Finding the end point

Finding the end point

• A redox indicator is a compound that changes color when it goes from its oxidized to its reduced state .

or

For ferroin, with E° = 1.147 V we expect the color change to occur in the approximate range 1.088 V to 1.206 V with respect SHE

A redox titration is feasible if the difference between analyte and titrant is > 0.2 V.

If the difference in the formal potential is > 0.4 V, then a redox

indicator usually gives a satisfactory end point.

• Starch is the indicator of choice for those procedures

involving iodine because it forms an intense blue complex with iodine.

• Starch is not a redox indicator; it responds specifically to the presence of I

• The active fraction of starch is amylose, a polymer of the sugar α-d-glucose.

• In the presence of starch, iodine forms I

chains inside the amylose helix and the color turns dark blue

Starch-Iodine Complex

Adjustment of analyte oxidation state

• Sometimes one needs to adjust the oxidation state of analyte before it can be titrated.

• Pre-oxidation: powerful oxidants that can be easily removed after preoxidation include peroxydisulfate, silver(II) oxide, sodium

bismuthate.

Definition:

• Disproportionation: a reactant oxidizes and reduces itself, such as H

O

in boiling water.

• Pre-reduction: Process of reducing an analyte to a lower oxidation state prior to performing a titration with an oxidizing agent.

• Amalgam: a solution of anything in mercury.

Oxidation with potassium permanganate

KMnO ₄ is a strong oxidant with an intense violet color.

In strongly acidic solutions, it is reduced to Mn ²⁺ .

In neutral or alkaline solution, it is reduced to brown solid MnO ₂ .

In strongly alkaline solution ( 2 M NaOH),

manganate ion (MnO _42- ) is produced.

Oxidation with Ce ⁴⁺

• Reduction of Ce

to Ce

proceeds cleanly in acidic solutions.

• The aquo ion, Ce(H

0)

, probably does not exist in any

of these solutions, because Ce(IV) binds anions (ClO

,

SO

NO

, Cl

) very strongly.

Methods Involving Iodine

• Iodimetry: a reducing analyte is titrated directly with iodine (to produce I

).

• iodometry, an oxidizing analyte is added to excess I

to produce iodine, which is then titrated with standard

thiosulfate solution.

• Iodine only dissolves slightly in water. Its solubility is enhanced by interacting with I

• A typical 0.05 M solution of I

for titrations is prepared by dissolving 0.12 mol of KI plus 0.05 mol of I

in 1 L of

water. When we speak of using iodine as a titrant, we

almost always mean that we are using a solution of I

plus excess I

.

Preparation and Standardization of Solutions

• Acidic solutions of I

are unstable because the excess I

is slowly oxidized by air:

• In neutral solutions, oxidation is insignificant in the absence of heat, light, and metal ions. At pH 11, ≳ triiodide disproportionates to hypoiodous acid (HOI), iodate, and iodide.

• An excellent way to prepare standard I

is to add a

weighed quantity of potassium iodate to a small excess

of KI. Then add excess strong acid (giving pH ≈ 1) to

produce I

by quantitative reverse disproportionation: