# Effect of Varying Concentrations of Hydrogen Peroxide and Potassium Iodate on the Rate of the Briggs-Rauscher Oscillating Reaction

The purpose of this investigation is to determine the effects of varying initial reactant concentrations on the overall rate of the reaction. This is then used to calculate the overall rate constant which was determined to be

$$0.576 \pm 0.012 \mathrm{l^2 mol^{-2} s^{-1}}$$

## The rate equation was found to be $$r = k[H_2O_2]^0[KIO_3]^2[C_3H_4O_4]^1$$

The Briggs-Rauscher reaction is a damped chemical oscillator that demonstrates a cycle of distinct colour changes from clear to blue-black. As the actual mechanisms for this reaction is far too complex, a simplified mechanism is considered. The rapid acceleration of this cycle is described by several independent mechanisms that ultimately vary the concentrations of iodine and its ions. The indicator, starch, binds these iodine and iodine ions to produce the vivid blue-black colour.
The overall reaction can be defined as follows:

eq.1

This can be separated into two different processes:

eq.2-3

When $\mathrm{[I^-]}$ is low, the radical mechanism in (3) consumes HIO slowly. As $\mathrm{[I^-]}$ increases, HIO is consumed faster. Any HIO that does not react by (3) is instead reduced by $\mathrm{H^+}$ as a part of the nonradical mechanism. When more HIO is used up by (3), excess iodide ions are formed, suppressing the radical mechanism. In summary, as (3) acts by one mechanism, (2) always opposes it and favour the alternate mechanism.

A more detailed analyses requires that the processes be broken up further into the radical and the nonradical reactions, both of which contains several intermediate steps.
When iodide ions are abundant, the nonradical reactions dominate. In the first step, iodide ions and iodate ions react slowly:

eq.4

Iodous acid is then further reduced:

eq.5

Hypoiodous acid is reduced by hydrogen peroxide:

eq.6

As (3) is so much faster than (2), $\mathrm{[I^-]}$ constantly decreases until the point where the radical mechanism (7 - 11) becomes dominant.

eq.7-11

Here, we see that the radical mechanism is autocatalytic and will therefore increase in rate. The colour changes can be described by the equations

eq.12-13

The amber colour arises from the $\mathrm{I_2}$ produced by (12). As HIO is converted to $\mathrm{I^-}$, the solution suddenly turns blue-black due to the iodines’ bonding with the starch indicator. This colour then fades as (13) consumes iodine faster than it is replenished. The radical mechanism then restarts this cycle. Since the only factors never regenerated are malonic acid and iodate ions, the reaction will continue until either one of these run out.

## Experimental Procedure

### Method

Preparation of the experiment involves three solutions of which solution B was prepared first. 4.3g of potassium iodate was weighted by difference and transferred to the 100ml conical flask. As sulphuric acid was only available in 2M the 1M solution was made by using 5ml acid and 5ml deionised water. The resultant 10ml 1M sulphuric acid was then added to the conical flask. To complete solution B, deionised water was used to make it up to 100ml. This was then set on the magnetic stirrer and heated to approximately 50 degrees to ensure that $\mathrm{KIO_3}$ dissolves.

Solution A was prepared by adding 30ml of 30% hydrogen peroxide and 70ml of deionised water into the 100ml graduated flask which was inverted a few times to ensure even mixture.

Solution C was prepared by first adding 0.1g soluble starch to approximately 90ml of boiling water which was stirred until cooled. Then 1.5g malonic acid and 0.4g manganese sulphate were added and dissolved. Deionised water was used to make the solution up to 100ml.

The reaction occurs when all three solutions were cooled to room temperature and moved to the fume cupboard. Solutions A and B were first added to the 500ml conical flask and stirred with the magnetic stirrer. Upon addition of solution C, the stopwatch was started and the time taken for each oscillation (blue to blue cycle) was recorded.

## Results

### Data

Note: Lack of data in some experiments were due to the reaction reaching steady state prematurely. Repeats of the 2x $\mathrm{KIO_3}$ variation were skewed due to difficulty in dissolving the reactants and were subsequently left out of the results.

Control: Normal Concentrations

2[Hydrogen Peroxide]

2[Potassium Iodate]

1.5[Potassium Iodate]

0.5[Potassium Iodate]

From the data above, the rate of the reaction is determined as the time taken per oscillation when gradient is zero; rate is constant.

## Conclusion

It is then possible to conclude that the overall reaction with respect to the concentration of hydrogen peroxide is zero; to potassium iodate is second and to malonic acid is first. This gives the overall rate equation

Substituting in the values from the baseline experiment

### Concentration of hydrogen peroxide:

30% of 30ml used = 9ml = 9 L
As density = $$\frac{mass}{volume}$$
$$1.45 = \frac{mass}{9}$$

Mass = 13.05g
Divide by the molecular mass:
$$\frac{13.05}{34.0147}$$ = 0.384 moles
Final volume = 300ml
Final concentration of hydrogen peroxide = 1.28 M

### Concentration of potassium iodate:

Mass used = 4.3g
= 0.02 moles
Divide by final volume:
$$\frac{0.02}{0.3}$$ = 0.067M

### Concentration of malonic acid:

Mass used = 1.5g
= 0.019 moles
Divide by final volume:
$$\frac{0.019}{0.3}$$ = 0.0625M

This yields the final rate equation

Where $$k$$ can then be calculated:

## Uncertainties

### Weighing by Difference

Uncertainty in the three-figure balance is $$\pm 0.001$$, this is divided by the mass measured then doubled to find the percentage uncertainty.

This gives a total percentage error of $$\pm 0.676%$$.

### Volume Measurements

Volumes were measured using class A flasks and measuring cylinders.

This gives a total percentage error of $$\pm 1.08%$$.

### Concentrations

Percentage errors in the concentrations of hydrogen peroxide and sulphuric acid are found to be 0.02% and 0.10% respectively. This totals 0.12%.

### Final Results

The values above adds to the final result of

## Evaluation

As the only measurement of importance is the duration of each oscillation, accuracy of the stopwatch reading affects the whole experiment. This is unfortunately undermined by many factors including reaction time lag, ambiguity in the colour changes and measurement frequency. During the 0.5 concentration of potassium iodate experiment, for example, the colour transitions initiated so quickly that the first six or so cycles were impractical to time. When varying the concentration of hydrogen peroxide, the cycles are so slow that it was difficult to establish the apex of each transition. Timing measurements could be improved by decreasing the human input. For example, a linear light source could be placed opposite a photoreceptor with the reaction flask in between. The receptor could be calibrated to trigger a stopwatch whenever a threshold value of light detected was met. Such a set up would be able to determine a constant apex in each cycle and measure at a wide range of frequencies while negating human reaction time lags.

Temperature had to be constant during the reaction. However, the magnetic stirrers on which the reaction flask sits are known to heat up when operating. This may act to speed up the reaction giving inaccurate results. To prevent this, a thin heat mat could have been placed between the stirrer and the reaction flask to block heat conduction while allowing magnetic field for the stirrer.
When varying the concentration of potassium iodate, it was quickly found that even when set on a stirrer hot plate, the reactants were difficult to fully dissolve. After around an hour, the solution will become slightly cloudy and make no further progress.

It is assumed that, for this investigation, results variation due to stirring speed is negligible.