Quantitative Analysis
Iodometric Titrations

Iodometric titration is an indirect volumetric analysis method where an oxidizing agent reacts with excess iodide ions (\(\text{I}^-\)) to liberate iodine (\(\text{I}_2\)). The liberated iodine is then titrated against a standard reducing agent, usually sodium thiosulfate (\(\text{Na}_2\text{S}_2\text{O}_3\)), using starch solution as an indicator.

$$\text{General Stoichiometry: } \frac{C_A V_A}{C_B V_B} = \frac{n_A}{n_B}$$ $$\text{Reaction: } \text{I}_2 + 2\text{S}_2\text{O}_3^{2-} \longrightarrow 2\text{I}^- + \text{S}_4\text{O}_6^{2-}$$ $$\text{Mass conc} = \text{Molar conc.} \times \text{Molar Mass} $$

Calculations

Example I

(a) What mass of iodine reacts with $22.80\text{ cm}^3$ of $0.011\text{ mol/dm}^3$ sodium thiosulphate(VI)?

(b) $20.0\text{ cm}^3$ of solution of iodine in potassium iodide solution required $23.50\text{ cm}^3$ of $0.045\text{ mol/dm}^3$ sodium thiosulphate(VI) solution for complete oxidation.
Calculate the:
I. Concentration of iodine in $\text{mol/dm}^3 $
II. Concentration of iodine in $\text{g/dm}^3 $

Equation for reaction:

$$\text{I}_2 + 2\text{S}_2\text{O}_3^{2-} \longrightarrow 2\text{I}^- + \text{S}_4\text{O}_6^{2-}$$
Solution
$$ \text{Find moles of thiosulphate ion} $$ $$n_{\text{S}_2\text{O}_3^{2-}} = \frac{\text{C} \times \text{V}(cm^3)}{1000}$$ $$ n = \frac{0.011 \times 22.80}{1000} $$ $$ = 0.0002508\text{ moles} $$ $$ \text{Determine moles of iodine } ({I}_2) $$ $$\text{From the balanced equation} $$ $$2\text{ moles of S}_2\text{O}_3^{2-} \rightarrow 1\text{ mole of I}_2$$ $$ \text{moles of iodine} = \frac{\text{S}_2\text{O}_3^{2-}}{2} $$ $$\text{Moles of iodine} = \frac{0.0002508}{2} $$ $$ = 0.0001254\text{ moles} $$ $$ \text{Calculate mass of iodine ($\text{I}_2$)} $$ $$\text{Molar mass of I}_2 = 127 \times 2 $$ $$ = 254\text{ g/mol} $$ $$\text{Mass of iodine} = \text{Moles} \times \text{Molar mass}$$ $$\text{Mass of iodine} = 0.0001254 \times 254 $$ $$ = 0.03185\text{ g}$$

Write out parameters

$$ OA = I_2 $$ $$ RA = Na_2S_2O_3 $$ $$V_{OA} = 20cm³ $$ $$V_{RA} = 23.50cm³ $$ $$C_{RA} = 0.045mol/dm³ $$ $$n_{OA} = 1 $$ $$n_{RA} = 2 $$

I. Concentration of iodine in mol/dm³

$$\frac{C_{\text{OA}} \times V_{\text{OA}}}{C_{\text{RA}} \times V_{\text{RA}}} = \frac{n_{\text{OA}}}{n_{\text{RA}}}$$ $$C_{\text{OA}} = \frac{C_{\text{RA}} \times V_{\text{RA}} \times n_{\text{OA}}}{V_{\text{OA}} \times n_{\text{RA}}}$$ $$C_{\text{OA}} = \frac{0.045 \times 23.50 \times 1}{20.0 \times 2}$$ $$C_{\text{OA}} = \frac{1.0575}{40} = 0.0264\text{ mol/dm}^3$$

II. Concentration of iodine in $\text{g/dm}^3$:

$$\text{mass conc.} = \text{molar conc} \times \text{Molar mass}$$ $$\text{Conc. in g/dm}^3 = 0.02644 \times 254 $$ $$ = 6.72\text{ g/dm}^3$$

Example II: $20.00\text{ cm}^3$ of iodine solution was titrated with $0.09\text{ mol/dm}^3$ sodium thiosulphate solution and the iodine reacted with $15.50\text{ cm}^3$ of sodium thiosulphate solution).

(a) Write equation for the reaction.
(b) State the indicator that should be used at the end point and the change at the end point.
(c) Calculate the concentration of the iodine solution in $\text{mol/dm}^3$ and $\text{g/dm}^3$.

Solution

(I) Equation for the reaction:

$$\text{I}_2 + 2\text{S}_2\text{O}_3^{2-} \longrightarrow 2\text{I}^- + \text{S}_4\text{O}_6^{2-}$$

(II) Indicator & End Point:
Indicator: Starch indicator solution
Color change at end point: Blue-black to colorless

(III) Concentration:

$$C_{\text{OA}} = \frac{C_{\text{RA}} \times V_{\text{RA}} \times n_{\text{OA}}}{V_{\text{OA}} \times n_{\text{RA}}}$$ $$C_{\text{OA}} = \frac{0.09 \times 15.50 \times 1}{20.00 \times 2}$$ $$C_{\text{OA}} = \frac{1.395}{40} = 0.0349\text{ mol/dm}^3$$

Concentration in $\text{g/dm}^3$:

$$\text{Conc. in g/dm}^3 = 0.0349 \times 254 $$ $$ = 8.87\text{ g/dm}^3$$

Example III: $20.00\text{ cm}^3$ of solution of iodine in potassium iodide solution was titrated against a solution of $0.01\text{ mol/dm}^3$ of sodium thiosulphate using soluble starch as indicator.

Burette reading 1st titre 2nd titre 3rd titre
Final burette reading $\text{S}_2\text{O}_3^{2-} (\text{cm}^3)$ 46.80 36.60 41.90
Initial burette reading $\text{S}_2\text{O}_3^{2-} (\text{cm}^3)$ 18.30 10.40 15.60
Volume of $\text{S}_2\text{O}_3^{2-}$ used $(\text{cm}^3)$ 28.50 26.20 $Y$

(I) Titre value of $Y$ for the third titration and hence determine the average titre value.
(II) Calculate the concentration of iodine in (a) $\text{mol/dm}^3$ and (b) $\text{g/dm}^3$. $$[\text{Na} = 23, \text{S} = 32, \text{O} = 16, \text{I} = 127]$$

Solution

(I) Value of $Y$ and Average Titre:

$$Y = 41.90 - 15.60 = 26.30\text{ cm}^3$$

Average volume of $\text{S}_2\text{O}_3^{2-}$ (using concordant 2nd & 3rd titrations):

$$V_{\text{RA}} = \frac{26.20 + 26.30}{2} = 26.25\text{ cm}^3$$

(II) Concentration of iodine:

(a) Concentration in $\text{mol/dm}^3$:

$$C_{\text{OA}} = \frac{C_{\text{RA}} \times V_{\text{RA}} \times n_{\text{OA}}}{V_{\text{OA}} \times n_{\text{RA}}}$$ $$C_{\text{OA}} = \frac{0.01 \times 26.25 \times 1}{20.00 \times 2}$$ $$C_{\text{OA}} = \frac{0.2625}{40} = 0.00656\text{ mol/dm}^3$$

(b) Concentration in $\text{g/dm}^3$:

$$\text{Conc. in g/dm}^3 = 0.0065625 \times 254 $$ $$ = 1.67\text{ g/dm}^3$$
Iodimetric titrations

Iodimetry is a direct titration method in quantitative redox analysis where a standard solution of iodine (\(\text{I}_2\)) is used as an oxidizing titrant to directly determine the concentration of reducing agents (such as thiosulfate, sulfite, arsenite, or ascorbic acid). Because elementary iodine is sparingly soluble in water, it is dissolved in an aqueous solution of potassium iodide (\(\text{KI}\)) to form the triiodide complex ion (\(\text{I}_3^-\)), which remains stable and reactive throughout the titration process.

Precautions in iodometry
  1. Add starch indicator near the end point: Starch should not be added at the beginning of the titration when the concentration of iodine is high, as iodine forms an insoluble, highly stable complex with starch that decomposes very slowly, leading to inaccurate end-point determination and an overestimation of titrant volume.
  2. Prevent iodine volatilization: Free iodine is volatile. Titration flasks should be kept covered with a watch glass or stoppered when not in use, and titrations should be conducted at cool or room temperature to avoid iodine loss due to evaporation.
  3. Prevent atmospheric oxidation of iodide: In acidic media, excess iodide ions (\(\text{I}^-\)) can be oxidized by atmospheric oxygen (\(\text{O}_2\)) to liberate extra iodine (\(4\text{I}^- + \text{O}_2 + 4\text{H}^+ \longrightarrow 2\text{I}_2 + 2\text{H}_2\text{O}\)). Titrations should be performed quickly and, if necessary, under an inert atmosphere (such as sodium bicarbonate addition to generate a $\text{CO}_2$ blanket).
  4. Maintain appropriate pH levels: Iodometric reactions must be carried out under neutral or weakly acidic conditions. Strong acid environments accelerate air oxidation of iodide, while strongly alkaline environments cause iodine to undergo dismutation into iodate (\(\text{IO}_3^-\)) and hypoiodite (\(\text{IO}^-\)), interfering with stoichiometry.
  5. Use freshly prepared starch solution: Starch solutions undergo bacterial decomposition over time, yielding glucose breakdown products that act as reducing agents or produce a reddish color change instead of the sharp blue-black-to-colorless transition.
  6. Ensure adequate reaction time for iodine liberation: In iodometric procedures where iodine is liberated (e.g., reactions involving $\text{Cu}^{2+}$ or $\text{Cr}_2\text{O}_7^{2-}$), the reaction mixture must be allowed to stand in the dark for 3–5 minutes before titrating to ensure complete reaction while preventing light-catalyzed air oxidation.
Summary