Persulphate (S2O82-) is reduced to sulphate (SO42-) in presence of iodide ions (I-) in aqueous solution. While persulphate is being reduced, iodide is being oxidised to iodine. The overall reaction equation being:
S2O82- + 2I- â†’ 2SO42- + I2
On the other hand, iodine can be converted to iodide in presence of thiosulphate (S2O32-) as follows:
I2 + 2S2O32- â†’ S4O62- + 2I-
If iodine indicator is added to the solution a blue colour of the starch-based iodine indicator complex develops after all thiosulphate has been consumed. This is because the starch-based iodine indicator complex is very stable in high concentrations of iodine. The charges transfers and the energy level spacings in the resulting complex correspond to absorptions in the visible part of the spectrum – blue. The strength of the resulting blue colour depends on the amount of iodine present.
Therefore, the development of the blue colour is useful as it indicates a constant point in the progress of the reaction which can be taken as being representative of the “end” of the reaction (Practical Booklet p.37).
Several kinetic parameters can be easily analysed and determined if those reactions are timed as they progress from start to end, hence the importance of the component time (t) in kinetic studies.
By adding a known amount of thiosulphate and iodine indicator, the reaction rate can be determined as it obeys the following rate equation:
This experiment was undertaken to determine the reaction rate, the order of reaction with respect to both iodide and persulphate, and the effect of variation in ionic strength on the rate constant of this reaction.
EXPERIMENTAL – Methods & Procedures
An exact copy of the methods provided in the practical booklet is attached to this document as no deviations were made from those given.
Time t’ (min:sec)
Tim t’ (seconds)
Table – Timing of reaction from the exact time of mixing to the time of the blue colour appearance
Treatment of results
Solution B ( + 0.5g iodine indicator)
[KI] (mol cm-3)
[Na2S2O3] (mol cm-3)
[KCl] (mol cm-3)
[K2S2O8] (mol cm-3)
Table – Concentrations of all substances in the final reaction mixture and water volume to account for dilution to a 50cm-3 final volume.
From the rate equation of this reaction:
Interpreting the above results, doubling [I-] while keeping [S2O82-] constant doubles the rate of reaction – halves t’ (time of appearance of the blue colour). Doubling [S2O82-] while keeping [I-] constant, also doubles the rate of reaction.
The below equation gives the relationship of the variation of rate constant (k) with ionic strength (I):
This equation is of type , where k0 is the limiting value of the rate constant at zero ionic strength. where zA and zB are the charges of the two ions involved in the rate-determining step at low ionic strengths.
Since [I-] remains constant during each kinetic run, the integrated rate equation is:
where a and (a-x) represent [S2O82-] at times t = 0 and t = t’ respectively.
Making k the subject:
Hence, for reactions 1, 4 and 5:
Per 2 moles of S2O32- consumed 1 mole of S2O82- oxidises iodide because any iodine produced is converted back to iodide until all thiosulphate is consumed. Therefore, adding a known amount of thiosulphate allows to determine the concentration of persulphate at the “end” of the reaction, when t=t’.
From 0.0005 mol cm-3 of S2O32- consumed, 0.0005/2 = 0.00025 mol cm-3 of S2O82- were consumed. Therefore the remaining persulphate at the “end” of reaction = 0.0040 – 0.00025 = 0.00375 mol cm-3
For reaction 2:
For reaction 3:
At the start of this run [S2O82-] = 0.00800 mol cm-3 and [S2O32-] = 0.0005 mol cm-3. At the end of the reaction 0.0005/2 = 0.00025 mol cm-3 of S2O82- were consumed. Hence, from the initial concentration of S2O82-: 0.00800 – 0.00025 = 0.00775 mol cm-3 of S2O82- were present at t=t’.
Rate Constant – Determination
The value of k can now be calculated for each mixture. The values were calculated and recorded in the table below:
Rate Constant (k) / mol cm-3s-1
Table – Values of k constant for each reaction mixture
For reaction 1:
t’ = 1172 s
Ionic Strength – Determination
The contribution of each electrolyte to the total ionic strength is given by:
Taking into account the dilution of the final volume to 50 cm3 of all electrolytes, and knowing that for 1:1 electrolytes, I=concentration, and for 2:1 electrolytes, I= 3 x concentration. Hence, for these mixtures:
The ionic strength was calculated and recorded in the table below:
Ionic Strength (I) / mol cm-3
Table – Ionic Strength per each reaction
Calculation of B and k0
In order to determine B and k0 a plot of v.s. logk was plotted.
Table – k constant and ionic strength for each reaction
Graph – Correlation between logk and ionic strength
The correlation is given as being y = 2.6305x – 3.7229.
logK0= -3.72 mol cm-3 s-1
k0= mol cm-3 s-1
k0= mol dm-3 s-1
B = 2.63
âˆ†Hâ€¡ & âˆ†Sâ€¡ Calculation
As stated above:
“(..) doubling [I-] while keeping [S2O82-] constant doubles the rate of reaction – halves t’ (time of appearance of the blue colour). Doubling [S2O82-] while keeping [I-] constant, also doubles the rate of reaction. “
Hence, in the rate equation:
x= 1 and y=1. The reaction is of first order with respect to and with respect to . As a result, the overall rate of reaction is x+y=2.
The overall rate of reaction is of second order in which the rate determining step is:
S2O82- + I- à IS2O83-
This is consistent with the rate observed experimentally by comparing mixtures 1 to 2 and 1 to 3.
The value of B (2.63) is rather high in comparison with the theoretical value but of same magnitude.
As stated above, B is given by where zA and zB are the charges of the two ions involved in the rate-determining step. Hence, theoretically .
The value of k0 determined experimentally is rather low in comparison to the published value.
k0= mol dm-3 s-1
Published value: mol dm-3 s-1
Some source of error must have influenced this result. Perhaps wrong timing and contaminated glassware would allow a rather faster/slower reaction which results in a different rate constant (k).
(Answer to Q1 above.)
The ionic strength is the same for reactions 1, 2 and 3. Reaction 4 shows a characteristic ionic strength for the concentrations of ions present in the mixture (Table 2). Accordingly, reaction 5 with no KCl added shows a rather low ionic strength as expected.
A variation of Ionic strength vs. Rate constant is notable. With reference to table 5, a decrease in ionic strength accompanies a decrease in the rate constant – the lower the ionic strength of a reaction the lower its rate constant (k).
It is important to mention that the ionic strength of each mixture remains constant up to time t’ (appearance of blue colour). This is perfectly understandable because the rate of production of anions is equal to the rate of their expenditure, hence manifesting a constant ionic strength. However, when all thiosulphate has been consumed the ionic strength decreases because an interruption in their equilibrium causes the consumption of iodine to stop.
In other words:
When the reactions reaches t’:
S2O82- + 2I- â†’ 2SO42- + I2
I2 + 2S2O32- â†’ S4O62- + 2I-
Thus, at t’ there is an increase in the amount of I2 and a drastic decrease in the amount of I-. Therefore:
which results in a low ionic strength.
(Answer to Q2 above)
For this experiment, the value of calculated is rather low but positive. All systems tend to progress in a direction of increasing entropy and therefore an increase in the disorder of the system in question.
For this bimolecular reaction in solution at 25°C, the entropy of activation of the rate determining step is positive meaning the final starch-based iodine indicator complex is rather stable because increases in entropy correspond to irreversible changes in a system. This is because the amount of work the system can do is limited because most of the energy was wasted as heat; therefore this reaction is thermodynamically irreversible reaction.
(Answer to Q3 above)
A reaction which proceeds more slowly with increasing ionic strength is
Aims were accomplished.
All kinetic parameters were determined although some major sources of error were present during this experiment.
Possible sources of error may include difficult or inappropriate timing (non-digital watches were used). Difficulty to quickly detect the end point of the reaction (development of blue colour) as the rate of reaction to some of the mixtures was very slow; contamination of glassware could also contribute for such errors.
Nevertheless, the order of reaction with respect to both iodide and persulphate was determined precisely; and the effect of decreasing ionic strength is known to decrease the rate of reaction.
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