Reduction- Oxidation Reactions 5th lecture. Ceric as titrant: Ce 4+

Reduction- Oxidation Reactions

5thlecture

Ceric as titrant: Ce4+

Although it could be used as self indicator it is preferable to use ferroin as indicator especially in case of det. of ferrous salts.

Ceric as titrant: Ce4+Properties

Ce4+ salts are strong oxidants in H2SO4

Ce4+ + e Ce3+ Yellow Colorless

They have wide range of oxidising power but they don’t oxidise HCl even in presence of Fe2+ salts

Ce4+ cannot be used in neutral or alkaline solution due to hydrolysis to hydrated ceric oxide

Ceric salts are much more stable than MnO4-

Ce4+ forms more stable complexes than Ce3+


Preparation and standardization of Ce4+ soluPrepared from primary standard Ce(NO3)6 (NH4)2 in conc H2SO4 or in 72% HClO4. If using other salts it should be standardized

(1) Against arsenious trioxide:2Ce4+ + H3 AsO3 + H2O 2Ce3+ + H3AsO4+ 2H+

(2) Against oxalate

In both cases , the reaction is slow it requires heat to 50°C, using ICl as catalyst and ferroin indicator

2Ce4++ H2C2O4 ↔ 2Ce3+ + 2CO2 + 2H+


(a)Direct titrations: determination of reducing agents Fe2+, AsO3

3- , C2O42-, H2O2, I-,

Fe(CN)64- using ferroin indicator

Color change from red to pale blue [ Fe (CN)6]4-+ Ce4+ Ce3++ [ Fe (CN)6]3-

Advantages: Better than MnO4

- as it is less subject to interference of organic matterIt is preferable to be used instead of MnO4

- in the determination of Fe2+ since we can use HCl.

Applications

H2O2 + 2Ce4+ 2Ce3+ + 2H+ + O2


Applications of Ce4+

(b) Back titrations: Determination of polyhydroxy

alcohols, aldehydes, hydroxy acids. example: glycerol, citric acid

C3H8O3+8Ce4++3H2O 3HCOOH+8Ce3++8H+

The excess Ce4+ is titrated against sodium oxalate or AsO3

3- using ICl as catalyst and ferroin as indicator at 50oC.

Potassium dichromate as titrant


It is a primary standard due to the stability of itssolution and is obtainable in high purity

Its oxidation potential is lower than KMnO4 and Ce4+ so it is limited in use

It does not oxidise Cl- into Cl2, oxalic acid ,ferrocyanide Its main application is the direct and indirect determination of Fe2+ ion

Properties

Many redox indicators are unsuitable:•because of their high oxidation potential, and •because of the deep green colour of Cr3+ which causes the colour change of the indicator to be less clear

The indicators usually used are:•diphenyl amine sulphonic acid. •4,7-dimethyl 1, 10 phenanthroline ferrous.


It can not serve as a self indicator reagent

Cr2O72- (Orange) + 14H+ + 6e 2Cr3+ (green) + 7H2O

Potassium dichromate as titrant Applications

1-Determination of Fe2+ Iron (internal indicator)

diphenylamineH2SO4

Fe2+

Titrate with Cr2O7

2-

E0 = 1.33 vFe3+

decreases the Fe3+/Fe2+ system potential so that Fe2+ ion will be oxidized before the indicator

and to remove the dark colour of Fe3+ ion giving a more clear colour change.

Role of H3PO4 or F-:

E° Fe3+/Fe2+ 0.77E° diphenylamine 0.76

H3PO4

E° ferroin 1.06Is there need for

H3PO4??


1-Determination of Fe2+ Iron external indicator

Ferricyanide:Fe2+ is titrated with dichromate in acidic medium.Occasionally remove a drop from the solution and add it to ferricyanide solu. a blue color of ferrous ferricyanide is formed. At the E.P. No more Fe2+ is present so no blue color is formed.Diphenylcarbazide: After oxidation of Fe2+ to Fe3+, the first exx of dichromate oxidizes the indicator and gives a red color.

2- Determination of some oxidising agents Add a measured exx of Fe2+ ion and back titrate

the exx. using Cr2O72- and diphenylamine as

indicator.



3-reducing agents 4-Organic compd 5-Pb2+ Na2SO3 glycerol PbO Add measured exx of Cr2O7

2- in presence of:

Sulphuric acid Sulphuric acid glacial HACICl as catalyst The excess dichromate is titrated iodometrically

3 SO32- + Cr2O7

2- + 8H+ 3 SO42- + 2Cr3+ + 4H2O

Cr2O72- + 6I- + 14 H+ 2Cr3+ + 3I2 + 7

H2O

Applications

3C3H8O3 + 7 Cr2O72- + 56 H+ 14Cr3+ + 9 CO2 +

40H2O2Pb2+ + Cr2O7

2- + H2O 2PbCrO4↓ (ppt)+ 2H+

Iodine as oxidant

The iodine/iodide half reaction is I2 + 2e 2I- (Eo = +0.535V)I- can be oxidized by systems I2 can oxidize systems of of higher oxidation potential lower oxidation potential MnO4

-/Mn2+ Sn4+/Sn2+

Cr2O72-/Cr3+ S4O6

2-/S2O32-

ClO3-/Cl- S/S2-

Iodine as oxidantProperties:

↑ E° ↓ E°

Iodometric method Indirect titration• Add KI to oxidizing agents, equivalent I2 is libarated andtitr with Na2S2O3

•To determine oxidizing agents

Iodimetric methodDirect titration with I2

•To determine reducing agents

Systems having oxidation potentials near to that of iodine/iodide e.g AsO4

3-/AsO33-, Fe3+/Fe2+

Their reactions with Iodine is directed forward or backword by control of experimental conditions.

i.e. Change in oxidation potential

Iodine as oxidantProperties:

1-the pH of the medium2-addition of complexing agents3-addition of precipitating agents

1-Effect of pH: The potential of: AsO4

3-/AsO33-= +0.57

I2/2I- = +0.54

To determine arsenite sample using Iodinethe pH of the solution should be adjusted to 8.3

by adding NaHCO3

I2 + AsO33- + H2O 2I- + AsO4

3- + 2H+

E AsO43-

/ AsO33-

=Eo – 0.059 / 2 log [AsO33- ] / [AsO4

3-][H+]2

↓ [H+] by addition of NaHCO3 ↓ the oxidation

potential of AsO43- / AsO3

3- system.NaHCO3 reacts with H+ giving CO2 and H2O shifting the reaction

to the right and prevent reversibility.At higher pH if using NaOH, I2 reacts with OH- producing OI- so

consuming more I2. Also OI- has oxidizing properties which differ than I2.

Factors affecting the potential of I2/I- system:

2-Effect of Complexing agents:Iodine as oxidant

When HgCl2 is added to the I2/I- system it forms

[HgI4]2- Thus:

removing the I- ions from the share of the reaction,

minimizing its concentration, increasing the ratio of I2 / [I-]2

increasing the oxidation potential of I2 /2I- systemSo I2 could determine AsO3

3-.

E= = Eo - Log [I-]2 / [ I2] 0.0592

I2 + 2 e 2I-

E° Fe3+/Fe2+= 0.77V E° I2/ 2I- = 0.54V Fe3+ + e Fe2+

E Fe2+ / Fe3+ = 0.559 - ][

][log

1

059.03

2

Fe

Fe

When pyrophosphate, EDTA or F- is added to the Fe3+/Fe2+ system it form [FeF6]3- or [Fe(PO4)6]3-

Thus: removing the Fe3+ ions from the share of the reaction,

minimizing its concentration, decreasing the ratio of Fe3+

/ Fe2+

lowering the oxidation potential of Fe3+/ Fe2+ below that of I2/2I- system.

Another example

How to determine Ferrous salts using Iodine?

Iodine as oxidant

Fe(CN)63- + e Fe (CN)6

4-

minimizing conc of ferrocyanide increasing ferri/ferro potential So Ferri/Ferro system can oxidize I- to I2

E = Eo - ][Fe(CN)

] (CN) Fe[log

1

059.0-3

6

-46

3- Effect of precipitating agents E° Ferri/Ferro= 0.36V E° I2/ 2I- = 0.54V

To determine [Fe(CN)6]3- ion iodometrically; Zn2+ should be present: it precipitate Zn2[ Fe(CN)6] ion

Iodine as oxidant

E° Cu2+/Cu+ = 0.46E° I2/2I- = 0.54

It is expected that I2 oxidizes Cu+ (cuprous), however, Cu2+

(cupric) oxidizes I-

Procedure: Cu2+ is treated with KI and the liberated I2 is titrated with S2O3

2- 2Cu2+ + 4I- I2 + Cu2I2 ↓ The precipitation of Cu2I2 increases the oxidation potential

of Cu2+ /Cu+ E = E0 - 0.059 log [Cu+] 1 [Cu2+] So Cu2+ oxidizes I- to I2

I2 tends to be absorbed on Cu2I2 so the reaction with S2O32-

is incompltete so add SCN- near the end point to form Cu2(SCN)2 which has no tendency to adsorb I2.

How to determine Cu2+ salts using KI ?

Iodine as oxidant

To reverse the reaction i.e. To allow iodine to oxidize cuprous.

Add tartarate or citrate which forms with cupric a stable complex so decreasing the oxidation potential of Cu2+/Cu+

E = E0 - 0.059 log Cu+

1 Cu+2

Titration methods:

Since iodine may be either reduced or produced by oxidation

DirectIodimetric method

IndirectIodometric method

Titrating agentIodine

for determination of reducing agents

I- is added to oxidizing agents,the librated I2

is titr. with Na2S2O3

Indicator(Starch)

Added at the beginning of titr.

Added near the end of titr (when the

brown color of I2 becomes pale)

E.P. permanent blue color

disappearance of blue color

Iodine as oxidant

Reductant +

starch

Iodine

E.P.

oxidant +

KI→I2

Na2S2O3

Add starch

Na2S2O3

Colorless E.P.

Iodine as oxidant

Detection of the end point in iodine titrations:1- The use of starch: Starch is used in the form of colloidal Solu giving a deep blue adsorbtion complex with traces I2

In exx I2 an irreversible blue adsorption complex is formed which is not changed

Starch consists of amylase and amylopectinI2 gives blue adsorption complex with amylase.In strong acid medium: starch hydrolyses giving products which give with iodine non reversible reddish color masking the end point change.

Iodine as oxidant

Detection of the end point in iodine titrations:1- The use of starch:

Starch indicator solution must be freshly prepared when it stands decomposition takes place and its sensitivity is decreased. A preservative can be added

Starch can not be used in alcoholic solu.because alcohol hinders the adsorption of I2 on starch

The sensitivity of the blue color decreases with temperature due to gelatinization of starch and volatility of Iodine

Iodine as oxidant

Iodine as oxidantDetection of the end point in iodine titrations:2- Use of organic solvent (CHCl3 or CCl4)

In presence of alcohol or conc acids, organic solvents are recommended as indicators.

These solvents dissolve iodine to give intensely coloured purple solution, so that a trace of I2 gives an intense colour, and the end point will be the appearance Or disappearance of the colour in the organic solvent layer.

I2 is soluble in CHCl3 or CCl4 90 times more than in H2O

It is important that the mixture be shaken well near the end point in order to equilibrate the iodine between the aqueous and organic phases to enable aqueous S2O3

2- to react with I2 in CHCl3

A- Error due to I2: (1)I2 is volatile especially at high temp and at a low Conc of I- ion so: ●Use stoppered glass containers ●Avoid elevated temp & cool during titratn ●Moisten the stopper with I-

I-+I2→ I3- (triiodide) less volatile and more stable

(2) I2 conc is changed if the solution gets in contact with rubber, organic matter, dust, SO2, H2S (3) I2 may undergo disproportionation into HOI and I-

I2 + H2O HOI + I- + H+

To overcome this difficulty the solution may be acidified to shift the reaction to the left.

Sources of error in iodimetry Iodine as oxidant

Sources of error in iodimetry

B- Error due to I- ion: I- ion is liable to atmospheric oxidation.

This is catalysed by light, heat, Cu2+, NO gas

The medium must be completely free from O2 so introduce CO2 (add little NaHCO3).

In titration which needs standing for time, standing should be away from light.

If we need acid medium, never use HNO3, it contains nitrous oxide.

Iodine as oxidant

4H+ + 4I- + O2 2I2 + 2H2O

Thiosulphate is affected by pH, the most favourable pH is 7 till pH 9

Under these conditions: S2O32- is oxidized to

S4O62-, where every 2 S2O3

2- is oxidized by 1 I2 to S4O6

2- (tetrathionate) 2S2O32-+I2 S4O6

2-+2I-

Under acidic conditions: thiosulphate is changed to bisulphite (HSO3

-) with the precipitation of S. Every 2 HSO3

- is oxidized by 2 I2 to 2HSO4

-

Therefore, The consumed I2 in acid medium is double that consumed in neutral medium.


C- Error due to S2O32- ion:


C- Error due to S2O32- ion:

In pH>9: I- changes to IO- (hypoiodite) oxidizing S2O3

2- to SO42- which is an incomplete reaction

Thiosulphate is decomposed during storage by thiobacteria, so:

●boiling water is used as a solvent,

●preservatives e.g. sodium benzoate, CHCl3, or HgI2 may be added.

●The pH is adjusted by adding borax, Na2CO3 or NaHCO3 to about pH 9 which inhibits bacterial action.

D- Error due to starch:

Starch may be decomposed by microorganisms into products e.g. glucose causes error due to its reducing action

other products gives nonreversible reddish color with I2 which masks the true end point.

To avoid this, preservatives e.g. H3BO3 and formamide are added.

Iodine as oxidantSources of error in iodimetry

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Reduction- Oxidation Reactions 5th lecture. Ceric as titrant: Ce 4+