Analytical Chem. 3 Course # 604 for GENERAL

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Analytical Chem. 3Course # 604 for GENERAL

Analytical Chem .2Course Code 306 for CLINICAL

2 + 1 Hours

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Team for General

Prof.Dr Mostafa Abdel-Aty

Prof.Dr. Abdel-Aziz Elbayoumy

Dr. Nesreen Khamees

Dr. Mohamad Khaled

Dr. Nesreen Talaat

Team for Clinical

Prof.Dr Mostafa Abdel-Aty

Dr Safaa Reyad

Dr Ali Yahia

Subjects

A- Redox

B- Electrochemistry

1-Conductometry

2-Potentiometry

3-Polarography

Subjects

-Redox

-Complexometry

Reference: Vogel’s textbook for quantitative chemical analysis

- Harris Quantitative Chemical Analysis

- Day and Underwood quantitative Chemical Analysis

- Skoog Quantitative Pharmaceutical Analysis

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OXIDATION-REDUCTIONTITRATIONS

REDOX

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REDOX

In the past, oxidation was defined as the reaction ofcompounds with oxygen while reduction as their reactionwith hydrogen.

But, no oxygen is involved in the oxidation of ferrous chlorideby chlorine gas according to the following equation:2 Fe2+ + Cl2 2 Fe3+ + 2 Cl-

Oxidation is the process, which results in the loss of one ormore electrons by atoms or ions.

Reduction is the process, which results in the gain of one ormore electrons by atoms or ions.

An oxidizing agent (chlorine) is one that gains electrons andis reduced to a lower valency condition.

A reducing agent (Fe2+) is one that loses electrons and isoxidized to a higher valency condition.

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• The following are examples of oxidizing agentsor oxidants which are of importance inquantitative analysis

• Examples of oxidising agents or oxidants arepotassium permanganate, potassiumdichromate, ceric sulphate, iodine, potassiumbromate and potassium iodate.

• Examples of reducing agents or reductants areFerrous sulphate, metallic iron, sodiumthiosulphate, sodium arsenite, oxalic acid,oxalates…etc.

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Equivalent weights: of an oxidant or reductant

is defined as that weight of the substance that reacts withor contains 1.008 g. of available hydrogen or 8.000 g ofavailable oxygen.

“available” is meant capable of being used in oxidation orreduction. The amount of available oxygen may beindicated by writing the hypothetical equation, e.g.,

2 K MnO4 acid K2O + 2 MnO + 5O2 K MnO4 gives up 5 atoms of available oxygen; hence its

equivalent weight is 2 K MnO4/10 or molecular weight/5.

2 K MnO4 alkaline K2O + 2 MnO2 + 3O

Therefore, equivalent weight of K MnO4 in alkaline medium= molecular weight/3.

K2 Cr2 O7 K2O + Cr2O3 + 3 OThe equivalent weight is molecular weight/6.

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Another method for calculating the equivalent weights ofoxidants and reductants is by using the

“ion-electron” equations.1. MnO4- Mn++2. To balance the equation ionically we make use of H+

(from water or acid present).3. To balance the equation electronically, 5 electrons must

be added to MnO4- sideMnO4- + 8 H+ + 5 e Mn++ + 4 H2OThe equation representing the reduction of MnO4- is now

balanced.The other partial equation representing the oxidation of

ferrous ions is as follows:Fe2+ Fe3+ balanced ionically

Fe2+ -e Fe3+ balanced electrically

The overall reaction will be:MnO4- + 8 H+ + 5 Fe2+ Mn++ +5 Fe3++ 4 H2O

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The equivalent weight of an oxidant or reductant is themolecular weight divided by the number of electronswhich one molecule of the substance gains or loses inthe reaction.

Thus equivalent Weight of MnO-4 = mol. Weight/5.& equivalent weight of Fe2+ = molecular weight/1.

“oxidation number” method, the procedure is as follows:

Equivalent weight = Molecular weight/change in theoxidation number of the element suffering oxidation orreduction. e.g.

The change in the oxidation number of the manganese isfrom +7 to +2. The equivalent weight of KMnO4 istherefore 1/5 mol.

84

62acid84

71 OSMnOMnK

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42

484

71 OMnmediumalkalineOMnK

It is clear that the equivalent weight of potassiumpermanganate in alkaline medium is 1/3 mol.

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Electrical properties of redox systems:

Electrode potential:• Suppose a metal rod dipped into a solution of one of its

salts, there is a tendency for the metal to dissolve thistendency is termed electrolytic (solution pressure). Thereverse tendency, namely, passage of metal cationsfrom the solution to be deposited on the metal is alsopossible (ionic pressure).

• In case of copper/copper sulphate system (Cu/Cu2+)the ionic pressure is greater than the solution pressure.

Cu2+ ions leave the solution to be deposited on thecopper rod. In this case, the solution acquires a negativecharge and the copper rod, a positive one (doubleelectric layer). Thus a certain potential differenceappears between the metal and the solution.

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• In case of Zn/Zn SO4 systemThe solution pressure is greater than the ionic

pressure. Zinc metal tends to dissolve formingZn2+ in solution, setting up an excess of positivecharges in the solution and of negative ones onthe metal rod (double electric layer); the netresult is also a potential difference but now it isof opposite sign.

The potential difference between the metal rod(electrode) and the solution is known as“electrode potential” abbreviated E.

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The potential difference between a metal and its ions isactually a measure of the tendency of the metal to beoxidized to metal ions or the tendency of the ions tobe reduced to metal atoms.

M - ne Mn+Mn+ + ne M

Nernest equation for electrode potential :Nernest formulated an equation relating the potential

difference-observed when an electrode is immersedin a solution of its own ions – to the concentration ofthe ions.

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Where:Et =Electrode potential at temperature tEo =A constant

dependent upon the system termed standard electrodepotential.R =Gas constant = 8.314.T =Absolutetemperature = (XoC + 273)F =Faraday = 96500coulombsLoge=Natural log. i.e. to the base 2.718 and isconverted to common log to base 10 by multiplying by2.303.n =Valency of the ions.(Mn+) =Molarconcentration of metal ions/liter.Nernest equation canbe simplified by introducing the known values of R andF, and converting the natural logarithms to base 10.

• Simply it will be

neot MLog

nFRTEE

no

o25 Mlog

n0.0591ECE

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Standard electrode potential:

Notes:• The sign of the potential is similar to the charge on the

metal electrode.• Standard electrode potential is a quantitative measure of

the readiness of the element to lose electrons (oxidized)giving its ions.

• When metals are arranged in the order of their standardelectrode potentials, the so called electrochemical seriesof the metals is obtained.

• The greater the negative value of the potential, thegreater is the tendency of the metal to pass into the ionicstate.

• A metal with a more negative potential will displace anyother metal below it in the series from its salt solution.Thus iron will displace copper or mercury from their saltsolutions.

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Standard oxidation potential:In a system containing an oxidant and its reduction product

(conjugate reductant) there will be an equilibriumrepresented as:

oxidant + ne conjugate reductante.g.Fe3+ + e Fe2+

The more powerful the oxidant, the weaker its conjugatereductant should be and vice versa.

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Notes:• The higher the standard oxidation potential of a given

system, the stronger the oxidizing power of its oxidizedform and the weaker the reducing power of its reducedform.

• The standard oxidation potentials indicate which ion willoxidize or reduce other ions at molar concentrations. Themost powerful oxidizing agents are those at the top (withhigher positive potential) and the most powerful reducingagents occupy the bottom (with higher negativepotential)

• If any two redox systems are combined, the stronger ofthe two oxidizing agents gains electrons from thestronger reducing agent with the formation of weakerreducing and oxidising agents.

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Nernest equation for oxidation potential:Nernest formulated an equation relating the oxidation

potential of the system to the concentration of bothoxidant and reductant as follows:

When [ox] = [red] then E25 = Eo = Stand. Ox. Pot.

Red

OxLogn

0.06oE25E

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Factors affecting oxidation potential(1) Common ion effect:

• The oxidation potential of MnO4-/Mn2+system,

varies with the ratio MnO4-/Mn2+ The oxidation

potential will decrease in presence of excess manganoussalt.• If Fe2+ is titrated with potassium permanganate inpresence of chloride ions; unless the oxidation potential ofMnO4

-/Mn2+ system is reduced Cl- will also be attacked bypotassium permanganate leading to higher results.• In such case, manganous sulphate in the form ofZimmermanns’ reagent is added to the solution to betitrated; permanganate is thus unable to oxidize chlorideions.

Mn

HMnOLog

50.059EE

84

o/MnMnO 4

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(2) Effect of increasing hydrogen ion concentration:• [H+] has a decided effect on the oxidation potential of

oxidizing agents containing oxygen.• The oxidation potential increases by increasing acidity

and decreases by decreasing it.

MnO4- + 4 H+ + 3 e MnO2 + 2 H2O

MnO4- + 8 H+ + 5 e Mn 2+ + 4 H2O

Cr2O72- + 14 H+ + 6 e 2 Cr 3+ + 7 H2O

2

14272

o/2CroCrCr

HOCrLog

60.059EE 2

72

33

AsO43- + 2 H+ + 2e AsO3

3- + H2O

AsO43- + 2 I- + 2 H+ AsO3

3- + I2 + H2O

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3. Effect of complexing agents:

• If HgCl2 solution is added to I2/2I- system where iodideions will be removed from the reaction as they form acomplex with mercuric ions (HgI4)2- (low dissociation).

Consequently, the oxidation potential of I2/2I- systemincreases in presence of HgCl2.

• Upon the addition of F- or PO43- to Fe3+/Fe2+ system.

Where Fe3+ ions are removed as the stable complexes(FeF6)3- or (Fe(PO4)2)3- and the oxidation potential ofFe3+/Fe2+ system is therefore reduced.• So I2 can be used to oxidise Fe2+ despite the close

oxidation potentials.I2 + Fe2+ Fe3+ + 2I-

F-

(FeF6)3-

HgCl2 (HgI4)2-

(Fe(PO4)2)3-PO43-

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4. Effect of precipitating Agents:

• The oxidation potential of Cu2+/Cu+ system is +0.15 v,therefore it is expected that cuprous compounds reduceiodine into iodide. However, Cu2+salts liberate iodine fromiodides. This is due to the low solubility of Cu2I2 (reducedform), therefore the concentration of the reduced form insolution is greatly reduced and the potential of the Cu2+/Cu+

system becomes greater than that of I2/2I-. A large excess ofpotassium iodide is needed to make this reversible reactionproceed quantitatively.

2 Cu2++ 4 I- Cu2I2 + I2However, in presence of much tartarate or citrate ions (that

form stable complex with cupric ions) iodine can oxidizequantitatively cuprous compounds.

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Redox titration curvesplot of ml of titrant against the potential E (volts)

A titrant such as 0.1 N ceric sulphate to 100ml of 0.1 Nsolution of ferrous sulphate, the change in potentialduring the titration can be either measured or calculatedusing Nernest equation as follows:

Upon adding 10ml of ceric sulphate, the ratio of(Fe3+)/(Fe2+) becomes 10/90

and When 50 ml of oxidant are added.

9010Log

10.059EE o

0.69v9010Log

10.0590.77

0.77V5050Log

10.0590.77E

With 90 ml titrantE = 0.8199 ml titrantE = 0.87 v

99.9 ml titrantE = 0.93 v

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At the equivalence pointwhen equilibrium is established,E (Fe3+ / Fe2+), will be equal to E (Ce4+/Ce3+) System.The potential at the equivalence point can be calculated

using either of the 2 half reactions:E = E1o + 0.059 Log [Fe3+] / [Fe2+]E = E2o + 0.059 Log [Ce4+]/ [Ce3+]

At the equivalence point the two potentials are identical.Moreover, [Fe2+] = [Ce4+] & [Fe3+] = [Ce3+]

The two equations are added:

2 E = E1o + E2o + 0.059 log

32

43

CeFe

CeFe

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E = E1o + E2oE = E1o +E2o / 2More generally:E ep = n1E1o + n2E2o / n1+n2

When 100 ml is added (equivalence point)E1o +E2o / 2 = 0.77 + 1.45 / 2 = 1.10 v

Addition of more of the oxidant beyond the equivalencepoint increases the ratio (Ce4+) / (Ce3+) with 100.1 mlWith 101 ml titrant E = 1.33 v110 ml titrant E = 1.39 v

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Inflection depends on:1- The difference between Eo of the two systems involved.2- Effect of dilution.3- Effect of addition of complexing or precipitating agents.

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Detection of the end point in redox titrations

1. No indicators• purple violet color of permanganate disappears owing

to reduction to the almost colorless Mn2+. When all thereducing agent has been oxidized a single excess dropof permanganate colors the whole solution a distinctpink.

• titrations with iodine solution may be performed withoutthe use of indicators because the dark brown color ofiodine disappears as a result of its reduction to iodideions. However, since the color of iodine solutions is notvery deep, titrations with iodine solution are best donein presence of an indicator-starch, which gives anintense blue color even with very small amounts of freeiodine.

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2. External indicators

Examples• The spot test method for the titration of ferrous

iron with potassium dichromate.Near the equivalence point, drops of solutionare removed and brought into contact withdilute freshly prepared potassium ferricyanidesolution on a spot plate. The end point isreached when the drop first fails to give a bluecolor.

• Titration of zinc ions with standard potassiumferrocyanide solution; here a solution of Uranylacetate or nitrate is the external indicator, andtitration is continued until a drop of the solutionjust imparts a brown color to the indicator.

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3. Internal Redox Indicators

Diphenylamine is a redox indicator(Eo = +0.76 v and n =2)the range of diphenylamine is 0.73. v - 0.79 v.At potential below 0.73 v the color of the reduced

form prevails (colorless).At E = 0.79 V the color of the oxidized form

predominates (blue-violet.Between 0.73 and 0.79 V the color of the solution

changes gradually from colorless to blue-violet.

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N

H

N

H

N

H

N

H

N

H

+ 2 H+ + 2e

+ +

Diphenylamine (colorless) Diphenylbenzidine

(colorless)

Diphenylbenzidine (violet)

+ 2 e

2

• The oxidation potential of a redox indicator should beintermediate between that of the solution titrated and thatof the titrant.• The range in which the indicator changes color must bewithin the limits of the sharp change of potential on thetitration curve so that the indicator error in titration shouldbe as small as possible.• 1% solution of diphenylamine in conc. H2SO4 is used asindicator.

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• Diphenylamine is unsuitable indicator for the titration offerrous iron with permanganate (potential break is from0.94 V- 1.47 V)

• It is also unsuitable for titration of Fe2+ with dichromate(potential break 0.94 V- 1.30 V) as the indicator color willchange when only about 50% of ferrous ions has beenoxidized (E = Eo Fe3+/Fe2+ = 0.77 V). If, however, ferricions are complexed by the addition of phosphate ions, itis then possible to lower the potential at which thechange begins. In presence of phosphate ions the colorchange of diphenlyamine is within the range of potentialbreak and diphenylamine is then quite suitable asindicator for titration of ferrous ions.

That is to make indicator potential falls betweenpotential break of oxidizing& reducing forms

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Dichromate can be titrated with ferrous iron even withoutaddition of phosphate in presence of redox indicatorswith higher values of Eo e.g. phenylanthranilic acid,Eo=1.08 V.Chelate of ferrous iron with 1.10 orthophenanthroline(ferroin) is intensely red and is converted by oxidationinto the pale blue ferric complex (ferriin):

N

N

N

N+ F e 2+3

F e 3+

31 ,1 0 - o rth o P h e n a n th ro lin e Iro n ( II) 1 ,1 0 - o rth o

P h e n a n th ro lin e

It is an excellent indicator for Ce4+. It has high Eo whichis affected by acidity. The only disadvantage of thisindicator is that it is somewhat expensive.

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ApplicationA. Fe2+

1. Fe2+ can be directly titrated in presence of dilutesulphuric acid with standard potassium permanganate. Ifsample was FeCl2, then one have to add Zimmermanreagent

2. Fe2+ can be directly titrated with standard dichromatesolution either in presence of diphenylamine indicator orby the use of potassium ferricyanide as externalindicator. If diphenylamine is used as internal redoxindicator, phosphoric acid must be added in order tolower the oxidation potential of Fe3+/Fe2+ system

3. Fe2+ can be directly titrated with ceric sulphate solutiontill the solution acquires a pale yellow color (that ofexcess titrant, self indicator). Methyl red can also beused as irreversible redox indicator till the red color ofthe indicator in acid solution is bleached or changed toyellow.

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B. Ferric Iron:Ferric ions can be directly titrated with titanous chloride

solution.The end point is detected by the use of methylene blue

which is reduced to the leuco-compound (colorless) bythe first excess of the titrant.Fe3+ + Ti 3+ Fe2+ + Ti4+

SCN- can also be used as indicator; the solution remainsred as long as Fe3+ ions are present.

Several methods are used for the reduction of Fe3+ toFe2+. The produced ferrous salt can be determined asunder ferrous iron.

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a. Reduction with stannous chloride:SnCl2 + HgCl2 Hg2Cl2 + SnCl4

b. Reduction with zinc metal and sulphuric acid:With granulated zinc in acid medium (slow reaction of with

is accelerated by the addition of few drops of coppersulphate solution) by gentle boiling. During boiling theflask should be fitted with a Bunsen valve to prevent theentrance of air i.e. prevents oxidation with atmosphericoxygen.

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c. Reduction with Amalgamated zinc:The amalgamated zinc (obtained by

treating zinc with mercuric chloridesolution) is an excellent reducing agent.

2 Fe3+ + Zno 2 Fe2+ + Zn2+

Reduction is done by passing the coldacidified solution of ferric salts through acolumn of amalgamated zinc (Jone’sreductor). The column is then washedwith 2.5% H2SO4 followed by water, thewashings are collected with the reducediron solution

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Notes:

• Reducing substances capable of reducing Fe3+ to Fe2+are determined by treating with an excess of ferric saltsolution, the produced Fe2+ is titrated with standardKMnO4 as before e.g. zinc powder, metallic iron.

Zno + 2 Fe3+ Zn2+ + 2Fe2+Feo + 2 Fe3+ 3 Fe2+Zinc oxide and iron oxides do not interfere.

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• Oxidizing substances capable of oxidizing Fe2+ to Fe3+are allowed to react with a known excess of standardferrous sulphate solution. The residual Fe2+ is thentitrated with standard permanganate.e.g. potassium persulphate, chlorate, manganese dioxide.

S2O82- + 2 Fe2+ 2 SO42- + 2 Fe3+ClO3- + 6 Fe2+ + 6 H+ Cl- + 6 Fe3+ + 3 H2OMnO2 + 2 Fe2+ + 4 H+ Mn2+ + 2 Fe3+ + 2 H2O

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C O O H

C O O H + 2 H 2 O + O 2 C O 2 + 3 H 2 O

C. Determination of oxalates:

Oxalic acid and oxalates are strong reducing agents i.e. can betitrated directly with permanganate or ceric solutions.

Titration with KMnO4 is done at 60o C in presence of dilutesulphuric acid. The reaction is slow at the beginning but once smallamount of Mn2+ is formed the reaction becomes very rapid.

2 MnO4- + 5 C2O4

2- + 16 H+ 2 Mn2+ + 10 CO2 + 8 H2O

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E. Determination of Metallic Iron in presence of Iron Oxide:

Fe/Fe2+ system has more negative electrode potential thanCu/Cu2+ or Hg/Hg2+ systems. Therefore metallic iron can displacecopper or mercury from their salt solutions.

Reduced iron is prepared by reducing ferric oxide by hydrogen.It can be assayed by shaking a known weight of reduced iron witheither mercuric chloride or copper sulphate solution on hot.

The produced ferrous salt is titrated with potassiumpermanganate. If mercuric chloride solution is used Zimmermann’sreagent should be added to avoid interference of chloride.

Fe + HgCl2 Hg + FeCl2Fe + CuSO4 Cu + FeSO4

Iron oxide does not interfere.

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F. Determination of Hydrogen peroxide:

The determination is based on the reaction of H2O2 with potassiumpermanganate:

5 H2O2 + 2 MnO4- + 6 H+ 5O2 + 2Mn2+ + 8H2O

The equation shows that in this reaction H2O2 acts as a reducingagent and is oxidized to oxygen.

H2O2 – 2e O2 + 2 H+

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Properties of some oxidizing agents:

Potassium permanganate: K MnO4 (Secondary standard substance

The solution must then be standardized using reducing agentse.g.a) iron wire: The A.R. quality is used. It is dissolved in H2SO4giving FeSO4 titrated with permanganate.b) Sodium oxalate: This is dissolved in sulphuric acid, thesolution is then titrated with KMnO4 at 70oC. The reaction iscatalyzed by the produced Mn2+ (autocatalytic):2 MnO4- + 5 H2C2O4 + 6H+ Mn2+ + 10 CO2 + 8H2Oc) Arsenious trioxide: This is dissolved in sodium hydroxidefollowed by acidification to 0.5 N with HCl, a drop of KI solution isadded as a catalyst (forming (ICl2-)) the solution is then titrated inthe cold with K MnO4:HAsO2

-+2 ICl2- + 2H2O H3AsO4 + I2 + 2H+ + 4 Cl-2 MnO4

- + 5 I2 + 20 Cl- +16 H+ 2Mn2+ + 10 ICl2- + 8 H2O

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Potassium Dichromate: K2Cr2O7

This is a strong oxidizing agent but of limited use. It is obtainable inhigh purity, thus it is a primary standard and its solution is stable.Another advantage is that it does not oxidize HCl. It can not oxidizeoxalic acid or ferrocyanide. Its main application is the determinationof ferrous ion. It cannot serve as a self indicator; the orange Cr2O7

2-

is reduced to the green Cr3+ ion.

Ceric (Ce4+).It has the following half reaction:Ce4+ + e Ce3+

Yellow orange colorlessAlthough it can be used as self indicator it is better to use a redoxindicator e.g. ferroin.Acid medium is needed to prevent the precipitation of CeO2.

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Redox reactions involving I2/2I- system.

The standard oxidation potential of I2/2I- system has the relativelylow value of +0.54 V. The fact that the system is about half –waydown the table of oxidation potentials shows that:

(a) There are several reducing agents which can be oxidized withfree iodine i.e. those having Eo < + 0.54 V e.g. :

Sn4+/ Sn2+ (Eo = +0.15 V),S4O6

2-/S2O32- (Eo = -0.08 V),

S/S2- (Eo = -0.55 V).

Sn2+ + I2 Sn4+ + 2I-2 S2O3

2- + I2 S4O62- + 2 I-

S2- + I2 So ppt + 2 I-This is the basis of iodimetric methods of analysis.

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(b) There is also a number of oxidizing agents which can be reducedby iodide ions i.e. having Eo> + 0.54 V e.g.:MnO4

- / Mn2+ (Eo = +1.5 V),Cr2O7

2-/2 Cr3+ (Eo = 1.3 V),ClO3

-/Cl – (Eo = + 1.45 V).

2MnO4- + 10 I- + 16 H+ 5 I2 + 2 Mn2+ + 8 H2O

ClO3- + 6 I- + 6 H+ 3 I2 + Cl- +3 H2O.

Cr2O72- + 14 H+ + 6 I- 2 Cr3+ + 7 H2O +3 I2

This involves iodometric methods of analysis.

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(c) Systems having oxidation potentials near to that of iodine/iodidesystem such as:AsO4

3-/AsO33- (Eo = + 0.57 V),

Fe3+/Fe2+(Eo=0.76V).Their reaction with iodine tends to go in the reverse direction and isdirected forward or backward by control of experimental conditions(affecting oxidation potential).

AsO33- + I2 + H2O AsO4

3- + 2 HIIf H+ ions are removed by the addition of AsO4

3-/AsO33- system is

lowered so that arsenite can be oxidized quantitatively by iodine. Ifhowever (H+) is increased, the oxidation potential of AsO4

3-/AsO33-

system increases and arsenate oxidizes iodide in presence of muchacid.

2Fe3+ + 2 I- 2 Fe2+ + I2Ferric ion oxidizes iodide quantitatively only in the presence of highiodide concentration (when the oxidation potential of I2/2I- system isdecreased). Iodine solution can oxidize Fe2+ salts in presence ofPO4

3- or F- that form stable complexes with Fe3+ lowering thereforethe oxidation potential of Fe3+/Fe2+ system.

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Cu2+/Cu+ system is +0.15 v,therefore it is expected that cuprous compounds reduce iodine intoiodide. However, Cu2+ salts liberate iodine from iodides. This is dueto the low solubility of cuprous iodide (reduced form), therefore theconcentration of the reduced form in solution is greatly reduced andthe potential of the Cu2+/Cu+ system becomes greater than that ofI2/2I-. A large excess of potassium iodide is needed to make thisreversible reaction proceed quantitatively.

2 Cu2+ + 4 I- Cu2I2 ppt + I2

However, in presence of much tartarate or citrate ions (that formstable complex with cupric ions) iodine can oxidize quantitativelycuprous compounds.