Analytical Chem .2
Course Code 306 for CLINICAL
2 + 1 Hours
١
Team for General
Team for Clinical
Prof.Dr Mostafa Abdel-Aty
Prof.Dr Mostafa Abdel-Aty
Prof.Dr. Abdel-Aziz Elbayoumy
Dr Safaa Reyad
Dr. Nesreen Khamees
Dr Ali Yahia
Dr. Mohamad Khaled
Dr. Nesreen Talaat
Subjects
Subjects
A- Redox
-Redox
B- Electrochemistry
-Complexometry
1-Conductometry
2-Potentiometry
Reference: Vogel’s textbook for quantitative chemical analysis
3-Polarography
- Harris Quantitative Chemical Analysis
- Day and Underwood quantitative Chemical Analysis
٢
١
- Skoog Quantitative Pharmaceutical Analysis
OXIDATION-REDUCTION
TITRATIONS
REDOX
٣
REDOX
In the past, oxidation was defined as the reaction of
compounds with oxygen while reduction as their reaction
with hydrogen.
But, no oxygen is involved in the oxidation of ferrous chloride
by chlorine gas according to the following equation:
2 Fe2+ + Cl2 
2 Fe3+ + 2 ClOxidation is the process, which results in the loss of one or
more electrons by atoms or ions.
Reduction is the process, which results in the gain of one or
more electrons by atoms or ions.
An oxidizing agent (chlorine) is one that gains electrons and
is reduced to a lower valency condition.
A reducing agent (Fe2+) is one that loses electrons and is
oxidized to a higher valency condition.
٤
٢
• The following are examples of oxidizing agents
or oxidants which are of importance in
quantitative analysis
• Examples of oxidising agents or oxidants are
potassium permanganate, potassium
dichromate, ceric sulphate, iodine, potassium
bromate and potassium iodate.
• Examples of reducing agents or reductants are
Ferrous sulphate, metallic iron, sodium
thiosulphate, sodium arsenite, oxalic acid,
oxalates…etc.
٥
Equivalent weights: of an oxidant or reductant
is defined as that weight of the substance that reacts with
or contains 1.008 g. of available hydrogen or 8.000 g of
available oxygen.
“available” is meant capable of being used in oxidation or
reduction. The amount of available oxygen may be
indicated by writing the hypothetical equation, e.g.,
2 K MnO4
acid
K2O + 2 MnO + 5O
2 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 O
The equivalent weight is molecular weight/6.
٦
٣
Another method for calculating the equivalent weights of
oxidants 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- side
MnO4- + 8 H+ + 5 e  Mn++ + 4 H2O
The 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
٧
The equivalent weight of an oxidant or reductant is the
molecular weight divided by the number of electrons
which one molecule of the substance gains or loses in
the 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 the
oxidation number of the element suffering oxidation or
reduction. e.g.
acid
K 1Mn 7O 4  8   Mn 2S  6O 4  8
The change in the oxidation number of the manganese is
from +7 to +2. The equivalent weight of KMnO4 is
therefore 1/5 mol.
٨
٤
K 1Mn 7O 4  8
alkaline
Mn 4 O2  4
medium
It is clear that the equivalent weight of potassium
permanganate in alkaline medium is 1/3 mol.
٩
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 this
tendency is termed electrolytic (solution pressure). The
reverse tendency, namely, passage of metal cations
from the solution to be deposited on the metal is also
possible (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 the
copper rod. In this case, the solution acquires a negative
charge and the copper rod, a positive one (double
electric layer). Thus a certain potential difference
appears between the metal and the solution.
١٠
٥
• In case of Zn/Zn SO4 system
The solution pressure is greater than the ionic
pressure. Zinc metal tends to dissolve forming
Zn2+ in solution, setting up an excess of positive
charges in the solution and of negative ones on
the metal rod (double electric layer); the net
result is also a potential difference but now it is
of opposite sign.
The potential difference between the metal rod
(electrode) and the solution is known as
“electrode potential” abbreviated E.
١١
The potential difference between a metal and its ions is
actually a measure of the tendency of the metal to be
oxidized to metal ions or the tendency of the ions to
be 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 immersed
in a solution of its own ions – to the concentration of
the ions.
١٢
٦
Et  Eo 
RT
nF
Log
M
n

e
Where:
Et =Electrode potential at temperature tEo =A constant
dependent upon the system termed standard electrode
potential.R =Gas constant = 8.314.T =Absolute
temperature = (XoC + 273)F =Faraday = 96500
coulombsLoge=Natural log. i.e. to the base 2.718 and is
converted to common log to base 10 by multiplying by
2.303.n =Valency of the ions.(Mn+) =Molar
concentration of metal ions/liter.Nernest equation can
be simplified by introducing the known values of R and
F, and converting the natural logarithms to base 10.
• Simply it will be
E25 o C  Eo 
 
0.0591
log Mn 
n
١٣
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 standard
electrode potentials, the so called electrochemical series
of the metals is obtained.
• The greater the negative value of the potential, the
greater is the tendency of the metal to pass into the ionic
state.
• A metal with a more negative potential will displace any
other metal below it in the series from its salt solution.
Thus iron will displace copper or mercury from their salt
solutions.
١٤
٧
Standard oxidation potential:
In a system containing an oxidant and its reduction product
(conjugate reductant) there will be an equilibrium
represented as:
oxidant + ne
e.g.
Fe3+ + e
conjugate reductant
Fe2+
The more powerful the oxidant, the weaker its conjugate
reductant should be and vice versa.
١٥
Notes:
• The higher the standard oxidation potential of a given
system, the stronger the oxidizing power of its oxidized
form and the weaker the reducing power of its reduced
form.
• The standard oxidation potentials indicate which ion will
oxidize or reduce other ions at molar concentrations. The
most powerful oxidizing agents are those at the top (with
higher positive potential) and the most powerful reducing
agents occupy the bottom (with higher negative
potential)
• If any two redox systems are combined, the stronger of
the two oxidizing agents gains electrons from the
stronger reducing agent with the formation of weaker
reducing and oxidising agents.
١٦
٨
Nernest equation for oxidation potential:
Nernest formulated an equation relating the oxidation
potential of the system to the concentration of both
oxidant and reductant as follows:
0.06
E25  E o 
Log
n
 Ox 
 Red
When [ox] = [red] then E25 = Eo = Stand. Ox. Pot.
١٧
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 manganous
salt.
• If Fe2+ is titrated with potassium permanganate in
presence of chloride ions; unless the oxidation potential of
MnO4-/Mn2+ system is reduced Cl- will also be attacked by
potassium permanganate leading to higher results.
• In such case, manganous sulphate in the form of
Zimmermanns’ reagent is added to the solution to be
titrated; permanganate is thus unable to oxidize chloride
ions.
8
١٨
٩
E
MnO 4  /Mn  
 Eo 
 

MnO 4 H 
0.059
Log
5
Mn  

(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
MnO4- + 8 H+ + 5 e
MnO2 + 2 H2O
Mn 2+ + 4 H2O
Cr2O72- + 14 H+ + 6 e
2 Cr 3+ + 7 H2O
E
Cr 2 o 7
2
/2Cr
3
 Eo 

 
 
Cr 2 O 7 2  H 
0.059
Log
2
6
Cr 3 
AsO43- + 2 H+ + 2e
AsO33- + H2O
AsO43- + 2 I- + 2 H+
AsO33- + I2 + H2O
14
١٩
3. Effect of complexing agents:
• If HgCl2 solution is added to I2/2I- system where iodide
ions will be removed from the reaction as they form a
complex with mercuric ions (HgI4)2- (low dissociation).
Consequently, the oxidation potential of I2/2I- system
increases 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 of
Fe3+/Fe2+ system is therefore reduced.
• So I2 can be used to oxidise Fe2+ despite the close
oxidation potentials.
HgCl2
(HgI4)2I2 + Fe2+
Fe3+ + 2IF-
٢٠
(FeF6)3-
١٠
PO43-
(Fe(PO4)2)3-
4. Effect of precipitating Agents:
• The oxidation potential of Cu2+/Cu+ system is +0.15 v,
therefore it is expected that cuprous compounds reduce
iodine into iodide. However, Cu2+salts liberate iodine from
iodides. This is due to the low solubility of Cu2I2 (reduced
form), therefore the concentration of the reduced form in
solution is greatly reduced and the potential of the Cu 2+/Cu+
system becomes greater than that of I2/2I-. A large excess of
potassium iodide is needed to make this reversible reaction
proceed quantitatively.
2 Cu2++ 4 I-  Cu2I2  + I2
However, in presence of much tartarate or citrate ions (that
form stable complex with cupric ions) iodine can oxidize
quantitatively cuprous compounds.
٢١
Redox titration curves
plot of ml of titrant against the potential E (volts)
A titrant such as 0.1 N ceric sulphate to 100ml of 0.1 N
solution of ferrous sulphate, the change in potential
during the titration can be either measured or calculated
using Nernest equation as follows:
Upon adding 10ml of ceric sulphate, the ratio of
(Fe3+)/(Fe2+) becomes 10/90
0.059
1
E  E o

 0.77
0.059
1

Log
Log
10
90
10
 0.69v
90
and When 50 ml of oxidant are added.
E  0.77 
٢٢
١١
E = 0.81
E = 0.87 v
E = 0.93 v
0.059
50
Log
 0.77V
1
50
With 90 ml titrant
99 ml titrant
99.9 ml titrant
At the equivalence point
when 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
Fe3 Ce4 
Fe2 Ce3 
٢٣
E = E1o + E2o
E = E1o +E2o / 2
More 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 equivalence
point increases the ratio (Ce4+) / (Ce3+) with 100.1 ml
With 101 ml titrant E = 1.33 v
110 ml titrant E = 1.39 v
٢٤
١٢
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.
٢٥
Detection of the end point in redox titrations
1.
•
•
٢٦
١٣
No indicators
purple violet color of permanganate disappears owing
to reduction to the almost colorless Mn2+. When all the
reducing agent has been oxidized a single excess drop
of permanganate colors the whole solution a distinct
pink.
titrations with iodine solution may be performed without
the use of indicators because the dark brown color of
iodine disappears as a result of its reduction to iodide
ions. However, since the color of iodine solutions is not
very deep, titrations with iodine solution are best done
in presence of an indicator-starch, which gives an
intense blue color even with very small amounts of free
iodine.
2. External indicators
Examples
• The spot test method for the titration of ferrous
iron with potassium dichromate.
Near the equivalence point, drops of solution
are removed and brought into contact with
dilute freshly prepared potassium ferricyanide
solution on a spot plate. The end point is
reached when the drop first fails to give a blue
color.
• Titration of zinc ions with standard potassium
ferrocyanide solution; here a solution of Uranyl
acetate or nitrate is the external indicator, and
titration is continued until a drop of the solution
just imparts a brown color to the indicator.
٢٧
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.
٢٨
١٤
2
H
H
H
N
N
N
Diphenylamine
(colorless)
+ 2 H+ +
2e
Diphenylbenzidine
(colorless)
H
H
N
+
N
+
+
2e
Diphenylbenzidine
(violet)
• The oxidation potential of a redox indicator should be
intermediate between that of the solution titrated and that
of the titrant.
• The range in which the indicator changes color must be
within the limits of the sharp change of potential on the
titration curve so that the indicator error in titration should
be as small as possible.
• 1% solution of diphenylamine in conc. H2SO4 is used as
٢٩
indicator.
• Diphenylamine is unsuitable indicator for the titration of
ferrous iron with permanganate (potential break is from
0.94 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 will
change when only about 50% of ferrous ions has been
oxidized (E = Eo Fe3+/Fe2+ = 0.77 V). If, however, ferric
ions are complexed by the addition of phosphate ions, it
is then possible to lower the potential at which the
change begins. In presence of phosphate ions the color
change of diphenlyamine is within the range of potential
break and diphenylamine is then quite suitable as
indicator for titration of ferrous ions.
That is to make indicator potential falls between
potential break of oxidizing& reducing forms
٣٠
١٥
Dichromate can be titrated with ferrous iron even without
addition of phosphate in presence of redox indicators
with 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 oxidation
into the pale blue ferric complex (ferriin):
N
3
N
N
+
F e 2+
N
1 ,1 0 - o r t h o P h e n a n t h r o lin e
F e 3+
3
Ir o n ( II) 1 ,1 0 - o r t h o
P h e n a n t h r o lin e
It is an excellent indicator for Ce4+. It has high Eo which
is affected by acidity. The only disadvantage of this
indicator is that it is somewhat expensive.
٣١
Application
A.
1.
2.
3.
٣٢
١٦
Fe2+
Fe2+ can be directly titrated in presence of dilute
sulphuric acid with standard potassium permanganate. If
sample was FeCl2, then one have to add Zimmerman
reagent
Fe2+ can be directly titrated with standard dichromate
solution either in presence of diphenylamine indicator or
by the use of potassium ferricyanide as external
indicator. If diphenylamine is used as internal redox
indicator, phosphoric acid must be added in order to
lower the oxidation potential of Fe3+/Fe2+ system
Fe2+ can be directly titrated with ceric sulphate solution
till the solution acquires a pale yellow color (that of
excess titrant, self indicator). Methyl red can also be
used as irreversible redox indicator till the red color of
the indicator in acid solution is bleached or changed to
yellow.
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) by
the first excess of the titrant.
Fe3+ + Ti 3+  Fe2+ + Ti4+
SCN- can also be used as indicator; the solution remains
red as long as Fe3+ ions are present.
Several methods are used for the reduction of Fe3+ to
Fe2+. The produced ferrous salt can be determined as
under ferrous iron.
٣٣
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 copper
sulphate solution) by gentle boiling. During boiling the
flask should be fitted with a Bunsen valve to prevent the
entrance of air i.e. prevents oxidation with atmospheric
oxygen.
٣٤
١٧
c. Reduction with Amalgamated zinc:
The amalgamated zinc (obtained by
treating zinc with mercuric chloride
solution) is an excellent reducing agent.
2 Fe3+ + Zno
2 Fe2+ + Zn2+
Reduction is done by passing the cold
acidified solution of ferric salts through a
column of amalgamated zinc (Jone’s
reductor). The column is then washed
with 2.5% H2SO4 followed by water, the
washings are collected with the reduced
iron solution
٣٥
Notes:
• Reducing substances capable of reducing Fe3+ to Fe2+
are determined by treating with an excess of ferric salt
solution, the produced Fe2+ is titrated with standard
KMnO4 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.
٣٦
١٨
• Oxidizing substances capable of oxidizing Fe2+ to Fe3+
are allowed to react with a known excess of standard
ferrous sulphate solution. The residual Fe2+ is then
titrated with standard permanganate.
e.g. potassium persulphate, chlorate, manganese dioxide.
S2O82- + 2 Fe2+
ClO3- + 6 Fe2+ + 6 H+
MnO2 + 2 Fe2+ + 4 H+
2 SO42- + 2 Fe3+
Cl- + 6 Fe3+ + 3 H2O
Mn2+ + 2 Fe3+ + 2 H2O
٣٧
C. Determination of oxalates:
Oxalic acid and oxalates are strong reducing agents i.e. can be
titrated directly with permanganate or ceric solutions.
C O O H
C O O H
+
2
H
2
O
+
O
2
C O
2
+
3
H
2
O
Titration with KMnO4 is done at 60o C in presence of dilute
sulphuric acid. The reaction is slow at the beginning but once small
amount of Mn2+ is formed the reaction becomes very rapid.
2 MnO4- + 5 C2O42- + 16 H+
٣٨
١٩
2 Mn2+ + 10 CO2 + 8 H2O
E. Determination of Metallic Iron in presence of Iron Oxide:
Fe/Fe2+ system has more negative electrode potential than
Cu/Cu2+ or Hg/Hg2+ systems. Therefore metallic iron can displace
copper 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 with
either mercuric chloride or copper sulphate solution on hot.
The produced ferrous salt is titrated with potassium
permanganate. If mercuric chloride solution is used Zimmermann’s
reagent should be added to avoid interference of chloride.
Fe + HgCl2
Hg + FeCl2
Fe + CuSO4
Cu + FeSO4
Iron oxide does not interfere.
٣٩
F. Determination of Hydrogen peroxide:
The determination is based on the reaction of H2O2 with potassium
permanganate:
5 H2O2 + 2 MnO4- + 6 H+
5O2 + 2Mn2+ + 8H2O
The equation shows that in this reaction H2O2 acts as a reducing
agent and is oxidized to oxygen.
H2O2 – 2e
٤٠
٢٠
O 2 + 2 H+
Properties of some oxidizing agents:
Potassium permanganate: K MnO4 (Secondary standard substance
The solution must then be standardized using reducing agents
e.g.
a) iron wire: The A.R. quality is used. It is dissolved in H2SO4
giving FeSO4 titrated with permanganate.
b) Sodium oxalate: This is dissolved in sulphuric acid, the
solution is then titrated with KMnO4 at 70oC. The reaction is
catalyzed by the produced Mn2+ (autocatalytic):
2 MnO4- + 5 H2C2O4 + 6H+
Mn2+ + 10 CO2 + 8H2O
c) Arsenious trioxide: This is dissolved in sodium hydroxide
followed by acidification to 0.5 N with HCl, a drop of KI solution is
added as a catalyst (forming (ICl2-)) the solution is then titrated in
the 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
٤١
Potassium Dichromate: K2Cr2O7
This is a strong oxidizing agent but of limited use. It is
obtainable in high purity, thus it is a primary standard and its
solution is stable. Another advantage is that it does not oxidize HCl.
It can not oxidize oxalic acid or ferrocyanide. Its main application is
the determination of ferrous ion. It cannot serve as a self indicator;
the orange Cr2O72- is reduced to the green Cr3+ ion.
Ceric (Ce4+).
It has the following half reaction:
Ce4+ + e
Ce3+
Yellow orange
colorless
Although it can be used as self indicator it is better to use a
redox indicator e.g. ferroin.
Acid medium is needed to prevent the precipitation of CeO2.
٤٢
٢١
Redox reactions involving I2/2I- system.
The standard oxidation potential of I2/2I- system has the relatively
low value of +0.54 V. The fact that the system is about half –way
down the table of oxidation potentials shows that:
(a) There are several reducing agents which can be oxidized with
free iodine i.e. those having Eo < + 0.54 V e.g. :
Sn4+/ Sn2+ (Eo = +0.15 V),
S4O62-/S2O32- (Eo = -0.08 V),
S/S2- (Eo = -0.55 V).
Sn2+ + I2
Sn4+ + 2I2 S2O32- + I2
S4O62- + 2 IS2- + I2
So ppt + 2 IThis is the basis of iodimetric methods of analysis.
٤٣
(b) There is also a number of oxidizing agents which can be
reduced by iodide ions i.e. having Eo> + 0.54 V e.g.:
MnO4- / Mn2+ (Eo = +1.5 V),
Cr2O72-/2 Cr3+ (Eo = 1.3 V),
ClO3-/Cl – (Eo = + 1.45 V).
2MnO4- + 10 I- + 16 H+
ClO3- + 6 I- + 6 H+
Cr2O72- + 14 H+ + 6 I-
5 I2 + 2 Mn2+ + 8 H2O
3 I2 + Cl- +3 H2O.
2 Cr3+ + 7 H2O +3 I2
This involves iodometric methods of analysis.
٤٤
٢٢
(c) Systems having oxidation potentials near to that of
iodine/iodide system such as:
AsO43-/AsO33- (Eo = + 0.57 V),
Their reaction with iodine tends to go in the reverse
direction and is directed forward or backward by control
of experimental conditions (affecting oxidation potential).
AsO33- + I2 + H2O
AsO43- + 2 HI
+
If H ions are removed by the addition of CO32- …
AsO43-/AsO33- system is lowered so that arsenite can be
oxidized quantitatively by iodine. If however (H+) is
increased, the oxidation potential of AsO43-/AsO33system increases and arsenate oxidizes iodide in
presence of much acid.
٤٥
Fe3+/Fe2+ (Eo=0.76V).
٤٦
٢٣
2Fe3+ + 2 I2 Fe2+ + I2
Ferric ion oxidizes iodide quantitatively only in the presence of
high iodide concentration (when the oxidation potential of I2/2Isystem is decreased). Iodine solution can oxidize Fe2+ salts in
presence of PO43- or F- that form stable complexes with Fe3+
lowering therefore the oxidation potential of Fe3+/Fe2+ system.
Cu2+/Cu+ system is +0.15 v,
therefore it is expected that cuprous compounds reduce iodine
into iodide. However, Cu2+ salts liberate iodine from iodides. This is
due to the low solubility of cuprous iodide (reduced form), therefore
the concentration of the reduced form in solution is greatly reduced
and the potential of the Cu2+/Cu+ system becomes greater than that
of I2/2I-. A large excess of potassium iodide is needed to make this
reversible reaction proceed quantitatively.
2 Cu2+ + 4 ICu2I2 ppt + I2
However, in presence of much tartarate or citrate ions (that form
stable complex with cupric ions) iodine can oxidize quantitatively
cuprous compounds.
Effect of increasing (OH-):
In strongly alkaline solutions iodine reacts with alkalies
according to the equation.
I2 + 2 OHIO- + I- + H2O
Hypoiodite ion (IO-) is stronger oxidant than iodine and
partially oxidizes thiosulphate to sulphate
S2O32-+4 IO- + 2 OH4I- + 2 SO42- +H2O
By the use of iodine in alkaline medium many mild
oxidations can be done in organic compounds
e.g. oxidation of glucose by IO-.
RCHO + IORCOOH + IThe excess IO- is converted to iodine by acidification
with dilute acid; titration of the excess iodine is carried out
as usual.
IO- + I- + 2H+
H2O + I2
٤٧
End Point Detection
a. Starch solution:
Starch solution forms an intense blue adsorption
compound with iodine.
1-2 ml of 1% aqueous solution of starch is added to
each 100 ml of the titrated solution.
The starch solution must be added near the end of the
titration when the titrated solution has a faint straw-yellow
color. If the starch is added earlier when there is still much
iodine in solution, the large amount of the iodine-starch
compound formed reacts slowly with the thiosulphate, so
that it is easy to add too much thiosulphate.
Starch can not be employed in alcoholic solution (that
hinders adsorbate formation), nor in strongly acidic
solutions (that hydrolyze starch and destroy the adsorbate).
٤٨
٢٤
b. Chloroform or carbon tetrachloride:
In alcoholic or strongly acidic solutions the end point is
detected by the use of either chloroform or carbon
tetrachloride. The solubility of iodine in chloroform is about
90 times greater than its solubility in water. When
chloroform is added to aqueous iodine solution, most of
iodine will dissolve – upon shaking- in the organic layer
that settles down and is colored deep violet.
٤٩
Applications
1- Iodimetry
A. Determination of sulphides, sulphites and thiosulphates:
These can be oxidized with iodine as follows:
S2- + I2
So + 2I2SO3 + I2 + H2O
SO42- + 2 HI
22 S2O3 + I2
S4O62- + 2I- Aqueous solution of sulphide or sulphite:
treated with excess standard iodine, the solution is then acidified
with 6 N HCl, mixed well, the residual iodine is then back titrated with
thiosulphate solution using starch as indicator. The solution must not
react basic as IO- may oxidize some sulphide ions into sulphate.
- Thiosulphate can be directly titrated
- Acid insoluble sulphids e.g. CuS, HgS can be decomposed by
treatment with acid and zinc.
CuS + Zno + 2 H+
Cuo + Zn2+ + H2S
The liberated H2S is absorbed in a measured excess of standard
iodine solution, the excess is then back titrated with std thiosulphste
soln..
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٢٥
2- Iodometry (for deten of oxidising agents)
a) Deten of free halogens
Chlorine water and chlorate
Cl2 + H2O
HOCl + HCl
When treated with KI both Cl2 and HOCl liberate iodine that can be
titrated with thiosulphate
ClO3- + 6I- + 6H+
I2 + Cl- + 3 H2O
b) Deten Cu2+ and Fe3+
2Cu2+ + 4ICu2I2 + I2
2Fe3+ + 2I2 Fe2+ + I2 (use Chloroform as ind.in acid med.)
Liberated I2 is titrated with std Na2S2O3 using suitable indicator
c) Dichromate, Aresenate, H2O2 and Permenganate
AsO43- + 2I- + 2H+
AsO33- + I2 + H2O (Conc HCl is added)
2+
Cr2O7 + 6I + 14H
2Cr3+ + 3I2 + 7H2O
+
2MnO4 + 10I + 16H
2Mn2+ + 5I2 + 8H2O
Liberated I2 is titrated with std Na2S2O3 using suitable indicator
٥١
Potassium Iodate:
Potassium iodate KIO3 is a strong oxidizing agent, the course of its
reaction with reducing agents being influenced by the conditions,
which are used for the reaction:
IO3- + 6 H+ + 6 e
3 H2O + I-
(1)
IO3 + 6 H + 5 e
3 H2O + 1/2 I2
(2)
3 H2O + I+
(3)
-
+
IO3- + 6 H+ + 4 e
In a weak acid medium (0.1 – 1 N HCl) the reaction of iodate with
potassium iodide stops when the iodate has been reduced to iodine
(equation2)
IO3- + 5 I- + 6 H+
3 I2 + 3H2O
In more concentrated acid solutions (4N – 6N HCl) the iodate is
reduced to I+ (reaction 3). Iodine monochloride is formed, ICl forms a
stable complex ion with chloride ion:
ICl + Cl-
(ICl2)
IO3- + 2I2 + 6 H+ + 10 Cl2 IO3 + 2I + 6H + 6 Cl
-
٥٢
٢٦
-
+
-
-
5 [ICl2 ]- + 3H2O
2 [ICl2]- +3H2O
The last reaction is the basis of a method for the determination of a
wide number of reducing substances and was first proposed by
Andrew's. In the above reactions I+ is not stable except in the presence
of high concentration of Cl- or CN- where it forms the fairly stable
[ICl2]- or ICN. Chloride ions are provided by the use of conc. HCl which
provides H+ as well. Lang employed CN- in a solution of lower acidity
(0.5 N-1N HCl), iodine cyanide is formed and this compound provides
the necessary stability.
iodine, iodide, arsenite and antimonite can be titrated with KIO3 (at 4 – 6 N
HCl or 0.5 – IN HCl in presence of CN-)
IO3- + 4 Cu+ + 6H+ + 2 Cl(also Hg+)
IO3- + 2 Sn2+ + 6H+ + 2 ClIO3-+ 2 As3+ + 6H+ + 2 Cl(also Sb3+)
ICl2-+4Cu2+ + 3H2O
ICl2-+2Sn4+ + 3H2O
ICl2-+2As5+ + 3H2O
٥٣
Use of strong oxidizing agent other than iodate in presence of
KI:
2 Ce4+ + I- + 2 ClICl2- + 2 Ce3+
+
MO2 + I + 4 H + 2 Cl
ICl2- + M2+ + 2H2O
(MO2 e.g. MnO2, PbO2, Pb3O4)
Back titration of the unreacted I- with std KIO3
Potassium Bromate KBrO3
It is a strong oxidant in presence of acids being reduced to Br:
BrO3- + 6H+ + 6e
Br- + 3 H2O
(Eo=+1.44v)
It is primary standard, its solutions are stable.
Standard BrO3-/Br- mixture is used as a stable source of
bromine. In neutral solution no bromine is liberated. However upon
acidification the mixture yields the equivalent amount of Br 2:
3 Br2 + 3H2O
BrO3-+ 5Br- + 6H+
Examples for the use of BrO3-/Br- mixture are the determination of phenols,
amines.
٥٤
٢٧
OH
OH
Br
+
Br
+
3 B r2
3 HBr
Br
Phenol
Br2 + 2 II2 + 2 S2O32-
2 , 4 , 6 ,-s y m -trib ro m o p h e n o l (S )
Y e llo w p p t
I2 + 2 BrS4O62- + 2 I-
The reaction is allowed to proceed in the dark for 15 minutes, the
excess Br2 can be converted to iodine by adding KI followed by
titration with sodium thiosulphate using starch as indicator.
Chloroform must be included to dissolve the precipitated
tribromophenol.
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٢٨