction) and the types of bonds breaking and forming.
- Reactions between ions in solution or between acids or bases have a faster rate due to the
presence of opposite charges and/or the ability to flow easily with less/no bonds.
- In molecular reactions, if the reactants are large or have strong bonds = reaction rate is slow
(Endothermic); if the reactants are very unstable = reaction rate is high (Exothermic; since ΔH is
higher than Ea it will provide enough Ea for the remaining reactants for effective collisions)
2) Concentration influences the Reaction Rate
- When reactants occur in solution, increasing its concentration leads to the greater number of
collisions per time due to the greater number of particles in the SAME volume = reaction rate
increases
- However as the reaction progress, the rate declines as the reactants are more likely to collide
with products.
6
6
http://chemhume.co.uk/ASCHEM/Unit%203/14%20Reaction%20rates/increased_concentration.jpg
3) Temperature
- When temperature increases, particles have more kinetic energy, increasing the frequency of
collisions and thus the number of effective collisions per time. At lower temperatures, the
kinetic energy is distributed less evenly, causing a small fraction of the collisions to be effective
while at higher temperatures, the kinetic energy is distributed more evenly, causing a bigger
fraction of the collisions to be effective.
7
4) Pressure
- If reactants are gases, increasing the pressure (by reducing the volume of a reaction container or
adding more reactant particles) increases the number of collisions/time, which in turn increases
the reaction rate.
5) Surface Area
- Smaller sizes of reactants have a greater exposed surface area compared to larger-sized
reactants with the exact total mass; thus a greater exposed surface area increases the frequency
of effective collisions and therefore, the reaction rate.
8
7
8
http://chem.wisc.edu/deptfiles/genchem/sstutorial/Text13/Tx133/tx133p1.GIF
http://www.ck12.org/book/CK-12-Chemistry-Second-Edition/section/18.4/
6) Catalyst
- A catalyst is a substance that increases the reaction rate by lowering the activation energy
required and that is not consumed/spent in the chemical reaction. Since more reactants than
have energy equal to/greater than the catalyzed activation energy, there are frequent effective
collisions and a faster reaction rate.
9
Homework: Read Section 6.2 and Take Notes
Extra Help (videos):
Learn about the collision theory and factors increase reaction rates through an analogy – Ted
Video10
If you don’t like analogies, click here11
9
http://www.dlt.ncssm.edu/tiger/diagrams/kinetics/Catalyst-2.gif
http://www.youtube.com/watch?v=OttRV5ykP7A
11
http://www.youtube.com/watch?v=syBoKVAng7E
10