ABSTRACT

This experiment investigated the partitioning behaviour of iodine between chloroform and
water in a two-phase system. A saturated solution of iodine in chloroform was mixed with
varying volumes of chloroform and a fixed amount of water, and the system was allowed to
reach equilibrium. After phase separation, the concentration of iodine in each layer was
determined by titration with standardized sodium thiosulphate solution. The concentrations of
iodine in the aqueous phase decreased from 1.39 to 0.77 mol/m³ across the bottles, while
those in the chloroform phase ranged from 141.0 to 99.0 mol/m³. The calculated partition
coefficients were 101.4, 84.6, and 128.6, confirming iodine’s preferential solubility in the
organic phase. Slight variations in these values were attributed to experimental errors, but the
overall results support the theoretical expectation of strong iodine affinity for chloroform.
INTRODUCTION

The distribution of a solute between two immiscible liquids is a fundamental concept in

physical chemistry and underpins various analytical and separation techniques, including
solvent extraction, chromatography, and pharmaceutical drug delivery systems (Skoog et al.,
2013). When a solute that is soluble in both phases is added to a biphasic system, it
distributes itself between the two layers in a fixed ratio at equilibrium, provided the solute
does not undergo chemical change and exists in the same molecular form in both solvents
(Atkins & de Paula, 2014).

This ratio is defined by the partition coefficient (or distribution coefficient), typically
expressed as:

𝐶𝐴
𝐾=
𝐶𝐵

Where CA and CB represent the equilibrium concentrations of the solute in solvent A and
solvent B, respectively. The partition coefficient is influenced by the nature of the solute and
solvents, temperature, and the presence of impurities or complexing agents (Christian &
O'Reilly, 1986).

The iodine–chloroform–water system is a classic model for studying phase partitioning due to
iodine's significantly higher solubility in non-polar organic solvents like chloroform
compared to polar solvents like water. This sharp difference allows for accurate
quantification of the iodine distribution using redox titration techniques, specifically titration
with standardized sodium thiosulfate solution in the presence of starch as an indicator
(Jeffrey et al., 1989).

Understanding the behaviour of iodine in this biphasic system not only illustrates the
principles of chemical equilibrium and solubility but also has practical implications in
extraction chemistry and environmental science, where the transport and fate of halogens in
water–organic interfaces are relevant (Skoog et al., 2013).

METHODOLOGY
A saturated solution of iodine in chloroform was prepared by dissolving iodine crystals in
chloroform until no more dissolved. The solution was filtered to remove undissolved solids.
Three reagent bottles were labelled 1, 2, and 3. Each was filled with 150 mL of distilled water
and varying volumes of saturated iodine solution and pure chloroform: Bottle 1 received 40
mL of iodine solution and no pure chloroform; Bottle 2, 30 mL iodine solution and 10 mL
chloroform; Bottle 3, 25 mL iodine solution and 15 mL chloroform. All bottles were sealed
and shaken vigorously for 30 minutes to ensure adequate mixing and phase contact.

After shaking, the mixtures were allowed to stand until complete phase separation occurred.
The chloroform (lower) and aqueous (upper) layers from each bottle were carefully separated
using a separating funnel, with attention to avoid cross-contamination between phases. The
separated layers were collected in clean, labelled beakers.

For each bottle, a 25 mL sample of the aqueous layer was transferred to a conical flask, and
3–4 drops of starch solution were added as an indicator. The sample was titrated with 0.01 M
sodium thiosulphate until the blue starch–iodine complex just disappeared. Titrations were
repeated until two concordant readings were obtained.

Similarly, for the chloroform layers, 5 mL of each organic sample was transferred to a
conical flask containing approximately 25 mL of water to extract the iodine into the aqueous
phase. Starch solution (3–4 drops) was added, and the mixture was titrated with 0.1 M
sodium thiosulphate until the endpoint was reached. The procedure was repeated to obtain
concordant results for each chloroform sample.

All titration values were recorded and used to calculate the concentration of iodine in both
aqueous and chloroform phases. These values were then used to determine the partition
coefficient for each mixture.

RESULTS AND DISCUSSION

Tables 1 and 2 present the titration volumes of sodium thiosulphate required to quantify
iodine in the aqueous and chloroform layers, respectively. Table 1 shows volumes of 0.01 M
thiosulphate used for the aqueous layer, while Table 2 shows volumes of 0.1 M thiosulphate
used for the chloroform layer
Table 1: volume of 0.01M thiosulphate titrated with aqueous layer

Bottle number Volume of 0.01M thiosulphate titrated

1 7.2
1 6.7
2 5.4
2 5.3
3 3.9
3 3,8

Table 2: volume of 0.1M thiosulphate titrated with chloroform layer

Bottle number Volume of 0.1M thiosulphate titrated

1 14.1
1 14.1
2 8.3
2 9.8
3 10
3 9.8

From Table 1, it is evident that the volume of thiosulphate required to titrate the aqueous
layer decreases across bottles 1 to 3. This indicates that the concentration of iodine in the
aqueous phase decreases as the volume of pure chloroform increases. Similarly, in Table 2,
the titration volume required for the chloroform layer also decreases from bottle 1 to bottle 3,
though the drop is more gradual. This suggests that more iodine is extracted into the organic
layer when the proportion of chloroform is higher.

However, the overall trend shows that while the total amount of iodine remains roughly
constant across the bottles, its distribution shifts increasingly toward the chloroform phase as
more chloroform is present. This behaviour aligns with the principle of phase partitioning,
where a solute distributes between two immiscible solvents in a fixed ratio governed by the
partition coefficient. Since iodine is significantly more soluble in chloroform than in water
due to its nonpolar nature, increasing the volume of chloroform enhances its capacity to
dissolve iodine, reducing the amount remaining in the aqueous phase.

Table 3: concentration of iodine in aqueous layer

Bottle number Concentration (mol/m3)

1 1.39
2 1.07
3 0.77
Table 4 concentration of iodine in chloroform layer

Bottle number Concentration (mol/m3)

1 141.0
2 90.5
3 99.0

The concentration of iodine in the aqueous phase decreased progressively across the three
bottles, with values of 1.39 mol/m³, 1.07 mol/m³, and 0.77 mol/m³ for bottles 1, 2, and 3
respectively. In the chloroform phase, the concentrations were 141.0 mol/m³, 90.5 mol/m³,
and 99.0 mol/m³ respectively. These values confirm that iodine preferentially dissolves in the
organic phase (chloroform) rather than in water, consistent with its nonpolar character and the
polar–nonpolar solvent distribution principle.

Table 5: the partition coefficients per bottle

Bottle number Partition coefficient

1 101.4
2 84.6
3 128.6

The corresponding partition coefficients, calculated as the ratio of iodine concentration in

chloroform to that in water, were 101.4, 84.6, and 128.6 for bottles 1 through 3 respectively.
Although the partition coefficient is theoretically constant under fixed temperature and
pressure, the observed values show some variation. This is likely due to experimental errors
such as inconsistent shaking, layer contamination, or titration inaccuracies. Despite these
deviations, the values are all within a reasonable range and reaffirm that iodine strongly
favours the chloroform phase. The overall results support the principle of phase partitioning
and demonstrate the practical use of titration to quantify solute distribution. The average
partition coefficient remains close to 100, which aligns well with literature values and
confirms the strong solubility of iodine in chloroform under the conditions used.
CONCLUSION

The experiment successfully demonstrated the partitioning of iodine between chloroform and
water in a two-phase system. It showed how a solute distributes itself between immiscible
solvents based on its relative solubility in each phase. The observed concentrations in both
layers allowed for the calculation of partition coefficients, which reflected iodine’s strong
preference for the organic phase. This behaviour aligned with theoretical expectations,
emphasizing the influence of molecular polarity on solute distribution. While minor
discrepancies were noted in the data, they did not significantly affect the outcome or overall
interpretation. The experiment also provided practical experience in titration techniques,
proper phase separation, and the quantitative treatment of equilibrium data, reinforcing key
analytical skills and concepts in physical chemistry.
REFERENCES

Atkins, P., & de Paula, J. (2014). Atkins' Physical Chemistry (10th ed.). Oxford University

Press. Oxford, United Kingdom.

Christian, G. D., & O'Reilly, J. E. (1986). Instrumental Analysis (2nd ed.). Allyn and Bacon.

Boston, United States.

Jeffery, G. H., Bassett, J., Mendham, J., & Denney, R. C. (1989). Vogel's Textbook of

Quantitative Chemical Analysis (5th ed.). Longman Scientific & Technical. Harlow, UK

Skoog, D. A., West, D. M., Holler, F. J., & Crouch, S. R. (2013). Fundamentals of Analytical

Chemistry (9th ed.). Cengage Learning. Boston, United States.

APPENDIX

Reaction Equation

I₂ + 2 S₂O₃²⁻ → 2 I⁻ + S₄O₆²⁻

1. Aqueous Layer (0.01 M Thiosulphate Titrant)

Moles of iodine: n(I₂) = ½ × Cthiosulphate × Vthiosulphate, with C = 0.01 mol/L

Sample volume = 25 mL = 0.025 L

CI aqueous = n(I₂) / 0.025 (mol/m³)

Bottle Avg Volume Moles I₂ Sample CI Aqueous

(mL) Volume (L) (mol/m³)
1 6.95 3.475 × 10⁻⁵ 0.025 1.39
2 5.35 2.675 × 10⁻⁵ 0.025 1.07
3 3.85 1.925 × 10⁻⁵ 0.025 0.77

2. Chloroform Layer (0.1 M Thiosulphate Titrant)

Sample volume = 5 mL = 0.005 L

n(I₂) = ½ × Cthiosulphate × Vthiosulphate, with C = 0.1 mol/L
CI chloroform = n(I₂) / 0.005 (mol/m³)
Bottle Avg Volume Moles I₂ Sample CI chloroform
(mL) Volume (L) (mol/m³)
1 14.1 7.05 × 10⁻⁴ 0.005 141.0
2 9.05 4.525 × 10⁻⁴ 0.005 90.5
3 9.9 4.95 × 10⁻⁴ 0.005 99.0

3. Partition Coefficients (K)

K = CI chloroform / CI Aqueous
Bottle CI Chloroform CI Aqueous Partition
(mol/m³) (mol/m³) Coefficient (K)
1 141.0 1.39 101.4
2 90.5 1.07 84.6
3 99.0 0.77 128.6