Collision Theory & Reaction Mechanisms

1. Collision Theory & Reaction Mechanisms How molecules actually react/interact with one another in a chemical reaction is explained through Collision Theory.

2. Collision Theory • Particles are in constant motion • A chemical reaction involves the effective collisions of particles ie. Sufficient energy with correct orientation (alignment of molecules so bonds can be broken or formed) • Rate of reaction depends on frequency of collisions and number of successful collisions

3. Successful vs. Unsuccessful Collisions

4. Rate of Reaction Factors & Collision Theory • Increasing concentration of chemicals and their surface area for reaction increases frequency of collision. • Catalysts and the nature of the reactants alter the number of successful collisions. • Increasing temperature changes both frequency of collisions and number of successful collisions.

5. Activation Energy In order for reactions to occur, molecules must collide with the right geometry (in order for the nuclei to bond) and with the right force of collision (activation energy). Activation Energy is this “right amount of energy” needed for reaction to occur. A slow reaction has a high activation energy while a fast reaction has a low activation energy.

6. Graphs of Ea & DH

7. An activated complex is the structure at the point in a chemical reaction where the reactants are transforming into the products.

8. Reaction Mechanisms Chemical reactions are not usually carried out in a single step process, as seen in a balanced chemical equation. Eg. 2C8H18(l) + 25O2(g) 16CO2(g) + 18H2O(g) • This implies that 27 molecules have to collide with the right geometry and energy to react! In fact, many simple steps must be completed before the overall reaction can proceed. This series of steps (called elementary steps) in order for a reaction to occur is known as the reaction mechanism.

9. For example, nitrous oxide (laughing gas) decomposes into nitrogen and oxygen gases according to the balanced chemical equation: 2N2O(g)  2N2(g) + O2(g) This implies that 2 molecules of N2O must collide with one another for reaction to occur (suggests 2nd order). However, experimental data indicates that this reaction is 1st order in N2O (dependent on only 1 molecule of N2O). This means that the slowest step in the series of possible steps must be the decomposition of just 1 N2O molecule.

10. Proposed Mechanism for this reaction consists of 2 elementary steps: 1. N2O(g)  N2(g) + O 2. N2O(g) + O  N2(g) + O2(g) _____________________________________________ Overall reaction: 2N2O(g)  2N2(g) + O2(g) • Step 1 comes from the rate law expression: rate = k[N2O] and must be the rate determining step (slowest step) • Step 1 is a unimolecular step; step 2 is a bimolecular step and “O” is known as a reactive intermediate

11. Eg. Predict a reaction mechanism for the reaction: X + 2Y + 2Z  XY2Z2 Given the rate law expression: rate =k[Y]2[Z]

12. 2Y + Z  Y2Z • Y2Z + X  XY2Z • XY2Z + Z  XY2Z2 Y2Z and XY2Z are known as: Reactive intermediates