Introduction.

For molecules to break and form bonds a chemical reaction must take place. Collision Theory describes the principles that must be followed for a chemical reaction to take place and is the concept used to predict rates in chemical reactions. It states that for a chemical reaction to take place three principles must be followed:
• There must be frequent collision between particles
• Particles must have the correct orientation
• Particles must possess energy equal to the activation energy of the reaction to produce a chemical change.

Depending on the rate at which particles collide a rate of reaction is set, this is the speed at which a chemical reaction happens. Reactions can take thousands of years while some can happen in less than a second. However, the rate of reaction can be increased or decreased by adjusting different factors, including:
• Temperature
• Concentration
• Surface area of particle
• Pressure
• A present catalyst
These factors change the rate of reaction by adjusting the rate at which particles collide, the chance for particles to collide with correct orientation and can allow particles to react with a lesser amount of energy by lowering the activation energy required. For example, particles move by vibrating very quickly and they are able to vibrate faster in warmer temperatures, so by increasing the temperature the speed of particles are increased and subsequently, the rate of reaction increases. This effect on the rate of reaction can be observed on a Maxwell Boltzmann Distribution graph, shown below in figure 1.

The iodine clock experiment was discovered by Hans Heinrich Landolt in 1886, it is an example of a chemical clock experiment, a chemical reaction in which the ‘clock chemical’ starts with a low concentration but is followed by a rapid increase in concentration which can produce dramatic changes. In the iodine clock experiment, there isn’t an immediate reaction but after a short time, depending on the concentration of iodate, the liquid made a very swift change of colour, which is due to the formation of a triiodide-starch complex. The reason this is a clock experiment that takes a certain amount of time is because there are many reactions that happen after the initial one that forms an iodine.

1. Initial Reaction
Solution reacts to form an iodine and bisulphate solution
2H+ + 5HSO3- + 2IO3  I2 + 5HSO4 + H2O

2. Iodine reacts with reactants in original solution to form iodide
H2O + HSO3- + I2  2I- + HSO4- + 2H+

3. Iodide reacts to form iodine
2I-  I2

4. Iodide and iodine react to form triiodide
I2 + I-  I3-

5. Final reaction
I3- + starch  I3- + starch complex (produces the blue-black colour)

The Iodine Clock experiment implements collision theory by demonstrating the effect of concentration on the rate of the reaction. An increase in concentration means that there are more particles in the same space, as shown in figure 2, thus there should be a higher chance that particles collide frequently. By controlling all variables and changing the concentration of the iodate solution the experiment evidences that concentration can have a drastic effect on the rate of reaction.



Hypothesis.

If the concentration of the iodate solution decreases, the rate of reaction should proceed to decrease.

Materials.

Apparatus
- 3 x measuring cylinder (10ml)
- 2 x beaker (100ml)
- 3 x pipet
- 1 x conical flask (250ml)
- 1 x Stopwatch/Timer

Chemicals
- 36ml/experiment x Distilled Water (H2O)
- 54ml/experiment x Solution A (IO¬3¬+) (aq)
Made by adding 4.3g of KIO3 to 2L of H2O
- 90ml/experiment x Solution B (starch) (aq)
Made by adding 4.0g of soluble starch to 1L of H2O and 0.8g of NaHSO3 in another 1L of H2O. The sodium bisulphite is then acidified with 10ml of 1M H2SO4 and mixed.


Procedure.

1. 10ml of Starch Solution was measured at room temperature using a pipet to transport the solution from a beaker to the measuring cylinder. The contents of the measuring cylinder were poured into the 250ml flask that was on top of a drawn X.
2. Then 10ml of Iodate Solution was measured at room temperature using a different pipet from a beaker to another measuring cylinder.
3. The 10ml of Iodate Solution was poured into the flask containing the Starch Solution, after the 10ml was deposited, the timer was started. The mixture was not mixed.
4. After the mixture created a cohesive blue/black colour and the X was unable to be seen the timer was stopped, and the time was recorded.
5. Steps 1-4 were repeated but the 10ml of Iodate Solution was replaced with a 10ml mixture of 9:1 Iodate Solution and distilled water.
6. Steps 1-4 were repeated 9 times in total while decreasing the ratio of Iodate by 1 and increasing the ratio of water by 1, e.g. a ratio of 8:2 Iodate Solution and distilled water followed by a ratio of 7:3 Iodate Solution and water.
7. A complete set of data was taken upon completion of steps 1-4 using a 1:9 ratio of Iodate Solution and distilled water.
8. The entire experiment from steps 1-7 was repeated to get a 2nd set of data.
9. The results were graphed on a scatter with a line of best fit added after.





Results.

Conc. Iodate (mol/L) Time Elapsed Result 1 (secs) Time Elapsed Result 2 (secs) Time Elapsed Average (secs)
0.01 21.69 19.85 20.77
0.009 22.38 21.37 21.875
0.008 24.35 24.56 24.455
0.007 36.03 32.12 34.075
0.006 46.97 40.32 43.645
0.005 65.39 53.37 59.38
0.004 86.75 70.06 78.405
0.003 109.91 105.66 107.785
0.002 268.37 375.09 321.73
Table 1. The time elapsed for the reaction to take place for each variable of concentration

After the concentration was lowered it was observed that the rate at which the mixture changed into a cohesive blue colour slowed and there was an exponential decrease in the rate at which the lower concentration mixtures reacted.

Figure 3, there is an evident outlier in the 0.002 concentration mixture

It is also evident that the 0.002 concentration mixture is an outlier as when it is inserted into the graph the line of best fit is not very accurate. However, without the last point the line of best fit works very well and most points lie accurately and precisely on the line.

Discussion.

The data above in table 1 shows an observable difference between the first set of results and the second but comparing them in figure 3 and 4 show that the rate of reaction between the first and average graph are very similar. This means that despite a noticeable difference between the numbers when this information is graphed what matters is the rate of reaction, which should be similar across experiments.

During the experiment it was also observed that as the concentration decreased less of the mixture was able to react at the same rate, this is supported by the Maxwell Boltzmann Distribution graph shown in introduction as it graph shows that when the concentration increases more particles are able to pass the activation energy threshold compared to a lower concentration mixture.

The hypothesis of this experiment was correct, this is supported by the information above in figures 3 and 4 which show a clear decrease in reaction rate when the concentration is lowered. It was also expected due to the understanding of collision theory where, if the concentration of the IO¬3¬+ solution was decreased there would be less frequent successful collisions between particles, thus decreasing the reaction rate.

Possible sources of error in this experiment are the temperature of the solutions during the procedure, as the temperature could increase or decrease the rate of reaction. A control for these errors would be to let the solutions sit in room temperature for an hour with a thermometer before beginning the experiment to make sure that the temperature of the solutions will not make any furt