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d - Block elements
A transition element may be defined as an
element whose atom in the ground state or ion in
common oxidation state has incomplete sub-shell, has
electron 1 to 9. It is called transition element due to
fact that it is lying between most electropositive ( s -
block) and most electronegative ( p - block) elements and
represent a transition from them. The general
electronic configuration of these element is
( 1 ).
1 to 10 0 to 2
n ns
The definition of transition metal excludes Zn , Cd
and Hg because they have complete d - orbital. Their
common oxidation state is , ,.
Zn Cd Hg They also do
not show the characteristics of transition element.
Element of group 3 ( Sc , Y , La and Ac ) and group 12
( Zn , Cd , Hg ) are called non typical transition element.
Table : 19.1 First transition or 3 d series :
Element Symbo
l
At. No. Electronic
configuration
Scandium Sc 21
d
orbitals are filled up
[ Ar ] 3 d
1
4 s
2
Titanium Ti 22 [ Ar ] 3 d
2
4 s
2
Vanadium V 23 [ Ar ] 3 d
3
4 s
2
Chromiu
m
Cr
*
24 [ Ar ] 3 d
5
4 s
1
Manganes
e
Mn 25 [ Ar ] 3 d
5
4 s
2
Iron Fe 26 [ Ar ] 3 d
6
4 s
2
Cobalt Co 27 [ Ar ] 3 d
7
4 s
2
Nickel Ni 28 [ Ar ] 3 d
8
4 s
2
Copper Cu
*
29 [ Ar ] 3 d
10
4 s
1
Zinc Zn 30 [ Ar ] 3 d
10
4 s
2
Table : 19.2 Second transition or 4 d - series
Element Symbol At. No. Electronic
configuratio
n
Yttrium Y 39
d
orbitals are filled up
[K r ] 4 d
1
5 s
2
Zirconium Zr 40 [ Kr ] 4 d
2
5 s
2
Niobium Nb
*
41 [ Kr ] 4 d
4
5 s
1
Molybdenum Mo
*
42 [ Kr ] 4 d
5
5 s
1
Technetium Tc 43 [ Kr ] 4 d
5
5 s
2
Ruthenium Ru
*
44 [ Kr ] 4 d
7
5 s
1
Rhodium Rh
*
45 [ Kr ] 4 d
8
5 s
1
Palladium Pd
*
46 [ Kr ] 4 d
10
5 s
0
Silver Ag
*
47 [ Kr ] 4 d
10
5 s
1
Cadmium Cd 48 [ Kr ] 4 d
10
5 s
2
Table : 19.3 Third transition or 5 d - series :
Element Symbo
l
At. No. Electronic
configuration
Lanthanu
m
La 57
d
orbitals
are filled
up
[ Xe ] 5 d
1
6 s
2
Hafnium Hf 72 [ Xe ] 4 f
14
5 d
2
6 s
2
Tantalum Ta 73 [ Xe ] 4 f
14
5 d
3
6 s
2
d and f - Block Elements
Chapter
Tungsten W 74 [ Xe ] 4 f
14
5 d
4
6 s
2
Rhenium Re 75 [ Xe ] 4 f
14
5 d
5
6 s
2
Osmium Os 76 [ Xe ] 4 f
14
5 d
6
6 s
2
Iridium Ir 77 [ Xe ] 4 f
14
5 d
7
6 s
2
Platinum Pt
*
78 [ Xe ] 4 f
14
5 d
10
6 s
0
Gold Au
*
79 [ Xe ] 4 f
14
5 d
10
6 s
!
Mercury Hg 80 [ Xe ] 4 f
14
5 d
10
6 s
2
Table : 19.4 Fourth transition or 6 d - series :
Element Symbo
l
At. No. Electronic
configuration
Actinium Ac 89
d
orbitals are filled up
[ Rn ] 6 d
1
7 s
2
Rutherfordiu
m
Rf 10
[ Rn ] 5 f
14
6 d
2
7 s
2
Hahnium Ha 10
[ Rn ] 5 f
14
6 d
3
7 s
2
Seaborgium Sg 10
[ Rn ] 5 f
14
6 d
4
7 s
2
Bohrium Bh 107 [ Rn ] 5 f
14
6 d
5
7 s
2
Hassium Hs 10
[ Rn ] 5 f
14
6 d
6
7 s
2
Meitnerium Mt 10
[ Rn ] 5 f
14
6 d
7
7 s
2
Ununnilium Uun 110 [ Rn ] 5 f
14
6 d
8
7 s
2
Unununium Uuu 111 [ Rn ] 5 f
14
6 d
9
7 s
2
Unubium Uub 112 [ Rn ] 5 f
14
6 d
10
7 s
2
Elements marked with asterisk have anomalous
configurations. These are attributed to factors like
nuclear-electron and electron-electron forces and
stability of half filled and full filled orbital.
All transition elements are d block elements but
all d block elements are not transition elements.
Physico-Chemical Properties of d-Block Elements
(1) Atomic radii : The atomic, radii of 3 d - series of
elements are compared with those of the neighbouring
s and p - block elements.
K Ca Sc Ti V Cr Mn
Fe Co Ni Cu Zn Ga Ge
The atomic radii of transition elements show the
following characteristics,
(i) The atomic radii and atomic volumes of d -
block elements in any series decrease with increase in
the atomic number. The decrease however, is not
regular. The atomic radii tend to reach minimum near
at the middle of the series, and increase slightly
towards the end of the series.
Explanation : When we go in any transition series
from left, to right, the nuclear charge increases
gradually by one unit at each elements. The added
electrons enter the same penultimate shell, (inner d -
shell). These added electrons shield the outermost
electrons from the attraction of the nuclear charge. The
increased nuclear charge tends to reduce the atomic
radii, while the added electrons tend to increase the
atomic radii. At the beginning of the series, due to
smaller number of electrons in the d - orbitals, the effect
of increased nuclear charge predominates, and the
atomic radii decrease. Later in the series, when the
number of d - electrons increases, the increased shielding
effect and the increased repulsion between the electrons
tend to increase the atomic radii. Somewhere in the
middle of the series, therefore the atomic radii tend to
have a minimum value as observed.
(ii) The atomic radii increase while going down in
each group. However, in the third transition series
from hafnium ( Hf ) and onwards, the elements have
atomic radii nearly equal to those of the second
transition elements.
Explanation : The atomic radii increase while
going down the group. This is due to the introduction of
an additional shell at each new element down the
group. Nearly equal radii of second and third transition
series elements is due to a special effect called
lanthanide contraction.
(2) Ionic radii : For ions having identical charges,
the ionic radii decrease slowly with the increase in the
atomic number across a given series of the transition
elements.
Elements
( m ):
Ionic
radius,( M
2+
)/ pm :
Pm :( M
3+
)/ pm :
Sc – 81
Ti 90 76
V 88 74
Cr 84 69
Mn 80 66
Fe 76 64
Co 74 63
Ni 72 –
Cu 69 –
Zn 74 –
Explanation : The gradual decrease in the values
of ionic radius across the series of transition elements
is due to the increase in the effective nuclear charge.
Outer Ele. Confi. and O. S. for 3 d- elements
Elements Outer
electronic
configuration
Oxidation states
Sc 3 d
1
4s
2
Ti 3 d
2
4s
2
V 3 d
3
4s
2
Cr 3 d
5
4s
1
- 1, + 2, + 3, + 4, + 5, + 6
Mn 3 d
5
4s
2
- 2, + 3, + 4, + 5, + 6, + 7
Fe 3 d
6
4s
2
Co 3 d
7
4s
2
Ni 3 d
8
4s
2
Cu 3 d
10
4s
1
Zn 3 d
10
4s
2
Explanation : The outermost electronic
configuration of the transition elements is ( n - 1) d
1 -
10
ns
2
. Since, the energy levels of ( n - 1) d and ns - orbitals
are quite close to each other, hence both the ns and ( n -
- d - electrons are available for bonding purposes.
Therefore, the number of oxidation states show by
these elements depends upon the number of d -
electrons it has. For example, Sc having a
configuration 3d
1
4s
2
may show an oxidation state
of + 2 (only s - electrons are lost) and + 3 (when d -
electron is also lost). The highest oxidation state
which an elements of this group might show is
given by the total number of ns and ( n - 1) d -
electrons.
The relative stability of the different oxidation
states depends upon the factors such as, electronic
configuration, nature of bonding, stoichiometry, lattice
energies and solvation energies. The highest oxidation
states are found in fluorides and oxides because
fluorine and oxygen are the most electronegative
elements. The highest oxidation state shown by any
transition metal is eight. The oxidation state of eight is
shown by Ru and Os.
An examination of the common oxidation states
reveals the following conclusions.
(i) The variable oxidation states shown by the
transition elements are due to the participation of outer
ns and inner ( n – 1) d - electrons in bonding.
(ii) Except scandium, the most common
oxidation state shown by the elements of first
transition series is +2. This oxidation state arises
from the loss of two 4 s electrons. This means that
after scandium, d - orbitals become more stable than
the s - orbital.
(iii) The highest oxidation states are observed in
fluorides and oxides. The highest oxidation state shown
by any transition elements (by Ru and Os ) is 8.
(iv) The transition elements in the + 2 and + 3
oxidation states mostly form ionic bonds. In compounds
of the higher oxidation states (compound formed with
fluorine or oxygen), the bonds are essentially covalent.
For example, in permanganate ion MnO 4
, all bonds
formed between manganese and oxygen are covalent.
(v) Within a group, the maximum oxidation state
increases with atomic number. For example, iron
shown the common oxidation state of + 2 and + 3, but
ruthenium and osmium in the same group form
compounds in the + 4, + 6 and + 8 oxidation states.
(vi) Transition metals also form compounds in
low oxidation states such as +1 and 0. For example,
nickle in, nickel tetracarbonyl, Ni ( CO ) 4 has zero
oxidation state. Similarly Fe in
5
( Fe ( CO ) has zero
oxidation state.
The bonding in the compounds of transition
metals in low oxidation states is not always very
simple.
(vii) Ionisation energies and the stability of
oxidation states :The values of the ionisation energies
can be used in estimating the relative stability of
various transition metal compounds (or ions). For
example, Ni
2+
compounds are found to be
thermodynamically more stable than Pt
2+
, whereas Pt
4+
compounds are more stable than Ni
4+
compounds. The
relative stabilities of Ni
2+
relative to Pt
2+
and that of
Pt
4+
relative to Ni
4+
can be explained as follows,
The first four ionisation energies of Ni and Pt
Metal ( IE 1
)
kJmol
,
( IE 3
)
kJmol
,
Etotal, kJ mol
(= IE 1
IE 4
)
Ni 2490 8800 11290
Pt 2660 6700 9360
Thus, the ionisation of Ni to Ni
2+
requires lesser
energy (2490 kJ mol
) as compared to the energy
required for the production of Pt
2+
(2660 kjmol
- 1
Therefore, Ni
2+
compounds are thermodynamically
more stable than Pt
2+
compounds.
On the other hand, formation of Pt
4+
requires
lesser energy (9360 kJ mol
) as compared to that
required for the formation of Ni
4+
(11290 kJ mol
Therefore, Pt
4+
compounds are more stable than Ni
4+
compounds.
This is supported by the fact that [ PtCl 6 ]
2 –
complex ion is known, while the corresponding ion for
nickel is not known. However, other factors which
affect the stability of a compound are,
(a) Enthalpy of sublimation of the metal.
(b) Lattice and the solvation energies of the
compound or ion.
(viii) Transition elements like Sc , Y , La and Ac do
not show variable valency.
(8) Electrode potentials ( E
o
) : Standard electrode
potentials of some half–cells involving 3 d - series of
transition elements and their ions in aqueous solution
are given in table,
Standard electrode potentials for 3 d - elements
Elements Ion Electrode reaction E °/ volt
Sc Sc
3+
Sc
3+
Sc – 2.
Ti Ti
2+
Ti
2+
Ti – 1.
V V
2+
V
2+
V
Cr Cr
3+
Cr
3+
Cr – 0.
Mn Mn
2+
Mn
2+
Mn – 1.
Fe Fe
2+
Fe
2+
Fe
Co Co
2+
Co
2+
Co – 0.
Ni Ni
2+
Ni
2+
Ni – 0.
Cu Cu
2+
Cu
2+
Cu
Zn Zn
2+
Zn
2+
Zn – 0.
The negative values of E ° for the first series of
transition elements (except for Cu
2+
/ Cu ) indicate that,
(i) These metals should liberate hydrogen from
dilute acids i.e., the reactions,
M + 2 H
M
2+
+ H
2
( g ); 2 M + 6 H
2 M
3+
3 H 2 ( g )
are favourable in the forward direction. In actual
practice however, most of these metals react with
dilute acids very slowly. Some of these metals get
coated with a thin protective layer of oxide. Such an
oxide layer prevents the metal to react further.
(ii) These metals should act as good reducing
agents. There is no regular trend in the E ° values. This
is due to irregular variation in the ionisation and
sublimation energies across the series.
Relative stabilities of transition metal ions in
different oxidation states in aqueous medium can be
predicted from the electrode potential data. To
illustrate this, let us consider the following,
M ( s ) M ( g ) ;
1
H Enthalpy of sublimation,
H sub
M ( g ) M ( g ) e ; H Ionisationenergy, IE
2
M ( g ) M ( aq )
hyd
H Enthalpyofhydration, H
3
Adding these equations one gets,
M ( s ) M ( aq ) e
sub hyd
H H H H H IE H
1 2 3
The H represents the enthalpy change required
to bring the solid metal M to the monovalent ion in
aqueous medium, M
(aq).
The reaction, M ( s ) M
(aq) + e
, will be
favourable only if H is negative. More negative is the
value is of H , more favourable will be the formation
of that cation from the metal. Thus, the oxidation state
for which H value is more negative will be stable in
the solution.
Electrode potential for a M
n +
/ M half-cell is a
measure of the tendency for the reaction, M
n +
(aq) + ne
M ( s )
Thus, this reduction reaction will take place if
the electrode potential for M
n +
/ M half- cell is positive.
The reverse reaction, M ( s ) M
n +
(aq) + ne
Involving the formation of M
n +
(aq) will occur if
the electrode potential is negative, i.e ., the tendency
for the formation of M
n +
(aq) from the metal M will be
more if the corresponding E ° value is more negative. In
other words, the oxidation state for which E° value is
more negative (or less positive) will be more stable in
the solution.
When an elements exists in more than one
oxidation states, the standard electrode potential ( E °)
values can be used in the predicting the relative
stabilities of different oxidation states in aqueous
solutions. The following rule is found useful.
The oxidation state of a cation for which
sub hyd
H H lE H or E° is more negative (for less
positive) will be more stable.
(9) Formation of coloured ions : Most of the
compound of the transition elements are coloured in
the solid state and /or in the solution phase. The
compounds of transition metals are coloured due to the
presence of unpaired electrons in their d - orbitals.
Explanation : In an isolated atom or ion of a
transition elements, all the five d - orbitals are of the
same energy (they are said to be regenerate). Under
the influence of the combining anion ( s ), or electron-
rich molecules, the five d - orbitals split into two (or
sometimes more than two) levels of different energies.
The difference between the two energy levels depends
upon the nature of the combining ions, but corresponds
to the energy associated with the radiations in the
visible region, ( 380 760 nm ). Typical splitting for
octahedral and tetrahedral geometries are shown in fig.
d-
orbitals
d xy
d yz
d zx
d xy
d yz
d zx
Octahexial field Tetrahedral
field
No field
The splitting of d - orbital energy levels in (a)
an octahedral, (b) a tetrahedral, geometry. This
dx
2
2
dz
2
dz
2
dx
2
2
called diamagnetic substances. Paramagnetism is due
to the presence of unpaired electrons in atoms, ions or
molecules.
The magnetic moment of any transition element
or its compound/ion is given by (assuming no
contribution from the orbital magnetic moment).
S S BM n n BM
s
where, S is the total spin ( n s ): n is the number of
unpaired electrons and s is equal to ½ (representing
the spin of an unpaired electron).
From the equation given above, the magnetic
moment( )
s
increases with an increase in the number
of unpaired electrons.
Magnetic moments of some ions of the 3 d-series
elements
Ion Outer
configuratio
n
No. of
unpaired
electrons
Magnetic moment
(BM)
Calculated observe
d
Sc
3+
3 d
0
0 0 0
Ti
3+
3 d
1
1 1.73 1.
Ti
2+
3 d
2
2 2.84 2.
V
2+
3 d
3
3 3.87 3.
Cr
2+
3 d
4
4 4.90 4.
Mn
2+
3 d
5
5 5.92 5.
Fe
2+
3 d
6
4 4.90 5.0-5.
Co
2+
3 d
7
3 3.87 4.4-5.
Ni
2+
3 d
8
2 2.84 2.9-3.
Cu
2+
3 d
9
1 1.73 1.4-2.
Zn
2+
3 d
10
0 0 0
In d - obitals belonging to a particular energy level,
there can be at the maximum five unpaired electrons in
d
5
cases. Therefore, paramagnetism in any transition
series first increases, reaches a maximum value for d
5
cases and then decreases thereafter.
(11) Formation of complex ions : Transition
metals and their ions show strong tendency for
complex formation. The cations of transition elements
( d - block elements) form complex ions with certain
molecules containing one or more lone-pairs of
electrons, viz., CO, NO, NH 3 etc., or with anions such
as, F
, Cl
, CN
etc. A few typical complex ions are,
2
2 6
2
34
4
6
[ Fe ( CN )] ,[ Cu ( NH )] ,[ Y ( HO )] ,
3
6
3
4 36
[ Ni ( CO )],[ Co ( NH )] [ FeF ]
Explanation : This complex formation tendency is
due to,
(i) Small size and high nuclear charge of the
transition metal cations.
(ii) The availability to vacant inner d - orbitals of
suitable energy.
(12) Formation of interstitial compounds :
Transition elements form a few interstitial compounds
with elements having small atomic radii, such as
hydrogen, boron, carbon and nitrogen. The small atoms
of these elements get entrapped in between the void
spaces (called interstices) of the metal lattice. Some
characteristics of the interstitial compound are,
(i) These are non-stoichiometric compounds and
cannot be given definite formulae.
(ii) These compounds show essentially the same
chemical properties as the parent metals, but differ in
physical properties such as density and hardness. Steel
and cast iron are hard due to the formation of
interstitial compound with carbon. Some non-
stoichimetric compounds are, VSe
(Vanadium
selenide), Fe 0.94 O and titanium nitride.
Explanation : Interstital compounds are hared
and dense. This is because, the smaller atoms of lighter
elements occupy the interstices in the lattice, leading to
a more closely packed structure. Due to greater
electronic interactions, the strength of the metallic
bonds also increases.
(13) Catalytic properties : Most of the transition
metals and their compounds particularly oxides have
good catalytic properties. Platinum, iron, vanadium
pentoxide, nickel, etc., are important catalysts.
Platinum is a general catalyst. Nickel powder is a good
catalyst for hydrogenation of unsaturated organic
compound such as, hydrogenation of oils some typical
industrial catalysts are,
(i) Vanadium pentoxide ( V 2 O 5 ) is used in the
Contact process for the manufacture of sulphuric acid,
(ii) Finely divided iron is used in the Haber’s
process for the synthesis of ammonia.
Explanation : Most transition elements act as
good catalyst because of,
(i) The presence of vacant d - orbitals.
(ii) The tendency to exhibit variable oxidation
states.
(iii) The tendency to form reaction intermediates
with reactants.
(iv) The presence of defects in their crystal
lattices.
(14) Alloy formation : Transition metals form
alloys among themselves. The alloys of transition
metals are hard and high metals are high melting as
compared to the host metal. Various steels are alloys of
iron with metals such as chromium, vanadium,
molybdenum, tungsten, manganese etc.
Explanation : The atomic radii of the transition
elements in any series are not much different from
each other. As a result, they can very easily replace
each other in the lattice and form solid solutions over
an appreciable composition range. Such solid solutions
are called alloys.
(15) Chemical reactivity : The d - block elements
(transition elements) have lesser tendency to react, i.e. ,
these are less reactive as compared to s - block elements.
Explanation : Low reactivity of transition
elements is due to,
(i) Their high ionisation energies.
(ii) Low heats of hydration of their ions.
(iii) Their high heats of sublimation.
Chromium containing compounds
Potassium dichromate, ( K 2 Cr 2 O 7 )
Potassium dichromate is one of the most
important compound of chromium, and also among
dichromates. In this compound Cr is in the hexavalent
(+6) state.
Preparation : It can be prepared by any of the
following methods,
(i) From potassium chromate : Potassium
dichromate can be obtained by adding a calculated
amount of sulphuric acid to a saturated solution of
potassium chromate.
KSO HO
orange
potassiumdichromate
HSO KCrO
yellow
K CrO
potassium chromate
2 4 2 4 2 2 7 2 4 2
2
K 2 Cr 2 O 7 Crystals can be obtained by concentrating
the solution and crystallisation.
(ii) Manufacture from chromite ore : K 2 Cr 2 O 7 is
generally manufactured from chromite ore ( FeCr 2 O 4 ).
The process involves the following steps.
(a) Preparation of sodium chromate : Finely
powdered chromite ore is mixed with soda ash and
quicklime. The mixture is then roasted in a
reverberatory furnace in the presence of air. Yellow
mass due to the formation of sodium chromate is
obtained.
2 4 2 2 3 2 3
4 FeCr O O 2 FeO 4 CrO
2 3 2 3 2 2 4 2
Cr O NaCO O NaCrO CO g
sodium chromate
FeCr O NaCO O FeO CO g Na CrO
2 4 2 3 2 2 3 2 2 4
The yellow mass is extracted with water, and
filtered. The filtrate contains sodium chromate.
The reaction may also be carried out by using
NaOH instead of Na 2 CO 3. The reaction in that case is,
FeCr O NaOH O NaCrO FeO HO
2 4 2 2 4 2 3 2
(b) Conversion of chromate into dichromate :
Sodium chromate solution obtained in step (a) is
treated with concentrated sulphuric acid when it is
converted into sodium dichromate.
Na CrO HSO NaCrO NaSO HO
sodium chromate sodiumdichromate
2 4 2 4 2 2 7 2 4 2
On concentration, the less soluble sodium
sulphate, Na 2
SO
4
.10 H
2
O crystallizes out. This is filtered
hot and allowed to cool when sodium dichromate,
Na 2 Cr 2 O 7 .2 H 2 O , separates out on standing.
(c) Concentration of sodium dichromate to
potassium dichromate : Hot concentrated solution of
sodium dichromate is treated with a calculated amount
of potassium chloride. When potassium dichromate
being less soluble crystallizes out on cooling.
Na CrO KCl KCrO NaCl
soddichromate potdichromate
.
2 2 7
.
2 2 7
Physical properties
(i) Potassium dichromate forms orange-red
coloured crystals.
(ii) It melts at 699 K.
(iii) It is very stable in air (near room
temperature) and is generally, used as a primary
standard in the volumetric analysis.
(iv) It is soluble in water though the solubility is
limited.
Chemical properties
(i) Action of heat : Potassium dichromate when
heated strongly. Decomposes to give oxygen.
2 2 7 2 4 2 3 2
4 K CrO s 4 KCrO ( s ) 2 CrO ( s ) 3 O
(ii) Action of acids
(a) In cold, with concentrated H 2 SO 4 , red crystals
of chromium trioxide separate out.
K CrO aq concHSO KHSO aq CrO s H O
2 2 7 2 4 4 3 2
On heating a dichromate-sulphuric acid mixture,
oxygen gas is given out.
2 2 7 2 4 2 4 2 43 2 2
2 K CrO 8 HSO 2 KSO 2 Cr ( SO ) 8 HO 3 O
(b) With HCl , on heating chromic chloride is
formed and Cl 2 is liberated.
K CrO HClaq CrCl KClaq HO Cl g 2 27 aq 3 aq 2 2
14 2 2 7 3
(iii) Action of alkalies : With alkalies, it gives
chromates. For example, with KOH ,
K CrO KOH KCrO HO
orange yellow
2 2 4 2 4 2
On acidifying, the colour again changes to orange-
red owing to the formation of dichromate.
Manganese containing compound
Potassium Permanganate, ( KMnO 4 )
Potassium permanganate is a salt of an unstable
acid HMnO 4
(permanganic acid). The Mn is an +7 state
in this compound.
Preparation : Potassium permanganate is
obtained from pyrolusite as follows.
Conversion of pyrolusite to potassium
manganate : When manganese dioxide is fused with
potassium hydroxide in the presence of air or an
oxidising agent such as potassium nitrate or chlorate,
potassium manganate is formed, possibly via potassium
manganite.
2 4 ] 2
2 2 3 2
MnO KOH KMnO HO
potassiummanganite
fused
K MnO O KMnO HO
2 3 2 2 4 2
HO
dark greenmass
potassiummanganate
KOH O KMnO
pyrolusite
MnO
fused
2 2 2 4 2
Oxidation of potassium manganate to potassium
permanganate : The potassium manganate so obtained
is oxidised to potassium permanganate by either of the
following methods.
By chemical method : The fused dark-green mass
is extracted with a small quantity of water. The filtrate
is warmed and treated with a current of ozone, chlorine
or carbon dioxide. Potassium manganate gets oxidised
to potassium permanganate and the hydrated
manganese dioxide precipitates out. The reactions
taking place are,
When CO 2
is passed
K MnO HO KMnO MnO KOH
potassiummanganate potassiumpermanganate
3 2 2 4
2 4 2 4 2
CO KOH KCO HO
2 2 3 2
When chlorine or ozone is passed
2 K MnO Cl 2 KMnO 2 KCl
2 4 2 4
K MnO O HO KMnO KOH O g
2 4 3 2 4 2
The purple solution so obtained is concentrated
and dark purple, needle-like crystals having metallic
lustre are obtained.
Electrolytic method : Presently, potassium
manganate ( K 2
MnO 4
) is oxidised electrolytically. The
electrode reactions are,
At anode:
MnO MnO e
green purple
4
2
4
At cathode:
H e H g
2
2 2
The purple solution containing KMnO 4
is
evaporated under controlled condition to get crystalline
sample of potassium permanganate.
Physical properties
KMnO 4
crystallizes as dark purple crystals with
greenish luster (m.p. 523 K ).
It is soluble in water to an extent of 6.5 g per 100 g
at room temperature. The aqueous solution of KMnO 4
has a purple colour.
Chemical properties : Some important chemical
reactions of KMnO 4 are given below,
Action of heat : KMnO 4 is stable at room
temperature, but decomposes to give oxygen at higher
temperatures.
KMnO s KMnO s MnO O g
heat
4 2 4 2 2
2
Oxidising actions : KMnO 4
is a powerful agent in
neutral, acidic and alkaline media. The nature of
reaction is different in each medium. The oxidising
character of KMnO 4 (to be more specific, of
4
MnO ) is
indicated by high positive reduction potentials for the
following reactions.
Acidic medium
MnO 8 H 5 e
4
Mn HO E V
o
4 1. 51
2
2
Alkaline medium
MnO 2 H O 3 e
4 2
⇌ MnO OH E V
o
4 1. 23
2
In strongly alkaline solutions and with excess of
4
MnO , the reaction is
MnO e
4
MnO E V
o
- 56
2
4
There are a large number of oxidation-reduction
reactions involved in the chemistry of manganese
compounds. Some typical reactions are,
In the presence of excess of reducing agent in
acidic solutions permanganate ion gets reduced to
manganous ion, e.g.,
Fe MnO H Fe Mn HO
2
3 2
4
2
5 8 5 4
An excess of reducing agent in alkaline solution
reduces permanganate ion only to manganese dioxide
e.g.,
NO MnO OH NO MnO HO
2 4 3 2 2
3 2 3
O
Cr
O
O
O
2 –
Chromate ion
O
Cr
O
O
O
131°
Cr
O
O
O
180
pm
161
pm
Dichromate
ion
The structure of chromate ( CrO
2 –
4
) and dichromate
( Cr 2
O 7
2 –
) ions
2 –
In faintly acidic and neutral solutions, manganous
ion is oxidised to manganese oxidised to manganese
dioxide by permanganate.
2 MnO 3 Mn 2 HO 5 MnO 4 H
2 2
2
4
In strongly basic solutions, permangante oxidises
manganese dioxide to manganate ion.
MnO MnO OH MnO HO
2
2
2 4 4
2 4 3 2
In acidic medium, KMnO 4 oxidises,
Ferrous salts to ferric salts
2 KMnO 3 HSO KSO 2 MnSO 3 HO 5 O
4 2 4 2 4 4 2
2 ] 5
2 3
4 2 4 2 4
FeSO HSO O Fe SO HO
KMnO HSO FeSO KSO MnSO Fe SO H O
2 3
4 2 4 4 2 4 4 2 4
2 8 10 2 5 8
Ionic equation
MnO H Fe Mn Fe HO
2
2 2 3
4
2 16 10 2 10 8
The reaction forms the basis of volumetric
estimation of Fe
2+
in any solution by KMnO 4.
Oxalic acid to carbon dioxide
2 KMnO 3 HSO KSO 2 MnSO 3 HO 5 O 4 2 4 2 4 4 2
2 ] 5
2 2 2
COOH O CO HO
KMnO HSO COOH KSO MnSO CO H O
4 2 4 2 2 4 4 2 2
2 3 5 2 10 8
Ionic equation
MnO H COOH Mn CO H O
2 2
2
4 2
2 6 5 2 10 8
Sulphites to sulphates
2 KMnO 3 HSO KSO 2 MnSO 3 HO 5 O
4 2 4 2 4 4 2
] 5
2 3 2 4
NaSO O NaSO
KMnO HSO NaSO KSO MnSO NaSO HO 4 2 4 2 3 2 4 4 2 4 2
2 3 5 2 5 3
Ionic equation
MnO H SO Mn SO HO
2
2
4
2 2
4 3
2 6 5 2 5 3
Iodides to iodine in acidic medium
2 KMnO 3 HSO KSO 2 MnSO 3 HO 5 O
4 2 4 2 4 4 2
2 2 5
2 2
KI HO O I KOH
2 2 ] 5
2 4 2 4 2
KOH HSO KSO HO
KMnO HSO KI KSO MnSO I HO
4 2 4 2 4 4 2 2
Ionic equation
MnO H I Mn I HO
2 2
2
4
2 16 10 2 5 8
Hydrogen peroxide to oxygen
2 KMnO 3 HSO KSO 2 MnSO 3 HO 5 O
4 2 4 2 4 4 2
5
2 2 2 2
HO O HO O
4 2 4 2 2 2 4 4 2 2
2 KMnO 3 HSO 5 HO KSO 2 MnSO 8 HO 5 O
Manganous sulphate ( MnSO 4
) to manganese
dioxide ( MnO 2 )
2 KMnO HO 2 KOH 2 MnO 3 O
4 2 2
3
4 2 2 2 4
MnSO HO O MnO HSO
KOH HSO KSO HO
2 4 2 4 2
4 4 2 2 2 4 2 4
2 KMnO 3 MnSO 2 HO 5 MnO KSO 2 HSO
Ionic equation
2 MnO 3 Mn 2 HO 5 MnO 4 H
2 2
2
4
Ammonia to nitrogen
2 KMnO HO 2 MnO 2 KOH 3 O
4 2 2
NH O N g H O
3 2 2
KMnO NH MnO KOH HO N g
4 3 2 2 2
Uses : KMnO 4 is used,
(i) As an oxidising agent. (ii) As a disinfectant
against disease-causing germs. (iii) For sterilizing
wells of drinking water. (iv) In volumetric estimation
of ferrous salts, oxalic acid etc. (v) Dilute alkaline
4
KMnO solution known as Baeyer’s reagent.
Structure of Permanganate Ion (MnO 4
-
) : Mn in
MnO 4
is in +7 oxidation state. Mn
7+
exhibits sp
3
hybridisation in this ion. The structure of MnO 4
is,
shown in fig.
Iron and its Compounds
(1) Ores of iron : Haematite
2 3
FeO
, Magnetite
3 4
FeO Limonite (. 3 )
2 3 2
FeO HO , Iron pyrites ( ),
2
FeS
Copper pyrities( )
2
CuFeS etc.
(2) Extraction : Cast iron is extracted from its
oxides by reduction with carbon and carbon monoxide
in a blast furnace to give pig iron.
Roasting : Ferrous oxide convert into ferric
oxide.
Fe O HO FeO HO
2 3 2 2 3 2
3 2
2 FeCO 2 FeO 2 CO
2 2 3
4 FeO O 2 FeO
Smelting : Reduction of roasted ore of ferric
oxide carried out in a blast furnace.
(i) The reduction of ferric oxide is done by carbon
and carbon monoxide (between 1473 k to 1873 k )
O
O O
-
O
Mn
Strucutre of MnO 4
-
ion
Cold conc.
3
HNO
makes iron passive due to the
deposit of a thin layer of iron oxide ( )
3 4
FeO
on the
surface.
Hot conc.
3
HNO reacts with iron liberating NO.
Fe HNO FeNO NO HO
3 33 2
4 (hotconc.) ( ) 2
(4) Iron does not react with alkalies.
(5) It displaces less electropositive metals ( e.g. ,
Cu , Ag etc.) from their salts
CuSO Fe FeSO Cu
4 4
(6) Finely divided iron combines with CO forming
penta carbonyl
5
Fe 5 CO Fe ( CO )
(7) Iron does not form amalgam with Hg.
(8) Iron is the most abundant and most widely
used transition metal.
Compounds of iron
(1) Oxides of Iron : Iron forms three oxides
2 3
FeO , FeO (Haematite),
3 4
FeO (magnetite also called
magnetic oxide or load stone).
(i) Ferrous oxide, FeO :It is a black powder, basic in
nature and reacts with dilute acids to give ferrous salts.
FeO HSO FeSO HO
2 4 4 2
; It is used in glass
industry to impart green colour to glass.
(ii) Ferric oxide :
2 3
FeO It is a reddish brown
powder, not affected by air or water; amphoteric in
nature and reacts both with acids and alkalis giving
salts. It can be reduced to iron by heating with C or CO.
Fe O 3 C 2 Fe 3 CO
2 3
2 3 2
Fe O 3 CO 2 Fe 3 CO
It is used as red pigment to impart red colour to
external walls and as a polishing powder by jewellers.
(iii) Ferrosoferricoxide (. ):
3 4 2 3
FeO FeOFeO It is
more stable than FeO and ,
2 3
FeO
magnetic in nature
and dissolves in acids giving a mixture of iron (II) and
iron (III) salts.
3 4 2 4
Fe O 4 HSO
(dil) FeSO Fe SO HO
4 2 43 2
(2) Ferrous sulphide FeS : It is prepared by
heating iron filing with sulphur. With dilute , 2 4
HSO it
gives. 2
HS
2 4
FeS HSO (dil) FeSO HS
4 2
(3) Ferric chloride :
3
FeCl (i) preparation : It is
prepared by treating 3
Fe ( OH )with HCl
Fe OH HCl FeCl HO
3 3 2
The solution on evaporation give yellow crystals
of FeCl HO 3 2
(ii) Properties : (a) Anhydrous
3
FeCl forms
reddish-black deliquescent crystals.
(b)
3
FeCl is hygroscopic and dissolves in HO
2
giving brown acidic solution due to formation of HCl
FeCl 3 HO Fe ( OH ) 3 HCl
(Brown)
3 2 3
(c) Due to oxidising nature
3
Fe ions
3
FeCl is
used in etching metals such as copper
3 2 2
Fe Cu Fe Cu aq
(d) In vapour state
3
FeCl exists as a dimer,
2 6
FeCl
Cl Cl Cl
Fe Fe
Cl Cl Cl
(e)
3
FeCl is used as stypic to stop bleeding from a
cut.
(4) Ferrous sulphate,
4
FeSO
HO
2
(Green
vitriol) : It is prepared as follow ,
2 4 4 2
Fe HSO FeSO H
(i) One pressure to moist air crystals become
brownish due to oxidation by air.
4 2 2 4
4 FeSO 2 HO O 4 Fe ( OH ) SO
(ii) On heating, crystals become anhydrous and on
strong heating it decomposes to
2 3 2
FeO , SO
and
3
SO
FeSO HO FeSO HO
4 2
heat
4 2
. 7 7
2 3 2 3
heating
4
2 FeSO FeO SO SO
Strong
(iii) It can reduce acidic solution of
4
KMnO
and
2 2 7
KCrO
(iv) It is generally used in double salt with
ammonium sulphate.
NH SO FeSO HO FeSO NH SO HO
4 2
Mohr's salt
4 2 4 4 2 4 42
Mohr’s salt is resistant to atmospheric oxidation.
(v) It is used in the ring test for nitrate ions
where it gives brown coloured ring of compound
4
FeSO NO
FeSO NO FeSO. NO
4 4
(vi)
4
FeSO is used in manufacture of blue black
ink.
(vii)
4 2 2
FeSO HO
is known as a name of
Fenton’s reagent.
(5) Mohr's salt. ( ). 6 :
4 42 4 2
FeSO NH SO HO It is a
double salt and is prepared by crystallising a solution
containing equivalent amounts of FeSO HO
4 2
. 7 and
4 2 4
( NH ) SO. It may be noted that Mohr’s salt contains
only
2
Fe ions without any trace of
3
Fe ions. In
contrast FeSO HO
4 2
always contains some
3
Fe ions
due to aerial oxidation of
2
Fe ions. Mohr salt is,
therefore, used as a primary standard in volumetric
analysis since a standard solution of
2
Fe ions can be
obtained directly by weighing a known amount of the
Mohr salt.
It acts as a reducing agent and as such reduces
acidified 4
KMnO and
2 2 7
KCrO solutions.
MnO Fe H Fe Mn HO
2
2 3 2
4
5 8 5 4
Cr O Fe H Fe Cr HO
2
3 3
(Frommohrssalt)
2 2
2 7
Copper and its Compounds
(1) Ores : Copper pyrites (chalcopyrite) ,
2
CuFeS
Cuprite (ruby copper) ,
2
CuO Copper glance ( )
2
CuS ,
Malachite[ ( ). ], 2 3
CuOH CuCO Azurite[ ( ). 2 ]
2 3
CuOH CuCO
(2) Extraction : Most of the copper (about 75%)
is extracted from its sulphide ore, copper pyrites.
Concentration of ore : Froth floatation process.
Roasting : Main reaction :
2 2 2 2
2 CuFeS O CuS 2 FeS SO.
Side reaction :
2 2 2 2
2 Cu S 3 O 2 CuO 2 SO
2 2
2 FeS 3 O 2 FeO 2 SO.
Smelting : (slag)
2 3
FeO SiO FeSiO
Cu O FeS FeO CuS
2 2
The mixture of copper and iron sulphides melt
together to form ' matte '( ) 2
Cu S FeS and the slag floats
on its surface.
Conversion of matte into Blister copper
(Bessemerisation) : Silica is added to matte and a hot
blast of air is passed (slag)
2 3
FeO SiO FeSiO. Slag is
removed. By this time most of iron sulphide is
removed.
2 2 2
Cu S 2 CuO 6 Cu SO
Blister copper : Which contain about 98% pure
copper and 2% impurities ( Ag , Au , Ni , Zn etc.)
Properties of copper : It has reddish brown
colour. It is highly malleable and ductile. It has high
electrical conductivity and high thermal conductivity.
Copper is second most useful metal (first being iron). It
undergoes displacement reactions with lesser reactive
metals e.g. with Ag. It can displace Ag from 3
AgNO. The
finally divided Ag so obtained is black in colour.
Copper shows oxidation states of +1 and +2. Whereas
copper (I) salts are colourless, copper (II) salts are blue in
colour. Cu (I) salts are less stable and hence are easily
oxidised to Cu (II) salts( 2 )
2
Cu Cu Cu
. This reaction
is called disproportionation.
(1) In presence of atmospheric
2
CO and moisture,
copper gets covered with a green layer of basic copper
carbonate (green layer) which protects the rest of the
metal from further acton.
(green lay er)
2 2 2 2 3
Cu O CO HO Cu ( OH ) CuCO
(2) In presence of oxygen or air, copper when
heated to redness (below 1370 K ) first forms red
cuprous oxide which changes to black cupric oxide on
further heating. If the temperature is too high, cupric
oxide changes back to cuprous oxide
Above 1370 (Black)
O
(Red)
2
Below 1370 K
2
2
Cu O CuO CuO
K
CuO Cu
Hightemp.
Cu O
2
(3) Action of acids. Non oxidising dil. acids such
as
2 4
HCl , HSO have no action on copper. However,
copper dissolves in these acids in presence of air.
Cu HCl CuCl HO
2 2 2
O(air)
2
1
2
With dil.
3
HNO
Cu
liberates NO
(nitric oxide)
Cu HNO CuNO NO HO
3 32 2
With conc.
3
HNO , copper gives
2
NO
Cu HNO CuNO NO HO
3 32 2 2
With hot conc.
2 4
HSO , copper gives
2
SO
Cu HSO CuSO SO HO
2 4 4 2 2
Compounds of Copper
(1) Halides of copper : Copper (II) chloride,
2
CuCl is prepared by passing chlorine over heated
copper. Concentrated aqueous solution of
2
CuCl is dark
brown but changes first to green and then to blue on
dilution.
On heating, it disproportionates to copper (I)
chloride and chlorine
2
Heat
2
2 CuCl 2 CuCl Cl
It is used as a catalyst in the Daecon’s process for
the manufacture of chlorine.
Copper (I) chloride, CuCl is a white solid
insoluble in water. It is obtained by boiling a solution
of
2
CuCl
with excess of copper turnings and conc. HCl
CuCl Cu 2 CuCl
HCl
2
It dissolves in conc. HCl due to the formation of
complex [ ]
2
HCuCl
[ ]
2
CuCl HCl HCuCl
It is used as a catalyst alongwith NH Cl
4
in the
preparation of synthetic rubber.
(2) Cuprous oxide Cu O
2
: It is a reddish brown
powder insoluble in water but soluble in ammonia
solution, where it forms diammine copper (I) ion.
2 [ ( )]
3 32
Cu NH CuNH. It is used to impart red colour
to glass in glass industry.
(3) Cupric oxide CuO : It is dark black,
hygroscopic powder which is reduced to Cu by
hydrogen, CO etc. It is used to impart light blue colour
to glass. It is prepared by heating copper nitrate.
3 2 2 2
2 Cu ( NO ) 2 CuO 4 NO O
forming NO
and
2
NO
respectively. Chlorine also reacts
with Ag to form AgCl.
2 Ag Cl 2 AgCl
2
Hot conc.
2 4
HSO reacts with Ag forming
2
SO
like Cu
Compounds of Silver:
(1) Silver oxide ( )
2
AgO
: It is unstable and
decomposes into Ag and 2
O on slow heating.
2 2
2 Ag O 4 Ag O
(2) Silver halides ( AgF, AgCl, AgBr and Agl ) :
Only AgF is soluble in HO 2
. AgCl is insoluble in HO
2
but dissolves in NHOH 4
2 2 3
NaSO and KCN solutions.
AgBr is partly soluble whereas Agl is completely
insoluble in NHOH 4
. Except AgF , all the remaining
three silver halides are photosensitive.
Diamine silver(I)chloride
4 32 2
AgCl 2 NHOH [ Ag ( NH )] Cl 2 HO
Pot. Dicy anoargentate(I)
2
AgCl 2 KCN K [ Ag ( CN )] KCl
Sod. Dithiosulphatoargentate(I)
2 2 3 3 2 32
AgCl 2 NaSO Na [ Ag ( SO )] NaCl
(3) Silver nitrate ( AgNO 3
) : Silver nitrate
( AgNO 3 ) is called lunar caustic silver nitrate on heating
above its m.p. (485 K ) decomposes to silver nitrite but
on heating to red heat gives silver.
2 2
Above 485 K
3
2 AgNO 2 AgNO O
2 2
Redheat
3
2 AgNO 2 Ag 2 NO O
When treated with alkali ,
3
AgNO
forms silver
oxide which in case of NHOH 4
dissolves to form
complex ion.
AgNO NaOH AgO NaNO HO
3 2 3 2
AgNO NHOH AgO NHNO HO
3 4 2 4 3 2
Ag O NHOH AgNH OH HO
2
Diamine silverhy droxide
2 4 32
4 2 [ ( )] 3
3
AgNO reacts with iodine in two ways
3
6 AgNO (excess) + l HO
2 2
3 3
AglO 5 AgI 6 HNO
3 2
5 AgNO 3 l (excess)
2 3 3
3 H O HIO 5 Agl 5 HNO
In contact with organic matter (skin, cloth, paper
etc.) 3
AgNO is reduced to metallic silver (black)
2 2 2 [ ]
3 2 3
AgNO HO Ag HNO O
oxidises
organic matter
3
AgNO gives different coloured ppt. with
different anions like
3
4
2
4
2
2 3
2
Cl , Br , I , S , SO , CrO , PO
etc.
3
AgNO
is used in the preparation of ink and hair
dyes.
Photography : The photographic plate is coated
with a colloidal gelatinised solution of AgBr. During
exposure, AgBr is reduced to metallic silver.
2
2 AgBr 2 Ag Br
The exposed film is developed. The developer
used is an alkaline solution of hydroquinone or quinol
which reduces some of the exposed AgBr to black silver.
Quinone
Quinol
6 4 2 6 4 2
C H OH AgBr Ag CHO HBr
The film is finally fixed by dipping in a solution of
sodium thiosulphate or hypo which removes unchanged
AgBr as complex ion.
AgBr 2 NaSO Na [ Ag ( SO )] NaBr
2 2 3 3 2 32
After taking a print of the photograph it is finally
toned by dipping in a dilute solution of gold chloride to
impart a beautiful golden colour or it is dipped in
potassium chloro platinate
2 6
KPtCl solution to get a
shining grey tinge.
AuCl 3 Ag 3 AgCl Au
3
Gold and its Compounds
(1) Occurrence of gold : Gold is mainly found in
native state either as vein gold, placer gold or alluvial
gold. It is also present to a small extent in the
combined state as sulphide, telluride and
arsenosulphide. It is considered to be the king of metal.
Some important ores of gold are:
(i) Calaverite, AuTe 2 (ii) Sylvanite, AuAgTe 2 and
(iii) Bismuth aurite,
2
BiAu
(2) Extraction of gold : (i) Mac-Arthur-Forest
Cyanide process : The powdered gold ore, after
concentration by Froth-floatation process , is roasted to
remove easily oxidisable impurities of tellurium,
arsenic and sulphur. The roasted ore is then treated
with a dilute solution of KCN in presence of
atmospheric oxygen when gold dissolves due to the
formation of an aurocyanide complex.
Au KCN HO O KAuCN 4 KOH
Solution
4 8 2 4 [ ( )]
2 2 2
The metal is then extracted by adding zinc dust.
ppt.
2 [ ( )] [ ( )] 2
2 2 4
KAuCN Zn K ZnCN Au
(ii) Plattner’s chlorine process : The roasted ore
is moistened with water and placed in wooden vats
with false perforated bottoms. It is saturated with
current of chlorine, gold chloride thus formed is
leached with water and the solution is treated with a
reducing agent such as
4
FeSO or HS
2
to precipitate
gold.
3 4 3 2 43
AuCl 3 FeSO Au FeCl Fe ( SO )
2 AuCl 3 HS 6 HCl 3 S 2 Au
3 2
The impure gold thus obtained contains
impurities of Ag and Cu. The removal of Ag and Cu from
gold is called parting. This is done by heating impure
gold with conc.
2 4
HSO
(or )
3
HNO
when Ag and Cu
dissolve leaving behind Au.
Cu HSO CuSO SO HO
2 4 4 2 2
Ag HSO AgSO SO HO
2 4 2 4 2 2
Properties of Gold: Gold is a yellow, soft and
heavy metal. Gold and Ag are called noble metals since
they are not attacked by atmospheric oxygen. However,
Ag gets tarnished when exposed to air containing traces
of HS 2
. Gold is malleable, ductile and a good conductor
of heat and electricity.
Pure gold is soft. It is alloyed with Ag or Cu for
making jewellery. Purity of gold is expressed in terms
of carats. Pure gold is 24 carats. Gold ’14 carats’ means
that it is an alloy of gold which contains 14 parts by
weight of pure gold and 10 parts of copper per 24 parts
by weight of the alloy. Thus the percentage of gold in
’14 carats” of gold is = 14 58. 3 %
24
100
Most of the jewellery is made from 22 carat gold
(91.66% pure gold). Gold is quite inert. It does not
react with oxygen, water and acids but dissolves in
aqua regia
3 2 2 ] 3
3 2
HCl HNO NOCl HO Cl
3 ] 2
3
Au Cl AuCl
Auric chloride Nitrosylchloride
2 9 3 2 6 3
3 3 2
Au HCl HNO AuCl HO NOCl
Oxidation states of gold: The principal oxidation
states of gold are + 1 and + 3 though + 1 state is more
stable than + 3.
Compounds of gold
(1) Auric chloride, AuCl 3
: It is prepared by
passing dry 2
Cl over finely divided gold powder at 573
K
2 3
2 3 AuCl
K
Au Cl
It is a red coloured crystalline solid soluble in
water and decomposes on heating to give gold (I)
chloride and 2
Cl
3 2
Heat
AuCl AuCl Cl
It dissolves in conc. HCl forming chloroauric acid
[ ]
3 4
AuCl HCl HAuCl
Chloroauric acid is used in photography for toning
silver prints and as an antidote for snake poisoning.
(2) Aurous sulphide, Au 2
S : It is prepared when
HS
2
is passed through an acidified solution of
potassium aurocyanide, [ ( )]
2
KAuCN
2 K [ Au ( CN )] HS AuS 2 KCN 2 HCN
2 2 2
It is a dark brown solid, not attached by dilute
mineral acids and hence is probably the most stable
gold compound.
Zinc and its Compounds
(1) Occurrence of zinc: Zinc does not occur in the
native form since it is a reactive metal. The chief ores
of zinc are (i) Zinc blende ( ZnS ) (ii) Calamine or zinc
spar ( ZnCO 3
) and (iii) Zincite ( ZnO )
(2) Extraction of zinc : Zinc blende, after
concentration by Froth floatation process, is roasted in
air to convert it into ZnO. In case of calamine, ore is
calcined to get ZnO. The oxide thus obtained is mixed
with crushed coke and heated at 1673 K in fire clay
retorts (Belgian Process) when ZnO gets reduced to
metallic zinc. Being volatile at this temperature, the
metal distils over and is condensed leaving behind Cd,
Pb and Fe as impurities. The crude metal is called
spelter. The metal may be refined either by electrolysis
or by fractional distillation.
Properties of Zn : Zinc is more reactive than
mercury. It is a good conductor of heat and electricity.
Zinc readily combines with oxygen to form ZnO. Pure
zinc does not react with non-oxidising acids ( HCl or
2 4
HSO but the impure metal reacts forming
2
Zn ions
and evolving
2
H
gas.
2 2
Zn 2 HCl ZnCl H
Hot and conc.
2 4
HSO attacks zinc liberating
2
SO
gas
Zn HSO ZnSO SO HO
2 4 4 2 2
Zinc also reacts with both dilute (hot and cold)
3
HNO and conc.
3
HNO liberating nitrous oxide ( )
2
NO ,
ammonium nitrate ( )
4 3
NHNO and nitrogen dioxide
2
NO respectively.
3
4 Zn 10 HNO (warm, dilute)
Zn NO NO HO
3 2 2 2
3
4 Zn 10 HNO
(coldvery dilute)
Zn NO NHNO HO
3 2 4 3 2
3
Zn 4 HNO
(hot and
conc.) Zn NO NO HO
3 2 2 2
Zinc dissolves in hot concentrated NaOH forming
the soluble sod. Zincate
2 2 4 2
Zn 2 NaOH 2 HO Na [ Zn ( OH )] H
or
2 2 2
Zn 2 NaOH NaZnO H
2 2 (red)
673
2
Hg O HgO
K
or by heating mercuric nitrate alone or in the
presence of Hg
2 2
red
Heat
3 2
2 Hg ( NO ) 2 HgO 4 NO O
When NaOH is added to a solution of
2
HgCl
yellow precipitate of HgO are obtained.
Hg Cl 2 NaOH HgO HO 2 NaCl
2
(y ellow)
2 2
Red and yellow forms of HgO differ only in their
particle size. On heating to 673 K , yellow form changes
to red form.
red
673
y ellow
HgO HgO
K
It is used in oil paints or as a mild antiseptic in
ointments.
(2) Mercuric chloride, HgCl 2 : It is obtained by
treating Hg with 2
Cl or by heating a mixture of NaCl
and 4
HgSO in presence of small amount of
2
MnO
(which oxidises any Hg (I) salts formed during the
reaction).
2 2 4
Heat
4
2
HgSO 2 NaCl HgCl NaSO
MnO
It is a white crystalline solid and is commonly
known as corrosive sublimate. It is a covalent
compound since it dissolves in organic solvents like
ethanol and ether.
It is extremely poisonous and causes death. Its
best antidote is white of an egg.
When treated with stannous chloride, it is first
reduced to white ppt. of mercurous chloride and then to
mercury (black).
4
white ppt.
2 2 2 2
2 HgCl SnCl HgCl SnCl
4
grey
2 2 2
Hg Cl SnCl 2 Hg SnCl
With ammonia it gives a white ppt. known as
infusible white ppt.
HgCl NH HgNH Cl NHCl
2 3 2 4
A dilute solution of
2
HgCl is used as an antiseptic.
(3) Mercuric iodide, HgI 2 : It is obtained when a
required amount of KI solution is added to a solution of
2
HgCl.
HgCl 2 KI HgI 2 KCl
(red)
2 2
Below 400 K ,
2
HgI
is red but above 400 K , it
turns yellow
(red)
2
HgI
(y ellow)
2
HgI
2
HgI readily dissolves in excess of KI solution to
form the
2
4
( HgI ) complex ion.
soluble colourlesssolution
2 4
Red ppt.
2
HgI 2 KI KHgI
An alkaline solution of [ ]
2 4
K HgI is called Nessler’s
reagent and is used to test
4
NH ions.
It gives a brown ppt. of NH Hg O Hg I
2
(Iodide of Millon’s base) with
4
NH ions.
2 K [ HgI ] NH 3 KOH
2 4 3
NH HgOHgl KI HO
2 2
It is used in ointments for treating skin
infections.
(4) Mercurous chloride, Hg 2 Cl 2 : It is obtained as
under :
(a)
3
white ppt.
2 32 2 2
Hg ( NO ) 2 NaCl HgCl 2 NaNO
(b)
2 2
Heatinanironretort
2
HgCl Hg HgCl (condenses
on cooling)
It is purified by sublimation.
Mercurous chloride is also called calomel. It is a
white powder insoluble in HO
2
. On heating, it
decomposes to give
2
HgCl and Hg.
Hg Cl HgCl Hg
2
Heat
2 2
It dissolves in chlorine water forming mercuric
chloride.
2 2 2 2
Hg Cl Cl 2 HgCl
With ammonia, it turns black due to the formation
of a mixture of finely divided black Hg and mercuric
amino chloride.
2 2 3
HgCl 2 NH
Hg NHHgCl NHCl
4
(black)
2
It is used to prepare standard calomel electrode
and as a purgative in medicine.
(5) Mercuric sulphide, HgS : The solubility
product of HgS is lower than that of ZnS and hence it
gets precipitated as black solid when HS
2
is passed
through an acidic solution of any mercury (II) salt.
above 400
K
below 400
K
HgCl HS HgS 2 HCl
2 2
It is insoluble in water and HCl but dissolves in
aqua regia (1 part conc.
3
HNO 3 parts conc. HCl )
Nascentchlorine
2
Nitrosy lchloride
Aqua regia
3
3 HCl HNO NOCl 2 HO 2 [ Cl ]
HgS Cl HgCl S
(Soluble)
2
2 | |
On sublimation, its colour changes to red and
hence it is used as a red pigment.
(6) Mercuric sulphate, HgSO 4
: It is obtained
when HgS is treated with conc. 2 4
HSO.
Hg HSO HgSO SO HO
2 4 4 2 2
It is a white solid which decomposes on heating to
give mercurous sulphate.
2 4 2 2
675
4
3 HgSO HgSO Hg 2 SO 2 O
K
It is used as a catalyst in the hydration of alkynes
to give aldehydes or ketones. It is also used as a
cosmetic under the name Vermillon and in ayurvedic
medicine as makardhwaj.
(7) Amalgams : Mercury forms alloys commonly
known as amalgams, with all metals except iron and
platinum. Hence it is transported in iron containers.
(8) Alloy of transition metal : See in table
discuss earlier in metallurgy.
Lanthanides and Actinides
Lanthanides and actinides are collectively called
f - block elements because last electron in them enters
into f - orbitals of the antepenultimate ( i.e. , inner to
penultimate) shell partly but incompletely filled in
their elementary or ionic states. The name inner
transition, elements is also given to them because they
constitute transition series with in transition series ( d -
block elements) and the last electron enters into
antepenultimate shell ( n - 2) f. In addition to incomplete
d - subshell, their f - subshell is also incomplete. Thus,
these elements have three incomplete outer shells i.e.,
( n – 2), ( n – 1) and n shells and the general electronic
configuration of f - block elements is ( n –
1 14 010 2
f ( n 1 ) d ns
(1) Lanthanides : The elements with atomic
numbers 58 to 71 i.e. cerium to lutetium (which come
immediately after lanthanum Z = 57) are called
lanthanides or lanthanones or rare earths. These
elements involve the filling of 4 f - orbitals. Their
general electronic configuration is,
1 14 010 2
[ Xe ] 4 f 5 d 6 s
Promethium ( Pm ), atomic number 61 is the only
synthetic (man made) radioactive lanthanide.
Properties of lanthanides
(i) These are highly dense metals and possess
high melting points.
(ii) They form alloys easily with other metals
especially iron. e.g. misch metal consists of a rare
earth element (94–95%), iron (upto 5%) and traces of
S , C , Ca and Al , pyrophoric alloys contain Ce (40–5%),
La + neodymium (44%), Fe (4–5%), Al (0–5%) and the
rest is Ca , Si and C. It is used in the preparation of
ignition devices e.g., trace bullets and shells and flints
for lighters and cigarette.
(iii) Oxidation state : Most stable oxidation state
of lanthanides is +3. Oxidation states + 2 and + 4 also
exist but they revert to +3 e.g.
2 2 2
Sm , Eu , Yb lose
electron to become +3 and hence are good reducing
agents, where as Ce
4+
, Pr
4+
, Tb
4+
in aqueous solution
gain electron to become + 3 and hence are good
oxidizing agents. There is a large gap in energy of 4 f
and 5 d subshells and thus the number of oxidation
states is limited.
(iv) Colour : Most of the trivalent lanthanide ions
are coloured both in the solid state and in aqueous
solution. This is due to the partly filled f - orbitals which
permit f–f transition. The elements with xf electrons
have a similar colour to those of (14 – x ) electrons.
(v) Magnetic properties : All lanthanide ions
with the exception of Lu
3+
, Yb
3+
and Ce
4+
are
paramagnetic because they contain unpaired electrons
in the 4 f orbitals. These elements differ from the
transition elements in that their magnetic moments do
not obey the simple “ spin only ” formula n n 2
eff
B.M. where n is equal to the number of unpaired
electrons. In transition elements, the orbital
contribution of the electron towards magnetic moment
is usually quenched by interaction with electric fields
of the environment but in case of lanthanides the 4 f -
orbitals lie too deep in the atom for such quenching to
occur. Therefore, magnetic moments of lanthanides are
calculated by taking into consideration spin as well as
orbital contributions and a more complex formula
4 S S 1 L L 1
eff
B.M.
which involves the orbital quantum number L and
spin quantum number S.
(vi) Complex formation : Although the lanthanide
ions have a high charge (+3) yet the size of their ions is
very large yielding small charge to size ratio i.e. , low
charge density. As a consequence, they have poor
tendency to form complexes. They form complexes
