Group Two
Collision Theory
In the early part of this century, a theory explaining reaction rates was developed from
the kinetic-molecular theory of gases. The theory, known as collision theory, holds that
for a chemical reaction to occur the reacting molecules must collide with each other. If
the molecules do not collide, there will be no reaction. The theory also states that the rate
of the reaction depends on the number of collisions between the reacting molecules. This
explains the dependence of reaction rate on the concentration of the reactants. The more
concentrated the reactants, the closer together they are, and the more likely they are to
collide and react.
Collision theory also explains why a reaction rate decreases with time. As soon as the
reaction begins, some reactant molecules are consumed and thus their concentration
decreases. As the concentration of the reactants are reduced, the reaction slows down
because collisions occur less frequently.
Why Every Collision Does Not Lead to a Reaction
Simple collision theory implies that every collision leads to a reaction. However,
calculations indicate that in a 1L volume of gas at 25  C and 1 atm pressure (SATP
conditions), there are approximately 1 x 1030 collisions every second! If every one of
these collisions were to result in a reaction, then all gas-phase reactions would occur
simultaneously. A similar situation would exist for liquids. Since not all gas and liquid
phase reactions are instantaneous at SATP conditions, we must conclude that not all
collisions lead to a reaction.
For a reaction to occur, reacting particles must collide with each other effectively. An
effective collision is one in which the colliding particles:
approach each other at the proper angle (i.e. have the proper orientation)
possess enough energy
Figure 2
For many simple reactions, they way the particles are facing
does not affect the possibility of a reaction. For other reactions
involving more complex particles, the way the collisions occur
is very important. Figure 2 shows that in order for NO to react
with ozone (O3), the nitrogen atom, not the oxygen atom, must
collide with one of the oxygen atoms of the ozone molecule.
Even if the particles collide with the correct orientation, a reaction will only occur if the
particles involved have sufficient energy. When two particles collide, the energy of the
collision has to be sufficient to cause the old bonds to break and new ones to form
(Figure 3).
Figure 3
Thus, two factors – frequency of collisions and the effectiveness of collisions – determine
how much reactant changes into product in a given unit of time.
Demo:
Overview: In this activity students will simulate the collision theory focusing on the
orientation of molecules.
Safety/Concerns: Remind students that though molecules have varying energies, we do
not want to injure anyone. Therefore, take care in not creating a “mosh” pit in the
classroom.
Activity:
- Move desks in room to create a large empty space
- Divide students into two different types of molecules
- 1 Type of “molecule” holds their arms out in front of them, parallel with the floor
- The other “molecule hold their arms out “like a bird”
- Both sets of molecules are to walk around the room, simulating molecules in
motion.
- When molecules collide, successful new compounds will be formed when one
hand from each of the two separate molecules touch.
Sample Questions
1. Describe the collision theory of reaction.
2. Although methane, the main component of natural gas, burns readily in oxygen,
the reaction is so slow at room temperature that it is not detectable. Explain this
observation.
3. Give two reasons why most molecular collisions do not lead to a molecule
reaction.