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ADVANCED
INORGANIC ANALYSIS

DR. S.K. AGARWALA
M.Sc., Ph.D.
Ex. Convener, R.D.C. and Syllabus Committee,
Ch. Charan Singh University,
Ex. Head, Chemistry Deptt.
Retd. Principal, Meerut College, MEERUT (U.P.).

DR. KEEMTI LAL
M.Sc., Ph.D.
Retd. Senior Reader
Chemistry Deptt.
D.N. College,
MEERUT (U.P.).

PRAGATIPRAKASHAN


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PRAGATIPRAKASHAN
Educational Publishers

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PRAGATI BHAWAN,
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Revised Edition : 2008

ISBN No. : 978-81-8398-523-9

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(Phone : 2641747) and Printed at : Arihant Printers, Meerut.


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CONTENTS
Qualitative Inorganic Analysis

1-162

Chapter

Page

1. Physical Principles (involved in the Analytical chemistry)


3-34

2. Analysis of Acidic Radicals

35-48

3. Tests for some Combination of Acidic Radicals

49-53

4. Removal ofInterfering Radicals

54-57

5. Reactions Involved in the Tests of Acidic Radicals

58-66

6. Tests of Basic Radicals

67-84

.,. Chemical Reactions Involved in the Tests of Basic Radicals

85-92

8. Analysis ofInsoluble Residues

93-95


9. Some Clues Regarding Mixture Analysis

96-97

10. Semi-Micro Analysis of the Mixture

98-104

11. Semi-Micro Method (Systematic Procedure)

105-109

12. Analysis of Mixture of Rare Metal Salts

110-119

13. SpotTestAnalysis

120-136

14. Chromatography

137-142

15. Preparation of Inorganic Compounds

143-162

'0U,"


Quantitative Inorganic Analysis

I

163-355

16. Volumetric Analysis

165-185

17. Acidimetry and Alkalimetry (Neutralization Titrations)
Na2C03 vs HCl
186

186-197

H 2S0 4 t's (NaOH + Na2C03)
H 2S0 4 vs (Na2C03 + NaHC0 3 )
192
Estimation ofNH 3
Problems
193

189
191

18. Oxidation-Reduction (Redox Titrations)
H 2 C 2 0 4 vs KMn04
198


198-207


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(vi)

FeS04(NH4hS04.6H20VS KMn04
199
FeS04' (NH4)S04.6H20 vs K2Cr207 (Internal indicator)
Fe 2+ and Fe 3+ vs KMn04
205
Problems
206

204

19. Iodimetry and Iodometry Titrations
Iodimetry :
12 vs Na2S03.5H20 208
As 20 3 VS 12
209
Iodometry :
K2Cr207 vs Na2S203.5H20 210
211
CuS04 vs Na2S203
212
Cu in brass
Available Cl 2 in Bleaching powder 213

Problems
214

208-215

20. Precipitation Titrations
AgN0 3 vs NH 4CNS 216
NaCl vs AgN0 3 217
AgN0 3 vs NaCl 218
Cl- in water sample 219
Problems
220

216-221

21. Complexometric Titrations
Mg 2+ vs EDTA 225
Ca2+ vs EDTA 226
Ca 2+ vs EDTA (Back titration method)
Problems
230

222-230

227

22. Some Inorganic Reagents used in Volumetric Analysis
Alkaline KMn04
231
NH 4V0 3 232

KI0 3 233
KBr03
236
Ce(S04h
239
TiCl 3 242
CrS04
244

231-246

23. Conductometric Titrations

247-260

24. Potentiometric Titrations

261-270

25. Flame Photometry

271-275

26. pHmetry

276-283


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(vii)

27. Colorimetry

284-291

28. Gravimetric Analysis

292-304

29. Some Gravimetric Estimations
(One Constituent)
Ag + ~ AgCl 305
Cl- ~ AgCl 306
Pb 2+ ~ PbS0 4 307
Pb 2+ ~ PbCr04
308

305-320

Cu+ ~ CuCNS
Fe 2+ ~ Fe203
Al 3+ ~ Al 20 3

309
310
312

Ni2+ ~ (DMG)zNi 313
Zn 2+ ~ZnNH4P04

314
2
+
Ba
~ BaS04
316
SO~- ~ BaS04
317
Ca 2+ ~ CaO 317

Mg2+ ~ MgNH 4P0 4 .6H 20
,,+
Mg'" ~ Mg 2P20 7
319

319

30. Estimation of Two Constituents (When present together)
Copper and Nickel 321
Copper and Zinc 322
Silver and Copper 323
Silver and Nickel 323
Silver and Zinc 324
Silver and Magnesium 324
Copper and Magnesium 325
Copper and Barium 325
Iron and Nickel 326
Iron and Magnesium 327

321-327


31. Estimation of Three Constituents (When present together) 328-333
Copper, Nickel and Zinc 328
Copper, Nickel and Magnesium 329
Silver, Copper and Nickel 329
Copper, Silver and Zinc 330
Silver, Nickel and Zinc 331
Silver-Nickel and Magnesium 332
Iron, Nickel and Zinc 332


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(viii)

32. Analysis of Alloys
Silver coin 334
Nickel coin 335
Solder and Type metal 336
Dolomite 337
Pyrolusite 339
Bleaching powder 341
Cl, Br and I in a mixture 342
Galena 344

334-345

33. Analytical Problems

346-355


Appendix :
Atomic weights 356
Molecular and equivalent weights 357
Indicators 359
Concentrated acids 359
Dilute acids 359
Bases 359
Solutions of other reagents 360
Testing paper 361
Solid reagents 361
Solvents 361
Gas reagents 361
Solubility products 362
Approximate pH values 362
Molecular and equivalent weights of more compounds
Solubility chart 363
Buffer solutions 364
Log and Antilog Tables

362

(i)-(iv)


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ADVANCED INORGANIC ANALYSIS

PART-I

QUALITATIVE INORGANIC ANALYSIS
(Includes: Macro Analysis of tile Mixture, Semi-micro Analysis
of the Mixture, Analysis of the Rare Metals Mixture, Spot Test
Analysis, Chromatography and Inorganic Preparations.)

~

Analysis is the back bone of Chemistry.
Theory Guides, Practical Decides.
~ Any Scientific Advance is Advance in Methods of Analysis.
~


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"This page is Intentionally Left Blank"


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PHYSICAL PRINCIPLES
(Involved in the Analytical Chemistry)

••• ••••• •

III INTRODUCTION
The term Analysis is commonly employed for the process of breaking up or
separation of a compound or a mixture into its constituents. The branch of chemical
analysis which aims to find out the constituents of a mixture or compound is known
as Qualitative Analysis. The other branch of analysis aiming to determine by weight

or by volume the exact quantities of the different constituents present in the
substance is known as Quantitative Analysis. For a complete analysis the qualitative
analysis must preceed the qwmtitative analysi.J. The former serves as a guide to the
methods to be followed in the latter analysis.
Analytical Chemistry

Qualitative Analysis

*

Quantitative Analysis

+

*

+

+

Volumetric Analysis
Gravimetric Analysis
The great importance of analytical chemistry is for the reason that it enables us
to identify and study the composition of an unknown substance. The identification
of a substance usually involves its conversion into a new substance possessing
characteristic properties with the help of one or more substances of known
composition. This change is called a chemical change or reaction. The substance
which is used to bring about such change is called a Reagent. The qualitative
analysis is carried out by Wet methods; when the identification is done from aqueous
solution and Dry method; when the identification is done in the dry state. The wet

methods have wider range of applicability and are more important than dry
methods.
Qualitative analysis may be carried out on different scales like Macro, Semi
micro or Micro. In Marco analysis 0.1 to 1.0 g of the substance is used and volume of
the solution in generally between 10 to 20 mL. In semi micro analysis 0.01 to 0.1 g of
the solid substance or 1 mL of the solution is employed. In micro analysis less than
0.01 g of the substance is used. There is no striking difference in micro and
(3)


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QUALITATIVE INORGANIC ANALYSIS

semimicro analysis. It lies only in the relative quantity of the substance employed for
analysis. For micro analysis special type of apparatus and techniques have been
developed.
Weight of Solid Analysed
Class
Macro
Meso or Semi-Micro
Micro
Sub-Micro or Ultramicro

1-0.1 g
0.1-0.01 g
0.01-0.001 g
0.001-0.0001 g


In macro analysis large samples are taken which often result in tedious and
time consuming operations. Micro, Semimicro and ultramicro analysis, on the other
hand are much faster and special techniques must be learnt before they can be
successfully employed.
Recently spot analysis has been devised to afford the greatest possible
economy of material, time, space and labour.
A 'spot test' is a test that can be made directly on the unknown, using only a
drop or two of the material.
Spot tests have two distinct disadvantages:
(a) Many spot tests do not give clear conclusive results.
(b) Many spot tests employ the use of costly organic reagents, that involve
complex organic reactions.
F. Feigl and his coworkers have considerably elaborated the spot test analysis
technique. These tests are carried out on a spot plate, using only a few drops of the
unknown solution. The reagents produce ex~emely pronounced colours and
exceedingly small quantities of a substance (10- 0 g) can be readily detected.
An important characteristic of spot methods is the simplicity of the technique and
apparatus by which very small amount of substance may be detected. Spot test analysis
is the scheme for quantitative and qualitative analysis of the compounds (Organic
and Inorganic) in which the sensitive and selective tests based on chemical reactions
are used with drop of the test or reagent solution present in micro and semi-micro
quantities.
The most essential parts of spot test analysis are :
CO Spotting of reactants.
(ii) Sensitive and Selective reagents.
(iii) Laboratory and equipment requirements.

III QUALITATIVE ANALYSIS
In this branch of analytical chemistry it is to learn the methods used for the

identification of an unknown compound or of a salt and to find out the elements or
radicals present in the given mixture of two or more compounds or in the given
mixture of two or more salts. A salt is formed by the interaction of an acid with a
base:
HCI + NaOH ~ NaCI + H 2 0
Acid

Base

Salt

Water

Every salt consists of two parts, usually radicals known as the Basic
Radical and the Acidic Radical. Sodium chloride is formed by the interaction of
hydrochloric acid and sodium hydroxide, as is evident from the above equation. In


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5

PHYSICAL PRINCIPLES

sodium chloride (NaC!), Na+ radical has come from sodium hydroxide (a base),
hence it is known as basic radical. The basic radical is also called the Positive
Radical since it always bears a positive charge. The CI- radical in sodium chloride
has come from hydrochloric acid (an acid), hence it is known as acidic-radical. The
acidic-radical is also called the Negative radical since it always bears a negative
charge.

When a salt is dissolved in water it undergoes ionization giving positive and
negative ions e. g.,
+
CINaCI ~ Na+
Sodium
chloride salt

--+

Sodium ion
Positive ion
Basic radical
Cation

Chloride ion
Negative ion
Acidic radical
Anion

High dielectric
constant

Electrostatic
Lines of Force Binding Ions
The ion that carries a positive charge is called the positive ion or the cation
(Na+), and the one that carries a negative charge the negative ion or the anion
(CI- ).
The detection of cations (Basic Radical) and anion (Acidic Radical) in a salt or in
a mixture is known as Inorganic Qualitative Analysis.


111 QUALITATIVE
PHYSICO-CHEMICAL PRINCIPLES INVOLVED IN
ANALYSIS
(a) Law of Mass Action: The law states that the velocity of a chemical
reaction is proportional to the product of the "active masses" of the reacting
substances. By the term 'active mass' is usually meant the concentration in mole L-1
Consider a simple reversible reaction,
A+B~C+D

Suppose the velocity with which A and B react together is vI. Let the active
masses of A and B be represented by [A] and [B], then
VI

= kI

[A] [B]

where kl is the constant of proportionality known as the velocity coefficient.
Similarly the velocity v 2 with which the back reaction occurs is given by
v2 = k2 [C] [D]

At equilibrium the velocity of the forwa.rd and backward reactions are the
same so that
kl [A] [B] = k2 [C] [D]


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6


QUALITATIVE INORGANIC ANALYSIS
[C][D]
[A] [B]

or

=!L =k
k2

k is termed as equilibrium constant of the reaction at a particular temperature. An
example of the application of the law may be given. The following equilibrium exists
in a dilute solution of hydrochloric acid at a given temperature :
HCI ~ H+ +CIApplying the law of mass action,
[H+][CI-]

----=K
[HCll

where K is the ionization constant or dissociation constant of hydrochloric acid at a
particular temperature. The ionization constant of an acid or a base is a measure of
the strength of the acid or the base respectively.

(b) Common Ion Effect: The degree of ionization of weak electrolyte is
suppressed by the addition of strong electrolyte containing a common ion.
For example, ammonium hydroxide ionizes in solution, thus:
NH 40H ~ NHt + OHOn applying the law of mass action,
[NHt] [OH-] = K
[NH 40H]
On the addition of ammonium chloride, ammonium ions are added to the
solution. The concentration of NH:4 increases, and since K is constant at any fIxed

temperature, there must be an increase in the concentration of NH 4 0H and a
decrease in the concentration ofOH-.
NH 4 CI ~ NHt
+ CICommon ion

Thus the ionization of NH 4 0H is diminished by the addition of NH 4CI which
furnishes the common ion, NHt .

The principle of common ion effect has a great importance in qualitative analysis.
Ionization of hydrogen sulphide is suppressed in the presence of hydrochloric
acid:
H 2S ~ 2H+ + S2HCl ~ H+

+CI-

Common ion

Ionization of ammonium hydroxide is suppressed in the presence of
ammonium chloride :
NH 40H ~ NHt + OHNH 4 CI ~ NHt

+ CI-

Common ion

Ionization of acetic acid is suppressed in the presence of sodium acetate :
CH 3 COOH ~ CH 3 COO- + H+
CH 3COONa ~ CH 3COO- + Na +
Common ion



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7

PHYSICAL PRINQPLES

*(c) Solubility Product and Precipitation: When a sparingly soluble
substance, say AB is kept in water for sometime at a definite temperature, the
following equilibrium is established :
AB
~
AB ~ A+ + BSolid (More)

Dissolved
01ery less)

~

ions

Applying the law of mass action,
[A + ][B-]
-'-----=K
[AB]Diss.

But the concentration of the unionized [AB] is constant at a given temperature
if excess of AB is present.
[A + ] [B-] =K [AB] Diss. = A constant, K sp.
Hence


Hence in equilibrium, the product of the ionic concentrations is
constant at a given temperature. This constant product [A+][B-] is
called the solubility product.
When the ionic product exceeds the solubility product, the solution is super
saturated and precipitation occurs, and if the ionic product is less than the solubility
product, the solution will be unsaturated and the precipitation will not occur.
This is also called as the theory of precipitation :
Ionic product < Solubility product, the salt dissolves.
Ionic product> Solubility product, precipitation takes place.
Since precipitation is governed by the principle of solubility product, the latter
receives important applications in the field of analytical chemistry.
Precipitation of the sulphides of Group II and IV: Precipitation of
the sulphide can occur only when the ionic product [M2+] [S 2-] exceeds the
solubility product of the sulphide (MS) at that temperature
[M2+][S 2-] > K sp (Solubility product).
Precipitation occurs

In the presence of acid (Hel) the ionization of H 2S
H 2 S ~ 2H+ + S2is suppressed due to the increase ofH+ ions (produced by acid) so that there are few
S 2- ions in solution and the solubility product of the sulphides of Group IV radicals
is not reached. It is however, enough to cause the precipitation of the lInd group
radicals e.g., CuS, CdS etc. which possess a low solubility product.
In the presence of NH 40H, the hydroxide ions obtained from it [NH40H~
NH4 + OH-] unite with the H+ ions produced from H 2S (H 2S ~ 2H+ + S 2- ) to
give unionized water, so that more of H 2S ionizes and thus the concentration ofS 2ions in solution increases. In this way it becomes high that the solubility product of
IV Group radicals e. g., ZnS, MnS etc. is exceeded and these precipitate out.
[M2+] [S 2- ] > K sp (Solubility product)
Precipitation occurs


* For a sparingly soluble binary electrolyte, the product of the total molar concentrations of the ions is
constant at constant temperature and this product is termed as the solubility product (K sp)'


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QUAUTATIVE INORGANIC ANALYSIS

Table 1.1: Solubility products of some chlorides
Substance

Solubility product

PbCl2

2.4 x 10-4
2.0 x 10-21
1.2 x 10-10

Hg~12

A£CI

Table 1.2: Solubility products of some sulphides
Substance

Solubility product


Substance

HgS

4 x 10-53
4.2 x 10-28

FeS

PbS

1.6 x 10-72
8.5 x 10-45
3.6 x 10-29

Bi~3

CuS
CdS

Solubility product

ZnS
MnS
NiS
CoS

1.5
1.0
1.4

1.4
1.0

x 10-19
x 10-20
x 10-15
x 10-24
x 10-27

Table 1.3: Solubility products of some
common hydroxides at 18°C
Substance

Solubility product

Fe (OH)3

1.1 x 10-36

Al(OH)3

8.5
1.8
4.0
3.4

Zn(OHh
Mn(OHh
Mg(OHh


x 10-23
x 10-14
x 10-14
x 10- 11

(d) Complex Ion Formation: Complex ion formation is of immense
importance in qualitative analysis
(a) to dissolve a precipitate alone or from a mixture of two,
(b) to check the precipitation of particular cation by complex ion formation.
For example:
(i) Silver chloride is soluble in ammonia solution due to the formation of the
complex ion, [Ag(NH3 h]+
AgCI + 2NH3 ~[Ag(NH3h]CI
or
Ag+ + 2NH3 ~ [Ag(NH 3 )2]+

The capacity of silver ions to form complex ion with ammonia is utilized in
separation of Ag+ from Pb++ or Hg~+
(ii) Separation of copper and cadmium is also based upon complex formation
with KCN solution. The complexes formed are K3[Cu(CN)4] and K2[Cd(CN)4]'
The complex salts ionize as follows :
K3[Cu(CN)4] ~ 3K+ + [Cu(CN)4]3[Cu(CN)4]3- ~ Cu+ + 4CN- (very low)
K2[Cd(CN)4] ~ 2K+ + [Cd(CN)4] 2[Cd(CN)4]2- ~ Cd 2+ + 4CN- (High)


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9

PHYSICAL PRINCIPLES


Applying the law of mass action
[Cu+][CN-]4 = K = 5 x 10-28
[{Cu(CN)4} 3-]
2
[Cd +][CN-]4 = K = 1.4 x 10-17
[{Cd(CN)4} 2-]

K is the instability constant of the complex ion and determines its stability. The
value of instability constant for copper complex ion is lesser than cadmium complex
ion and hence more Cd 2+ ions are available in the solution than Cu 2+ ions. When
hydrogen sulphide gas is passed through a solutio~ containing copper and cadmium
cyanide complex ion, copper complex remains unaffected while cadmium complex
ion gives cadmium ions which combine with S 2- ions (available from H 2S) to form
a yellow precipitate of CdS.
This leads to the precipitation of cadmium as cadmium sulphide in presence of
copper.
[Cd(CN)4]2- ~Cd2+ + 4CNH2S~ 2H+

+ S2-

Cd 2+ + S2- ~ CdS,!..
Yellow ppt.

(e) Oxidizing and Reducing Agents: Oxidation is a process which
results in the loss of one or more electrons by atoms or ions. An oxidizing agent is
one that gains electrons and is reduced to a lower oxidation state e. g.,
KMn04' K2Cr207, HN0 3 , H 20 2 etc.
2KMn04 + 3H 2S0 4 ~ K2S0 4 + 2MnS04 + 3H 20 + 50
Mn7+ (Pink)

Mn 2+ (Colourless)
Reduction is a process which results in the gain of one or more electrons by
atoms or ions. A reducing agent is one that loses electrons and becomes oxidized to a
higher oxidation state e. g., SO 2, H 2S, SnC1 2, HI etc.
SO§- + H202~SO~- + H 20
(f) Concentrations of Reagents :
Molar and Normal Solutions: Since in qualitative analysis one deals
continuously with solutions of electrolytes, so concentrations of both the unknown
substance and the reagents should be expressed exactly and non-ambiguously.
There are two popular methods of expressing the concentrations of solutions.
(i) Molar Solutions: A molar solution (lM) is one which contains one
gram molecular weight (Formula weight), i. e., one mole of solute per litre of
solution.
For example, a 1 M solution result if 58.45 g of NaCI, or 98.08 g of
H 2S0 4 or 126.1 g of H 2C 20 4 . 2H 20 is dissolved in a litre of solution.
When one is concerned with the reactions of the solute then molarity system of
defining the concentration of a solution is adopted.
(li) Normal Solutions: A normal solution (1 N) is one which contains
one gram equivalent weight of solute per litre of solution. The different


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10

QUAUTATIVE INORGANIC ANALYSIS

concentrations of solutions, are expressed as multiples or fractions of the normal
solutions, e. g. , five times normal is represented as 5 N and one tenth of normal as
0.1 N (N/I0).

The gram equivalent weight is defined as that weight of substance which will
react with or displace 1.008 g of hydrogen, 8 g of oxygen, 35.5 g of chlorine or that
quantity of any element which reacts with these weights of hydrogen, oxygen and
chlorine respectively. Usually the equivalent weight of an element is equal to its
atomic weight divided by its valency.
The equivalent weight of a compound is that weight which contains one gram
equivalent weight of the component taking part in the reaction under consideration.
Thus with acids it is ionizable and replaceable hydrogen, with bases it is generally
the cation or the anion. In general, the gram equivalent weight is dependent on the
particular reaction in which the substance takes part.
It often happens that the same compound possesses different equivalent
weights in different chemical reactions.
The gram equivalent weight of an acid is that weight of it which contains one
replaceable hydrogen atom, i. e., 1.008 g of hydrogen.
The equivalent weight of monobasic acids (e.g., HCI, HBr, HI, HN0 3 or
CH 3COOH) is identical with their molecular weights.
A normal solution of monobasic acids will, therefore, contain 1 gram
molecular weight (1 mole) of substance in a litre of solution.
The equivalent weight of a dibasic acid (e. g., H 2S0 4, H 2C 20 4) or of a tribasic
acid (e. g., H3P04) is likewise.!. and .!., respectively of the molecular weight.
2
3
The gram equivalent weight of a base is that weight of it which contains one
replaceable hydroxyl group, i. e., 17.008 g of ionizable hydroxyl group. 17.008 g of
hydroxyl group is equivalent to 1.008 g of hydrogen.
The equivalent weight of NaOH, KOH and NH 40H are gram molecular weight
(1 mole) and of Ca(OHh, Sr(OH)2 and Ba(OHh is half gram molecular weight
(.!. mole).
2


The gram equivalent weight of a normal salt is that weight of it which contains
one gram equivalent weight of the cation or anion. This quantity will be molecular
weight of the salt divided by the total valency of cation or anion. Thus, the
equivalent weight ofKCI is 1 mole, ofNa2S04 is.!. mole; of AlCl 3 is.!. mole; ofSnCl 4
.

2

3

is.!. mole and ofCa3(P04h is 1/6 mole.
4
The gram equivalent weight of oxidizing and reducing agents :
(i) The equivalent weight of an element taking part in an oxidation-reduction
(redox) reaction is the atomic weight divided by change in oxidation number.
(ii) When an atom in any complex molecule suffers a change in oxidation
number, the equivalent weight of the molecule is the molecular weight divided by
the change in oxidation number of the oxidized or reduced elements. If more than
one atom of the reactive element is present, the molecular weight is divided by the
total change in oxidation number.


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11

PHYSICAL PRINOPLES

For example, the equivalent weight of KMn04 is different in different media.
In Acidic medium

7+
2+
2KMn04 + 3H 2S0 4 ~ K2S0 4 + 2MnS04 + 3H 20 + 5[0]
Mole. wt. of KMn04
Eq. wt. 0 f KM n 4 - - - - - - - - - - - - ' Change in Ox. No. of Mn

°

=

In Neutral medium
7+
2KMn04 + H 20

158 = 158 = 31.6
7-2
5

4+
2KOH + 2Mn02 + 3[0].
Mole. wt. of KMn04
Eq. wt. of KMn04 =
Change in Ox. No. of Mn
~

=

158

7-4


= 158 = 52.6
3

In Alkaline medium
7+
6+
2KMn04 + 2KOH ~ 2K2Mn04 + H 20 +[0]
Mole. wt. of KMn04
Eq. wt. 0 f KMn 4 = - - - - - - - Change in Ox. No. of Mn

°

=

158

= 158

7-6
(g) Scale of Acidity or pH (Measure of Acidity or Alkalinity): pH
of a solution plays an important role in the precipitation of metal cation. pH of a
solution is a measure of H+ ion concentration and is defined as "the numerical value
of the negative power to which 10 must be raised in order to express the hydrogen ion
concentration. "
[H+] = 10-°10-110-210-310-410-510-610-710-810-910-1°10-1110-1210-1310-14
pH = ,0

1


2

3

4

5

6

v

7
'

Acidic

8

Neutral '

9

10

11

12

13


14

v

Alkaline

Mathematically it is expressed as
[H+] = lO-pH
pH = -log [H+] = log _1_
[H+]
Therefore pH may be defined as, "the logarithm of the reciprocal of the H + ion
concentration.

or

Table 1.4: pH of some common solutions
Solution

pH

Character

0.1 N HCI
Double seven (Cold drink)
Grape fruit juice
Soda water

1.0
3.1

3.1
4.0

Acidic
Acidic
Acidic
Acidic


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QUAUTATIVE INORGANIC ANALYSIS

Solution

pH

Character

Blood
0.1 N-NaOH
Sodium carbonate
Pure water

7.2
13.0
8.0
7.0


Alkaline
Alkaline
Alkaline
Neutral

*Buffer solution: A Buffer solution possesses the following characteristics:
(0 pH of such a solution does change either on dilution or on keeping for a long
time.
(iO pH of such a solution is not altered by a small addition of either an acid or
alkali.
The important buffer solutions are :

CH 3 COOH + CH 3 COONa; NH 40H + NH 4CI etc.
(h) Complex Formation: Before explaining complex formation, it is
most essential to understand simple, double and complex salts.
Simple salt: It is generally formed by the neutralization process, i. e., by
the reaction of acid and alkali. These simple salts ionize when dissolved in water.
NaOH + HCI
Alkali

~

Acid

NaCl

NaCI
Simple Salt


~ Na+

Simple salt

+ H 20

Water

+ Cl-

~

IOns

Double salt: It has been observed that when two or more normal salts are
mixed in requisite proportions and allowed to crystallize together, a double salt is
formed.
FeS04 + (NH4)2S04 + 6H 20

Simple salt

Simple salt

Water

~

FeS04 . (NH4hS04 ·6H 20
Double salt


Double salt gives in aqueous solution, the test of all its constituents ions, i. e.,
Fe 2+, SO~- and NHt. Such substances are called Double salts or Lattice
Compounds.
Complex salt: It has been observed tbat when FeS04 solution and KCN
solution are mixed together and evaporated, K4[Fe(CN)6] is formed which in
aqueous solution does not give test for the component Fe 2 + and CN- ions but gives
test for K+ and the complex ferrocyanide ion, [Fe(CN)6]4K4 [Fe(CN)6] ~ 4K+ + [Fe(CN)6]4The dissociation of ferrocyanide ion, [Fe(CN)6]4- is so negligible, as not to
confIrm to the usual tests of the ferrous ion.
The formation of new complex ion gives new chemical and physical character to
the substance.

A complex ion is defined as a charged radical which is formed by
the combination of a simple cation with one or more neutral
molecules or one or more other simple ions.
* Buffer solution possess reserve acidity or reserve alkalinity.


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13

PHYSICAL PRINCIPLES

e. g., when hydrogen ion reacts with the normal anions of polybasic acids, we have
H+ + CO~- ~ HCO

s

H+ + SO~- ~ HS0 4
H+ + PO~- ~ HPO~When sodium hydroxide solution is added in small quantities to a solution

containing Zn 2+ , a white precipitate of zinc hydroxide is formed. This Zn(OHh
dissolves in excess of NaOH solution to give a complex salt, sodium zincate
Na2 (Zn02)' 2H 20 or Na2[Zn(OH)4]'
Zn(OH)2 + 20H-~[Zn(OH)4]2The formation of this complex ion [Zn(OH)4]2- , reduces the Zn 2+
concentration with the result that the ionic product of[Zn 2+] and [OH-]2 is less
than the solubility product and so zinc hydroxide dissolves. Similar behaviour of
NaOH solution is observed towards aluminium ions in solution to form Na[Al(OH)4]
or Na[Al0 2]· 2H 20.
Al(OHh + OH- ~ [Al(OH)4r
Chromic hydroxide behaves similarly and the complex ion may be written as
[Cr(OH)4r. Stannous salts give initially a precipitate of stannous hydroxide which
dissolves in excess of alkali to form sodium stannite, Na2[Sn02] aq. or
Na[SnO ·ONa] aq.
Sn(OH)2 + OH- ~ [Sn(OH)3r
Stannic salts yield Na2[Sn03]' 3H 20 or Na2[Sn(OH)6] complexes e.g.,
Sn(OH)4 + 20H- ~ [Sn(OH) 6] 2Plumbite [Pb(OH)3r and plumbate [Pb(OH)6]2- ions are also evident in
solutions.
Complex ions are often formed by the interaction of inorganic ions with
organic ions or molecules. The combination of ferric and oxalate ions; cupric and
tartarate ions may be cited as examples.
Fe 3+ + 3C20~- ~ [Fe(C 20 4 h]3Cu 2+ + 2C4H40~- ~ [Cu(C 4H40 6 h]2The ferric oxalate complex is simply stable. The Fe 3+ ion can be reduced (by
the addition of sufficient C 20 ~- ions) to so low value that the solution does not give
the red colour of the ferrithiocyanate Fe(SCN) 3 with a soluble thiocyanate.
Certain non-volatile organic compounds containing hydroxyl group such as
tartarates citrates etc. form complex ions with various metals e. g.,
j
Cu 2+, Fe +, Al3+, Cr 3+ etc. These complex ions are very stable and yield such
concentrations of the simple metal ions that they often fail to respond that the
non-volatile organic matter is destroyed before proceeding with the systematic
detection of the metals.

Complex ions are formed by Cu 2+, Cd 2+, Ni2+, C0 2+ and Zn 2+ with NH3
molecule.


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14

QUAUTATIVE INORGANIC ANALYSIS

Many halide complex ions are known which may include
[HgCI 4 ]2- , [SnC1 6 ]2-, [SbCI 6 ]3-, [FeF6 ]3- ,[AlF6 ]3-, [AuCI 4 ]3- and [PtCI 4 ] 2Complex sulphide ions are that of As, Sb and Sn.
As 2S 3 + 3S 2- ~ 2AsS~As 2S S + 3S 2- ~ 2AsS~­
Sb 2S 3 + 3S 2- ~ 2SbS~­
Sb 2S 5 + 3S 2- ~ 2SbS~SnS2 + S2- ~ SnS~A complex formation is accompanied by the following changes:
A sudden change in the solubility: It has been observed that the
solubility of AgCN appreciably increases by the addition of KCN, due to complex
formation, K[Ag(CNh].
AgCN + KCN ~
K[Ag(CN)2]
Pot. argentocyanide (complex)
OR Pot. dicyanoargentate(I)

Drop in conductivity: Complex ion formation results by the combination
of two or more ions and hence total number of ions in solution decreases and
consequently complex ion becomes heavier and there is a marked drop in the
conductivity of the solution.
Colour change: Sometimes a colour change is accompanied by complex
formation.
G.T. Morgan and Drew proposed the term chelate for those cyclic

structures which arises from the union of metallic atoms with organic or inorganic
molecules. The rings of such compounds are called chelate rings and the
phenomenon as the chelation.
The chelation ring system may be formed by groups which have more than one
point of attachment to the metal, i. e., bidentate or tridentate groups etc.
,H2\/iH2
CH2

I

M

t

CH2

I

CH2-NH-CH2
Bidentate
chelating agent

Tridentate chelating
agent

Normally chelated complexes are more stable than the simple non-chelated
complexes. The more rings that are formed, the more stable are the complexes.
H2-NH2~

I


[C

CH2-NH2~

/"NH 2-CH 2
Cu

2+

I

~NH2-CH2 }

Its enhanced stability has been attributed due to chelate ring.


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PHYSICAL PRINCIPLES

2+
NH3

1
r

H3N~Cu


+--NH3

NH3

(Complex)
(Non -chelated)

III REACTIONS INVOLVING COMPLEX FORMAnON
Complex formation plays an important part in the analytical chemistry.
Complex formation is of great importance in qualitative and quantitative
(gravimetric and volumetric) analysis.

QUALITATIVE ANALYSIS
The complex forming reactions are as follows :
(i) Reactions with NaOH: Aluminium and zinc salt solutions form
sodium aluminate and sodium zincate with excess of NaOH, respectively.
AlCl 3 + 3NaOH ~ Al(OH) 3 .J.. + 3NaCI
Al(OHh + NaOH ~ NaAl0 2 + 2H 20
Sod. aluminate

ZnC1 2 + 2NaOH ~ Zn(OH)2.J.. + 2NaCI
Zn(OHh + 2NaOH ~ Na2Zn02 + 2H 20
Sod. zincate

Similarly chromium hydroxide dissolves in excess of NaOH forming a
chromite, chromium salt in the presence of oxidizing agents
CrCl 3 + 3NaOH ~ Cr(OHh .J.. + 3NaCI
2Cr(OHh + 2NaOH


~

Na2Cr204 + 4H 20

Sod. chromite

such as Br2 water, H 20 2 forms a soluble yellow chromate ion.
Br2 + 2NaOH ~ 2NaBr + H 20 + [0]
2CrCl 3 + lONaOH + 3[0] ~
2Na2Cr04
+ 6NaCI + 5H 20
Sod. chromate (yellow)

These reactions are employed in the separation of aluminium from iron and
chromium or zinc from manganese.
(ii) Reactions with Potassium Iodide: Solution of iodine in potassium
iodide: Iodine dissolves readily in KI solution forming a complex salt, 'potassium
tri-iodide.' KI3 dissociates into K+ and 13, But the solution of iodi.ne in potassium
iodide behaves as if the solution is of iodine.
KI+I2 ~KI3~K+ +13

The action of potassium iodide on mercuric salt: Mercuric chloride
reacts with KI forming a scarlet precipitate of mercuric iodide which dissolves in
excess of KI forming K2Hg1 4 .


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16


QUAUTATIVE INORGANIC ANALYSIS

HgCl 2 + 2KI
HgI2 + 2KI

~
~

HgI2 + 2KCI
K2HgI4
Complex

When this complex is dissolved in an excess of NaOH, the solution is known as
"Nessler's reagent" which is employed for the detection of ammonia. It gives brown
colour or precipitate with ammonia.

~2HgI4 + 3NaOH} NH40H ~ Hg~
"V"

Hg

Nessler's reagent

NH2

0

1

I


Iodide of Million's base
Brown ppt.

+ 4KI + 3NaI + 3H20
(iii) Reactions with Potassium Cyanide: Separation of Copper and
Cadmium: When in copper and cadmium salts solution KCN is added in excess, both

copper and cadmium form complexes. These two complexes differ appreciably in
their stability.
CUS04 + 2KCN ~ Cu(CNh,j, + K2S0 4
Yellow ppt.

2Cu(CNh ~ CU2(CNh,j, + (CNh

t

White ppt.

CU2(CNh + 6KCN

~

2K 3[Cu(CN)4]

Pot. tetracyanocuperate(I)

CdS0 4 + 2KCN
Cd(CNh + 2KCN


Cd(CNh + K2S0 4
~
,K 2[Cd(CN)4]
~

Pot. tetracyanocadmiate(II)

The copper complex is much more stable than the cadmium complex with the
result when H 2S is passed in the aqueous solution of these two complexes, cadmium
alone gets precipitated.
K3[Cu(CN)4] ~ 3K+ + [Cu(CN)4]3[Cu(CN)4]3- ~ Cu+ + 4CN- (dissociation is negligible)
K2[Cd(CN)4] ~ 2K+ + [Cd(CN)4]2[Cd(CN)4]2- ~ Cd 2+ + 4CN- (dissociation is very high)

Separation of Nickel and Cobalt: Cobalt salt forms a complex salt,
potassium hexacyanocobaltate(II) , K4[Co(CN)6] with excess of KCN which in
slightly acid solutions in air is oxidized to another stable complex, potassium
hexacy;mocobaltate(III) , K3 [Co(CN)6], while nickel only forms nickel cyanide,
Ni(CNh which on boiling with NaOH and bromine water is converted into black
oxide of nickel.
CoCl 2 + 2KCN ~ Co (CN)2 + 2KCI
Co(CNh + 4KCN ~ K4[Co(CN)6]
2K 4 [Co(CN)6] + H 20 + 0 ~ 2K 3[Co(CN)6] + 2KOH
NiCl 2 + 2KCN ~ Ni(CN)2 + 2KCI


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PHYSICAL PRINCIPLES


(iv) Reactions with Potassium Nitrite: Cobalt reacts with potassium
nitrite to form potassium hexanitrocobaltate(II), K4[Co(N0 2 )6] which is oxidized
to potassium cobaltinitrite, a yellow precipitate results.
CoCl 2 + 2KN0 2 ~ Co(N0 2 h + 2KCI
Co(N0 2 )2 + 4KN0 2 ~
K4[Co(N0 2 )6]
2K 4[Co(N0 2 )6] + 2CH 3COOH + O(air)

Pot. hexanitrocobaltate(II)

~

2r~3[Co(N02)6]

Yellow ppt.
Pot. hexanitrocobaltate(III)

,j, + 2CH 3COOK + H 20

Under the conditions nickel only gives a water soluble double salt,
Ni(N0 2 h.4KN0 2·
(v) Reactions with Sodium Bicarbonate: When the solution of cobalt
and nickel salts is treated with excess of sodium bicarbonate cobalt forms a pink
coloured complex sodium cobalto-carbonate, which on treatment with bromine
water is oxidized to green coloured sodium-triscarbonato cobaltate(III). Nickel does
not form a complex with NaHC0 3 but on heating with bromine water, it is oxidized
to black nickelic oxide.
CoCl 2 + 2NaHC0 3 ~ Co(HC0 3 h + 2NaCI
Co(HC0 3 h + 4NaHC0 3 ~ Na4[Co(C03)3] + 3H 20 + 3C0 2 t

Sod. triscarbonato
cobaltate (II)

2Na4[Co(C03h] + Br2

~

2Na3[CO(C03h] + 2NaBr
Sod. triscarbonato
cobaltate (III) (green)

Br2 + H 20 ~ 2HBr + [01
NiCl 2 + 2NaHC0 3 ~ NiC0 3 + 2NaCI + H 20 + CO 2 t
2NiC0 3 + 3H 20 + [0] ~ 2Ni(OHh + 2C0 2 t
2Ni(OHh ~ Ni 20 3 ,j, + 3H 20
Black ppt.

(vi) Reactions with Ammonia: AgCI and AgBr dissolve in dil. and
concentrated ammonia forming respective complexes.
AgCI + 2NH3 ~ [Ag(NH 3 h]Cl
AgBr + 2NH3 ~ [Ag(NH 3 h]Br
Ag1 with liquid ammonia forms Ag(NH 3 )1
Ag1 + NH3 ~ Ag(NH 3 )I
Copper salts give intense blue colour with ammonia.
Cu(N0 3 h + 2NH 40H ~ Cu(OHh,j, + 2NH 4N0 3
Cu(OHh + 4NH3 ~ [Cu(NH 3 )4](OHh ~ [Cu(NH 3 )4]2+ + 20HBlue colour

Tetra-ammine copper (II) ion [Cu(NH 3 )4]2+ is blue.
Cadmium salts also form complex with ammonia similar to copper.
Cd(N0 3 h + 2NH 40H ~ Cd(OHh,j, + 2NH 4 N0 3

Cd(OH)2 + 4NH3 ~ [Cd(NH 3 )4] (OHh
Mercurous chloride with ammonia gives a black coloured mass (metallic
mercury and NH 2HgCl). The whole mass appears to be black on account of the
finely divided mercury.


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18

QUAliTATIVE INORGANIC ANALYSIS

Hg 2Cl2 + 2NH 40H ~ !fgNH2CI + H~ + NH 4CI + 2H 20
Bl~ck
(vii) Reactions with Ammonium Poly-sulphide : In the second
group of Qualitative Analysis the sulphides of the copper group, (i. e., Hg, Pb, Bi, Cu

and Cd) are separated from those of the arsenic group (As, Sb and Sn) by means of
yellow ammonium sulphide. Yellow ammonium sulphide contains excess of sulphur
and is represented by the formula (NH 4 h S x'
Arsenic group sulphides form the following complex and so go in the solution:
Arsenic:
As 2S 3 + 3(NH 4 hS ~ 2(NH 4 h AsS 3
As 2S S + 3(NH4hS ~ 2(NH 4 hAsS 4
Antimony:
Sb 2S 3 + 3(NH4hS ~ 2(NH 4 )3 SbS 3
Sb 2S s + 3(NH4hS ~ 2(NH 4 h SbS 4
Tin:
SnS + (NH4hS2 ~ (NH4)2SnS3
SnS2 + (NH 4 h S ~ (NH 4 h SnS 3

These thioarsenites, thioarsenates, thioantimonites, thioantimonates and
thiostannates are soluble complexes from which original sulphides of arsenic,
antimony and tin can be precipitated simply by making the solution acidic.
(viii) Reactions with Potassium Ferrocyanide: Copper salt solution
in presence of acetic acic,i gives a chocolate red precipitate of copper ferrocyanide.
2Cu(N0 3 h +
K4[Fe(CN)6]
~ CU2[Fe(CN)6] ,j,. + 4KN0 3
Pot. hexacyanoferrate(II)

Copper ferrocyanide
Chocolate red ppt.

Ferric salt solution gives a deep blue colour with potassium ferrocyanide due
to the formation of ferri-ferrocyanide.
4FeCl 3 +
3K 4 [Fe(CN)6]
~ Fe4[Fe(CN)6h + 12KCI
Pot. hexacyanoferrate(II)

Ferri - ferrocyanide
Deep blue colour

(ix) Reactions with Sodium Thiosulphate: Sodium thiosulphate
solution dissolves silver chloride, bromide and iodide.
2AgCI + Na2S203 ~ Ag 2S 20 3 + 2NaCI
SNa2S203 + 3Ag 2S 20 3 ~ 2NaS[Ag3(S203)4]
Complex

When sodium thiosulphate solution is added to copper sulphate solution in

cold, the blue solution of copper sulphate becomes successively brownish green,
yellowish green and finally yellow.
CuS04 + Na2S203 ~ CuS203 + Na2S04
2CuS203 + 2Na2S203 ~ 2Na. CuS 20 3 + Na2S406
Sodium cuprous thiosulphate
(Comp/ex)

If the above solution is boiled, cupric sulphide is formed.
2NaCuS203 ~ 2CuS,j,. + Na2S206
Black ppt.

Na2S206 ~ Na2S04 + S02

t