Kinetics between Ferric and Iodide Ions

*Christian Paolo Asequia, **Ercille Mae Pacamo,

***Rene Susette Ann Pontillas, **** Mariza Silagan

Department of Chemical Engineering

Xavier University-Ateneo de Cagayan

Corrales Avenue, Cagayan de Oro, Philippines

Juliet Dalagan, PhD

Department of Chemistry

Xavier University-Ateneo de Cagayan

Corrales Avenue, Cagayan de Oro, Philippines

Abstract: The experiment aims to determine the order of the

reaction rate of the reactants in the reaction: 2Fe3+

+ 3I-

2Fe2+

+ I3-. Also, the experiment aims to determine its over-all

reaction and its overall reaction rate. lastly, It aims to

determine the effects of adding a concentration of another

reactant to the solution. The method used was the method of

initial rates, which varies the concentration of one of the

reactants while the other is being held constant. Through this,

the time it takes to convert the product to reactant was

determined and also its respective concentration. Through this,

the rate of the reaction was determined. The logarithm of the

varied concentration of one reactant was plotted with the

logarithm of its rate. The equation of the line then was

determined and the slope, which corresponds to the order of

the reaction, was obtained. For the ferric ion, the order was

found out to be one and for the iodide, it was found out to be

2. The over-all reaction order was then found out to be three.

Furthermore, through series of calculations, the reaction rate

constant average was determined and it was found out to be

54.31. Lastly, through calculations, it was found out that the

addition of another substance reduces the rate of reaction of

the solution. Errors contributed to the experiment especially

the errors committed by the experimenters.

Keywords: Rate Constant, Rate Law, Order of Reaction,

Reaction Rate

I. INTRODUCTION

In order to write the expression of the rate of the

reaction, it is necessary to experimentally determine the

relationship between the rate and the concentration of each

reactant. More often than not, the rate of the reaction will

increase if the concentration of the reactant will be increased.

Increasing the population of the reactant also increases the

likelihood of a successful collision. According to the collision

theory, an assumption is made that for a reaction to proceed,

molecules of the reactants must collide with sufficient energy

and proper orientation. [1]

The Rate Law states that for any general reaction

aA + bB cC + dD , the rate is proportional to [A]m[B]

n, that

is:

Rate = k [A]m[B]

n [Equation 1]

The expression is the general expression for the rate

law where k is the rate constant. The dependence of the rate of

the reaction on the concentration can often be expressed as a

direct proportionality in which the concentrations may appear

to be zero, first or second power. The power to which the

concentration of a substance appears in the rate law is the

order of reaction with respect to that substance. In the general

reaction, the order of the reaction is: m + n [2]

The clock reaction approach can be used for the

kinetic study of ferric and iodide ions. This is named after the

analogy between chemical changes in such reactions and an

alarm clock. In a clock reaction, chemical change becomes

visible only after the reaction has reached a certain extent.

There are three steps in each clock reaction. The first is a slow

formation of a chemical intermediate. The second step is a fast

consumption of the intermediate by the limiting reagent and

the last step takes place after the limiting reagent has been

consumed which then causes the color change. The three steps

can be easily presented through:

A + B T ; slow

T + L X ; fast

T + I S ; fast

; Where initial reaction mixture contains variables A,

B, L and I. l is the limiting reagent and I is the indicator. S

represents chemical species that sends out the signal for color

change.

The kinetics of the reaction transformed into a clock reaction

is easily investigated by the initial rate method. All that has to

Figure 1.Calculation of the average rate.

be measured is the time elapsed from the mixing of two

solutions to sudden color change. For the oxidation of iodide

by ferric ions, the reaction rate can be defined as:

[Equation 2]

The initial reaction rate can then be approximated by:

[Equation 3]

With ∆ [Fe3+

] being the change in the concentration

of ferric ions in the initial period of the reaction. If ∆t is the

time measured, then [Fe3+

] is the decrease in ferric ion

concentration from the moment of mixing to the moment of

complete thiosulfate (limiting reagent) consumption. The

chemical reaction is then defined by:

( )

( ) ( )

( ) [Equation 4]

Therefore from the stoichiometric reactions, it follows that:

-∆ [Fe3+] = [S2O3

2-] [Equation 5]

And consequently,

vo = k[Fe3+

]ok [I

-]

yo [Equation 6]

From the latter equation, it will then follow that:

[Equation 7]

If the initial concentration of only one reactant is

varied while the initial concentrations of the others are held

constant, it is possible to determine the order of the reaction

with respect to the reaction of the reactant with varied

concentration [3]

The reaction to be studied in the experiment is the

oxidation of I- by the Fe

3+ ions according to the equation:

2Fe3+

+ 3I- 2Fe

2+ + I3

- [Equation 8]

II. EXPERIMENTAL SECTION

The experiment was divided into three parts. The first

part is the reaction order with respect to Fe3+

. A total of ten

(10) mixtures were prepared. Five of which is with varying

concentrations of Fe3+

and the other five is where the

Potassium iodide concentration was kept constant. Contents of

the mixture were swirled and were allowed to reach

equilibrium for 10 to 15minutes. The initial temperature of

each mixture was recorded and each mixture for each run was

rapidly added and swirled until the appearance of the blue

color. At the instant the mixtures were mixed together, the

time was recorded until the appearance of the blue color.

Temperature of the blue mixture was also recorded.

The second part of the experiment was the

determination of the reaction order with respect to I-. The

procedure in the first part was done again but this time, the

concentration of the ferric ion was kept constant and the

concentration of the iodine was varied. A total of three runs

were made in this part of the experiment.

The third part of the experiment was the effects of

Fe2+

. The first run of the first part was done but the HNO3

solution was replaced with 0.002M of Fe(NH4)2(SO4)2 in

HNO3 solution.

III. RESULTS AND DISCUSSION

The experiment aims to investigate the oxidation of

the iodide ions by the ferric ions. Several mixtures were

prepared with varying concentrations of the substances

involved for the desired reactions. Moreover, the chemical

changes necessary in this clock reaction are the reactions

mentioned earlier. The limiting reagent is the Sodium

thiosulfate and the starch is the indicator. The signaling

species in this clock reaction is the starch-pentaiodine

complex or the starch-I5- complex which is blue.

Table 1. Change in Time and Initial Rate of Reactions

Table 1 shows the change in time obtained from the

experiment. Runs 1 to 5 pertains to the first part of the

experiment in which the concentration of Iodide was kept

constant. Runs 6 to 8 is for the determination of the reaction

order with respect to I- in which the concentration of the Ferric

Ion was held constant. Since the rates of the reaction of the

ferric and iodide ions are dependent with respect to their

respective concentrations, we can express the rate as:

Where k is the reaction rate constant and a is the reaction

order of Fe3+

and b is the reaction order of I-.

To be able to determine the reaction order of the Fe3+

ion, the

slope of the line from the graph of log rate versus the log of

[Fe3+

] wherein the concentration of ferric ions was varied at

different runs.

Run Change in Time

(s)

M Thiosulfate Rate Log rate

1 48.66 0.0004 4.11x10-6 -5.38614

2 34.24 0.0004 5.84x10-6 -5.2335

3 27.78 0.0004 7.20x10-6 -5.1427

4 25.15 0.0004 7.95x10-6 -5.09951

5 24.67 0.0004 8.11x10-6 -5.09114

6 169.43 0.0004 1.18x10-6 -5.92796

7 52.20 0.0004 3.83x10-6 -5.41664

8 14.97 0.0004 1.34x10-6 -4.87419

y = 0.6193x - 3.8652 R² = 0.9561

-5.5

-5.4

-5.3

-5.2

-5.1

-5

-3 -2.5 -2 -1.5 -1 -0.5 0

Log

[R

ate

]

Log [Fe3+]

Log [Rate] vs. Log [Fe3+]

y = 1.562x - 1.7523

R² = 0.887

-7

-6

-5

-4

-3

-2

-1

0

-3 -2.5 -2 -1.5 -1 -0.5 0

Log

[Rat

e}

Log [I-]

Log [Rate] vs. Log [I-]

Table 2. Log [Fe3+

] from Initial Concentration of [Fe3+]

The average value of the initial and final concentrations of the

ferric ions was used because of the presence of sodium

thiosulfate that yields a thiosulfate complex with a reversible

reaction as can be seen from the equation:

And another redox reaction that has an order of zero and

another with a reaction order of two that is negligible and

won’t change anything from the reaction that was being

investigated.

The logarithm of the initial rate was plotted against the

logarithm of the average ferric ion to be able to obtain the

equation of the line and more importantly, the slope of the line

which corresponds to the order of the reaction with respect to

the ferric ions. The equation of the line was found out to be

y=0.6193x-3.8652, and the slope is then determined which is

0.6193. Since there is no reaction order of 0.6, it was rounded

up and therefore giving a value of the order of the reaction of

approximately equal to 1. A first order of reaction has a rate

proportional to the concentration of one of the reactants,

which in this case is the ferric ion. [4]

The second part of the experiment corresponds to

runs 6 to 8 of table 1. This was done to determine the reaction

order with respect to iodide therefore the concentration of the

ferric ion was held constant throughout the process.

Table 3. Log [I-] from the Initial Concentration of iodide

Table 3 shows the varied concentration of iodide with

the concentration of ferric ion was held constant. The

logarithm of the iodide ion was obtained to be plotted against

the logarithm of the initial rate and obtain the equation of the

line.

Figure 3. Log rate vs Log [I-]

The plot of the logarithm of the rate of reaction was

plotted against the logarithm of the average concentration of

Iodide. The equation of the line was determined and found out

to be y = 1.562x-1.7523. Though this, the slope of the line can

directly be determined which then gives the value of 1.562.

This slope of the line corresponds to the order of the reaction

with respect to iodide. But since there is no order of reaction

which is 1.562, the value was rounded up and therefore will

give a value of the order of reaction approximately equal to 2,

which indicates a second order of reaction with respect to the

concentration of iodide. A second order of reactions implies

that the initial rate of the reaction quadruples. [5]

By determining the reaction order of the two

reactants, the overall order of the reaction can be determined

by summing up the two values of the orders which then will

lead to an overall reaction order of three which implies that the

rate of reaction increases eightfold. Moreover, by obtaining

the orders of reaction, the rate constant denoted with k will be

determined through the rate equation:

Rate = k[Fe3+

][I-]

2

Run [Fe3+]i [Fe3+]f [Fe3+]ave log

[Fe3+]av [I-]ave log rate

1 0.004 0.0036 0.0038 -2.42022

0.004 -5.3861

2 0.006 0.0056 0.0058 -2.23657 0.004 -5.2335

3 0.008 0.0076 0.0078 -2.10791 0.004 -5.1427

4 0.010 0.0096 0.0098 -2.00877 0.004 -5.0991

5 0.012 0.0116 0.0118 -1.92812 0.004 -5.0914

Run [Fe3+] [I-]ave log [I-]ave log rate

6 0.004 0.002 -2.69897 -5.92796

7 0.004 0.006 -2.22184 -5.41664

8 0.004 0.008 -2.09691 -4.87419

Figure1. Log [Rate] vs Log [Fe3+]

Table 4. Average Rate Constant

Table 4 shows the calculation of the average rate

constant for the chemical reaction. Through the use of the

identified rate law with the orders of the respective reactants

known, the rate constant was found out to be 54.31.

The last part of the experiment is to identify the

effect of Fe2+

on the rate of reaction of the ferric and iodide

ions. The change in time that was recorded in the experiment

was 76.77 seconds. The same concept as the run 1 of the first

part was followed for the calculation of the rate,

[ ]

M/s

Comparing the rate that was obtained in the first run

of the first part, the rate was reduced by approximately 63%

which indicates that with the addition of Fe2+

to the solution

reduces the rate of the reaction of Iodide and Ferric ions.

IV. ERROR ANALYSIS

The errors in the experiment may have surfaced and

caused a deviation from the true values. The change of time

may have caused an error since it varies from one person to

another. Another error that may have surfaced in the

experiment is the end point interpretation since it was done

by different persons and again may vary from one person to

another. Lastly, in the preparation of the solutions, there

might be excess or lack of the amount of substance in the

mixture that may have caused the change in concentration.

Moreover, impurities in the equipment used also may play a

big factor in the execution of the experiment since this may

cause a change in the concentration of the solutions and the

experiment depends greatly on the concentration of the

solution.

V. CONCLUSION

The order of the reactants was obtained in the

experiment. For the ferric ion, it was found out to be one or

the first order. On the other hand the order of reaction with

respect to iodide is 2 which is in second order. The over-all

order of the reaction was also determined through summing up

the orders of reaction of the reactants and it lead to a third

order of the reaction. Moreover, the reaction rate constant was

also obtained and it gave a reaction rate constant of 54.31.

REFERENCES

[1]

HTTP://WWW.CHM.DAVIDSON.EDU/VCE/KINETICS/METHODOFI

NITIALRATES.HTML (DATE ACCESSED: FEBRUARY 22,2014)

[2] http://www.chem4kids.com/files/react_rates.html (date accessed: 02/22/2014)

[3] http://www.kbcc.cuny.edu/academicdepartments/ (date

accessed 02/22/2014)

[4]

http://chemwiki.ucdavis.edu/Physical_Chemistry/Kinetics/Rat

e_Laws/The_Rate_Law (data accessed: 03/01/2014)

[5]

http://www.chemguide.co.uk/physical/basicrates/experimental

.html (date accesed: 03/01/2014)

APPENDICES

I. CALCULATIONS

(a) Determination of the Concentration of Sodium Thiosulfate

(b) Determination of Initial Rate

[Rate] = 4.11x10-6

M/s

(c) Calculation of Iodide Concentration

Rate M [Fe3+] M [I-] k Kave

4.11x10-6 0.0038 0.004 67.60

54.31

5.84x10-6 0.0058 0.004 62.94

7.20x10-6 0.0078 0.004 57.69

7.95x10-6 0.0098 0.004 50.72

8.11x10-6 0.0118 0.004 42.94

1.18x10-6 0.004 0.002 73.78

3.83x10-6 0.004 0.006 26.61

1.34x10-6 0.004 0.008 52.19

(d) Calculation of Ferric Ion Concentration

(e) Calculation of Ferric ions final Concentration

(f) Calculation of Average Ferric Ion Concentration

(g) Calculation of the rate Constant, k:

k = 67.60