Collision Theory

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Background and General Aspects

• Chemists propose that for a reaction to occur, reacting particles (atoms, molecules, or ions) must collide with one another, this is known as collision theory
• This proposition is supported by the kinetic molecular theory of gases. For example, it is estimated that in a 1 mL sample of gas at STP, approximately 1028 collisions among the molecules take place every second.
• If each collision resulted in a reaction, all the reactions would be complete in about a nanosecond (10 -9s).
• In fact, these reactions occur much more slowly. Thus, it is reasonable to infer that only a small fraction of collisions between reactants results in a reaction.

Effective Collisions

• for a collision to result, it must be effective-which means that it must satisfy the specific conditions:

1. the orientation of the reactants must be favourable[pic 3]

• reacting particles must collide with the proper orientation relative to one another. This is known as having the correct collision geometry.
• the example below shows three of the many possible ways in which NO (g) and   O3(g) can collide. Only one of the three possibilities has the correct collision geometry for the reaction to occur.
• As shown in the diagram, only a specific orientation of the two reactants leads to a reaction.

1. The collision must occur with sufficient energy  

• Reactant particles must collide with sufficient energy for a reaction to take place. This is called activation energy (Ea).
• The collision energy depends on the kinetic energy of the colliding particles.

• [pic 4]The dotted line in the diagram indicates the activation energy and the shaded part indicates the fraction of the total collision’s that have energy greater than or equal to Ea.
• Suppose that a set of reactants collide 10,000 collisions/s and 1/50 collisions are effective. The reaction rate would be 10,000 x 1 reaction/50 collisions, or 200 collisions/s.
• Ea does not change when the temperature changes.
• However, if the rate would change if the number of total collisions/s change.

Transition state theory

• As a reaction, the potential energy of the reactants increases as they get closer to each other.
• If the collisions does not meet the proper conditions to an effective collision will change in configuration and are then said to be in their transition state.
• From the transition state, the process can go forward to form products or go backward and re-form reactants.

Potential energy diagrams

• Potential energy diagrams are used to represent the changes of potential energy as a chemical reaction progresses.
• It shows the activation energy and enthalpy changes.

Exothermic reaction

• in the diagram below, the products have a lower energy than the reactants have. This decrease in potential energy indicates that energy was released in the reaction

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Endothermic reaction

• In the diagram below, the products have more potential energy, indicating that energy was absorbed. Thus the reaction was endothermic.

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Activation Energy and Enthalpy

• the activation energy of a reaction cannot be predicted from its enthalpy change.
• The enthalpy change is determined only by calculating the difference in the potential energy of the reactants and products.it is independent of any process that occurs during the reaction.
• The activation energy of a reaction is determined by analyzing the reaction rate at a various temperature.
• In general, reactions, regardless of whether they are exothermic or endothermic, with low activation energies tend to proceed quickly at room temperature while reactions with high activation energies tend to proceed slowly in room temperature.
• For instance, gasoline, a highly flammable substance, does not burst spontaneously into flames because it requires a spark, a small energy input, to overcome Ea and initiate combustion. Once ignited, the gasoline continues to burn because the energy released from the initial reaction exceeds the activation energy of the other gasoline molecules, creating a chain reaction until the gasoline molecules are exhausted.

Activation Energy for reversible reactions

• many reactions can proceed forwards and reverse.
• For example: forward: CO(g) + NO2(g)- NO(g)         ∆Hr=226 kj

                      Reverse: CO2(g) + NO(g)-CO(g) + NO2(g)     ∆Hr=+226 kj

• a potential energy can represent the reaction in both the forward and reverse directions.
• The relationship between the forward Ea, reverse Ea, and the enthalpy change can be represented by the equation: Ea (fwd)- Ea (rev) = ∆ Hr
• For exothermic reactions Ea (fwd) is less than Ea (rev), and for endothermic reactions, Ea (fwd) is more than Ea (rev)
• During the transition state, there is a chemical species that exist referred to as an activated complex.
• An activated complex is neither a product nor reactant. It is a temporary arrangement of atoms that form as bonds are breaking and new bonds are forming.
• Because the activated complex contains partial bonds, it is highly unstable. It can break down to either form or re-form reactants.

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