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What is collision theory?

Collision theory, as pointed out by wh00t, is a theory used by chemists to predict the rate at which chemical reactions will occur (or the effect that different conditions will bring upon the rate for non gas phase reactions). Collision theory works by assuming that all substances are made up of distinct particles (this agrees with current theories - these particles are, generally speaking, molecules). These particles are all moving around in some way (in a gas, they rush around randomly, in a liquid, they move around in a random fashion, but slower, and in a solid they are merely allowed to vibrate). If you mix two chemicals, the particles continue to move around as normal but now there is a chance that they may hit each other, eg if we have O2 and H2 gasses, mixed together, we may have something like this:

  O2
   \|           <-- H2

       H2 -->
                  H2         _
        _         |         /|
        /|       \|/       O2
       H2

Where the arrows indicate the direction in which a particular molecule is travelling.

What happens when they collide?

At some point it is likely that one of the O2 molecules will collide with a H2 molecule, and if it does, it's possible that a reaction will occur. There are a few conditions on the possibility of a reaction:

Activation energy

Every reaction has a particular activation energy -- this is the minimum total energy that the colliding particles need to cause a reaction to happen - the energy coming from their speed (kinetic energy) or some form of internal energy. It's fairly easy to understand this if you think about it - there's unlikely to be any effect if two low energy particles hit each other in the same way there's very little effect if you slowly lower ha hammer onto a nail.

The activation energy is generally the energy needed to break *all* of the bonds that need to break to react the two chemicals, so in the case of a reaction between O2 and H2...

O = O
H - H

There is a O to O double bond (there are two chemical bonds between the two O atoms) and a H to H single bond to be broken in order to react the two together. The product of the react will be H2O and a spare O molecule, so the reaction is best written with 2 H2 molecules so there's no left over stuff:

2 H2 + O2 --> 2 H2O

And the activation energy for the reaction is the energy to break the double O=O bond and 2 H-H bonds. The activation energy for some reactions can be lowered by introduction of a catalyst - this provides an alternate 'route' for the reaction to take place by breaking it down into smaller steps.

Orientation

In some chemical reactions, the orientation of the molecules involved may have a significant effect on the reaction, for instance, if we try to add an OH- ion to a bromomethane molecule...

   ______   H                H
  /      \| |                |
OH- + Br -- C -- H  ->  H -- C -- OH + Br-
      |\_/  |                |
            H                H

The OH- ion can only attach the bromomethane from one side or no reaction takes place, therefore since for the most part, the collisions will probably be in the wrong orientation, there is a reduced rate of reaction.

Other stuff collision theory shows

As well as giving some insight into the mechanism of a chemical reaction, collision theory also yields some other important factors in the rate of a reaction...

Pressure

If a reaction occurs in the gasseous phase, you can increase the rate of reaction by increasing the pressure -- this is neatly explained because by increasing the pressure, you are increasing the number of particles that pass through a given volume of gas in a given time (if you take a small cube and look in it, with increased pressure, there will be more particles going through it). Since there are more particles going through the space in a given time there is more chance of a collision between them.

Concentration

If the reaction occurs inside a liquid, the concentration of reactants works in a similar way to pressure in a gas - more concentration means more particles in a given volume and therefore more chance of a collision.

Surface area

If one of the reactants is a solid, surface area comes into play - with less surface area the 'target area' for the reacting particles to hit is smaller so there is less chance of a collision. In addition, there is no way the reactants can get at the matter in the middle of the solid so the rate is slowed because there are effectively less reactants. With more surface area (eg breaking up into a powder) the rate is faster because there is a bigger 'target area' for reactants and more of the solid reactant is available to collide with.

Temperature

At higher temperatures, the particles have more kinetic energy ie are moving faster, therefore there are more particles moving through a given space in a given time, and so more collisions and a faster rate of reaction.

In addition, with a higher temperature, because the particles have a higher kinetic energy, any collisions that occur are more likely to cause a reaction as the energy of the particles is more likely to be above the activation energy. The energy of particles actually matches the Maxwell-Boltzmann distribution, and looks something like this (A is a low temperature, B is a higher temperature):

   /|\    |
    |     |
  Number  |        ___ A
    of    |      _-   -__         x = particles in high
Molecules |     -        --_         temp with >
          |    |            \        activation energy
          |   /              \    X = particles in low
          |   |       __B     |      temp with >
          |  /     _--  -__   \      activation energy
          |  |   _/        --_ |
          |  |  /             -+_         Activation
          | /  /               \ \        | Energy
          | | /                 \ --___B  
          | |/                   \     ---+____
          |/ |                    \        xxxx----_____
          ||/                      -\_A   |xxxxxxxxxxxxx
          |/                          ---__xxxxxxxxxxxxx
          ||                              |XXXX----xxxxx
          |---------------------------------------------
                                      Kinetic Energy--->

As you can see from the graph, there are more particles with enough energy to react in a hotter environment (B) than in a colder one (A) - and so there are more successful collisons, and a faster rate.

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