Collision TheoryCollision theory is used to explain how different variables affect the rate of a reaction. Its basic premise is that for a reactants to form into products, the reactant particles must collide however only a small percentage of these particles have the orientation and have enough kinetic energy to overcome the energy barrier to produce said products.In order for molecules to react, the colliding molecules must possess enough kinetic energy to overcome the bonds and repulsive energies of the reactants. The activation enthalpy is the minimum amount of kinetic energy required by a pair of colliding particles before the reaction will occur and is denoted Ea. The higher the Ea of a reaction, the smaller the amount of collision present and the slower the reaction as more energy is needed by the colliding molecules. However, the opposite is also true, the smaller the Ea of a reaction, the greater the amount of collisions present and the faster the reaction as less energy is needed by the colliding molecules. Effect of orientation of moleculesThe orientation of unsymmetrical molecules can determine whether a reaction will occur as without proper orientation, the reaction will not occur. An example of how orientation effects reactions can be seen below. The reaction can only occur if the hydrogen end of the H-Cl bond approaches the Carbon-Carbon double bound else they simply bounce off each other because of repulsion caused by the bonds.

Figure 5 - The Collision of Hydrogen Chloride with ethane to form chloroethane :(Carbon is shown in black, hydrogen in white, and chloride in green.)

From this we can conclude that an orientation favourable for breaking the bonds and allowing rearrangement is needed for a successful reaction.

Activation enthalpy and Enthalpy profilesAn enthalpy profile is when a plot is drawn of a reactions progress against the enthalpy (or energy) profile. They show the energy changes that take place during a reaction. The highest peak on the graph shows the point in the reaction where the old bonds are stretched and broken and new bonds have started to form and the activation enthalpy needed can be seen from the size of the peak. The curve on the graph below applies to a simple one step reaction however many reactions take place in a series of steps and drawing a single curve would be meaningless.

The difference in endothermic and exothermic reactions can be seen in their respective enthalpy profiles as the energy of the products is higher than the reactants towards the end of a reaction in an endothermic reaction due to heat being absorbed from the surrounding whereas the opposite is true for the exothermic reactions. As the iodine clock reaction is an exothermic reaction, its enthalpy profile would be similar to the exothermic enthalpy profile.

Effect of Temperature:

Temperature is a measure of the average kinetic energy of particles. Kinetic energy is the energy possessed by objects in motion and can be found using the equation:

As temperature increases, molecules move faster and collide more frequently because of the increased average kinetic energy. At any temperature the kinetic energy of particles in a system are spread over a wide range. The Boltzmann is the distribution function for the distribution of an amount of energy between identical but distinguishable particles.

The Boltzmann distribution curve can be seen in figure 6.(Figure 6)Only the particles represented in the area shaded green have a high enough energy to react where as those in the area shaded in blue in the graph will simply bounce apart elastically due to the repulsion caused by the charges in bonds. So to speed up the reaction there needs to be an increase in the number of particles possessing activation energy during collision to cause a reaction as they possess enough energy to break bonds and initiate the reaction. The Boltzmann distribution for a graph with the temperature T and the graph of an increased temperature labelled T + t is shown below:

From the graph you can see that by increasing the temperature that the number of molecules with an energy greater or equal to the activation energy has increased as the peak of the graph has moved towards the right, closer to the activation energy which means that more successful collisions will take place resulting in an increased reaction rate. Increasing the temperature may not have a difference on the individual particles within a system but there is a bigger difference to the proportion of particles that have enough energy to react.Kinetic Energy (Initial momentum before collision)

(final momentum after collision

L

Consider a single gas particle of the mass and travelling at a velocity in a two dimensional box of the length L. The initial momentum, , of the particle is Once the particle bounces off the wall and heads in the opposite direction the momentum is The change in momentum is therefore, .

Since force is the change in momentum per unit time, the force on one wall following the collision can be expressed as

The time between collisions, , can be found as the time taken to travel the given distance can be expressed as distance over the velocity. The distance is as the particle must travel to the wall and back thus

We can then substitute into our previous equation to get Suppose that a large number of molecules, , of the mass are moving independently in the box. Each particle will have a different velocity so the equation is now Now by calculating the mean value of the squares of all the velocities = Substituting this into the equation for the force we get

Pressure is defined as force over area: thus Now the area is and multiplied by we get which is the volume of the box, , so the equation becomes or However this only considers particles moving along one dimension so we now considering the problem in three dimensions.In two dimensions Pythagorean Theorem states that for a right angle triangle the length of the hypotenuse is:

YX

Z

Extending this into three dimensions we get: ZxCYBy

Combining the two equations we get:

Now as the particles move randomly it means that the mean values for the velocities X2, Y2 and Z2 are all equal so: Substituting this into our equation we now have the pressure of the gas in three dimensions:

is the number of particles which is equal to Avogadros number, , multiplied by the number of moles, .

For one mole, therefore: Which can be written as follows using the equation to work out Kinetic Energy :

Now using the Ideal Gas equation where is the molar gas constant (8.31J mol-1K-1): Therefore, Now which is the Boltzmann constant,, hence Where This equation allows to calculate the average kinetic energy of the particles in the reaction.

Effect of Concentration:

Concentration is the number of molecules in a given volume of a solution.In the iodine clock experiment, the concentration of a chemical is found using the following equation:

Concentration (moldm-3) = moles (mol) / total volume of solution (dm-3)

By altering the moles or the total volume of the solution the concentration of a chemical can be altered which can in turn have an effect on the rate of reaction.

A mole is a unit of measurement used in chemistry to express the amount of a chemical substance that contains the number of particles equivalent to the number of atoms in 12 grams of pure Carbon-12 (12C). The number of particles in 12 grams of 12C is 6.02214179 x1023 which is known as Avogadros number. So essentially one mole contains 6.02214179 x1023 so by altering the moles of a chemical we can increase or decrease the number of particles of that chemical within a given volume.

By increasing the concentration we can increase the rate of reactions as for a reaction to occur, two particles must collide with the minimum amount of kinetic energy needed to overcome the bonds and repulsion. Increasing the concentration means that there is a greater number of particles of that chemical in a given volume as seen in figure 5 which means that the frequency of collisions will increase. As the frequency of collisions will increase due to particles being more crowded together, the chance that a collision between two particles has the correct conditions - as discussed previously - needed for a reaction to occur will increase which can be seen in figure 6 leading to an increased the rate of reaction.

(Figure 5) (Figure 6)

In the case of most reactions, by increasing the concentration, the rate of reaction increase. However, in certain multi-step reactions where the reaction occurs in a series of small steps, increasing the concentration of one constituent may not increase the overall rate of the reaction. Take the following reaction:

The pace at which A splits into X and Y dictates the rate of as it is the slowest step in the reaction; this is known as the rate-determining step. If you increase the concentration of A, the frequency of the rate-determining step will increase due to the abundance of A and lead to an increase in overall rate. However, increasing the concentration of B will increase the frequency of the second step but makes little difference to the overall rate as for the second step to occur as it depends on the rate-determining step to occur first and the frequency of that has not changed.

The reason for the difference in the reaction speed for the multi-step reaction is due to the different activation enthalpies needed for the different steps. The rate-determining step requires colliding particles to have a higher total kinetic energy to overcome the energy barrier for the reaction to occur leading to a slower reaction rate; this can be seen in figure 7. This is why increasing the concentration of the chemicals in the rate-determining step has a greater effect on the overall rate. Whereas the second step in the reaction requires a smaller total kinetic energy between colliding particles so can occur much faster as seen in figure 8 but it still depends on the speed at which products are made in the rate determining step; the faster the products are made, the more frequently the second step can occur so increasing the concentration of the chemicals in the second step is less likely to affect the overall rate.

(Figure 11)

(Figure 12)CatalystCatalystWhen a catalyst is added to a reaction, the change in energy can also be seen in a enthalpy profile as the catalyst provides a different pathway for the reaction so that less energy is needed during the collision for products to be made. The enthalpy profile of an endothermic reaction with a catalyst can be seen below. After adding a catalyst to my reaction I will also draw an enthalpy profile to compare to my previous one.

Rate equation and Orders of Reactions