The d-block of the periodic table contains the elements

of the groups 3-12 in which the d orbitals are

progressively filled in each of the four long periods.

The f-block consists of elements in which 4 f and 5 f

orbitals are progressively filled. They are placed in a

separate panel at the bottom of the periodic table. The

names transition metals and inner transition metals

are often used to refer to the elements of d-and

f-blocks respectively.

There are mainly four series of the transition metals,

3d series (Sc to Zn), 4d series (Y to Cd), 5d series (La

and Hf to Hg) and 6d series which has Ac and elements

from Rf to Cn. The two series of the inner transition

metals; 4f (Ce to Lu) and 5f (Th to Lr) are known as

lanthanoids and actinoids respectively.

Originally the name transition metals was derived

from the fact that their chemical properties were

transitional between those of s and p-block elements.

Now according to IUPAC, transition metals are defined

as metals which have incomplete d subshell either in

neutral atom or in their ions. Zinc, cadmium and

mercury of group 12 have full d

configuration in their

ground state as well as in their common oxidation states

and hence, are not regarded as transition metals.

However, being the end members of the 3d, 4d and 5d

transition series, respectively, their chemistry is studied

along with the chemistry of the transition metals.

The presence of partly filled d or f orbitals in their

atoms makes transition elements different from that of

The

- and

Block Elements

The

and

Block Elements

After studying this Unit, you will be

able to

• learn the positions of the d– and

f-block elements in the periodic

table;

• know the electronic configurations

of the transition (d-block) and the

inner transition (f-block) elements;

• appreciate the relative stability of

various oxidation states in terms

of electrode potential values;

• describe the preparation,

properties, structures and uses

of some important compounds

such as K

and KMnO

;

• understand the general

characteristics of the d– and

f–block elements and the general

horizontal and group trends in

them;

• describe the properties of the

f-block elements and give a

comparative account of the

lanthanoids and actinoids with

respect to their electronic

configurations, oxidation states

and chemical behaviour.

Objectives

Iron, copper, silver and gold are among the transition elements that

have played important roles in the development of human civilisation.

The inner transition elements such as Th, Pa and U are proving

excellent sources of nuclear energy in modern times.

UnitUnit

Unit

2020-21

216Chemistry

the non-transition elements. Hence, transition elements

and their compounds are studied separately. However,

the usual theory of valence as applicable to the non-

transition elements can be applied successfully to the

transition elements also.

Various precious metals such as silver, gold and

platinum and industrially important metals like iron,

copper and titanium belong to the transition metals series.

In this Unit, we shall first deal with the electronic

configuration, occurrence and general characteristics of

transition elements with special emphasis on the trends

in the properties of the first row (3d) transition metals

along with the preparation and properties of some

important compounds. This will be followed by

consideration of certain general aspects such as electronic

configurations, oxidation states and chemical reactivity

of the inner transition metals.

THE TRANSITION ELEMENTS (d-BLOCK)

The d–block occupies the large middle section of the periodic table

flanked between s– and p– blocks in the periodic table. The d–orbitals

of the penultimate energy level of atoms receive electrons giving rise to

four rows of the transition metals, i.e., 3d, 4d, 5d and 6d. All these

series of transition elements are shown in Table 8.1.

In general the electronic configuration of outer orbitals of these elements

is (n-1)d

1–10

1–2

. The (n–1) stands for the inner d orbitals which may

have one to ten electrons and the outermost ns orbital may have one

or two electrons. However, this generalisation has several exceptions

because of very little energy difference between (n-1)d and ns orbitals.

Furthermore, half and completely filled sets of orbitals are relatively

more stable. A consequence of this factor is reflected in the electronic

configurations of Cr and Cu in the 3d series. For example, consider the

case of Cr, which has 3d

configuration instead of 3d

; the

energy gap between the two sets (3d and 4s) of orbitals is small enough

to prevent electron entering the 3d orbitals. Similarly in case of Cu, the

configuration is 3d

and not 3d

. The ground state electronic

configurations of the outer orbitals of transition elements are given in

Table 8.1.

8.18.1

8.1

Position in thePosition in the

Position in the

Periodic TablePeriodic Table

Periodic Table

8.28.2

8.2

ElectronicElectronic

Electronic

ConfigurationsConfigurations

Configurations

of the d-Blockof the d-Block

of the d-Block

ElementsElements

Elements

Sc Ti V Cr Mn Fe Co Ni Cu Zn

Z 21 22 23 24 25 26 27

28 29 30

4s 2 2 2 1 2 2 2 2 1 2

3d 1 2 3 5 5 6 7 8 10 10

1st Series

Table 8.1: Electronic Configurations of outer orbitals of the Transition Elements

(ground state)

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217 The d- and f- Block Elements

The electronic configurations of outer orbitals of Zn, Cd, Hg and Cn

are represented by the general formula (n-1)d

. The orbitals in

these elements are completely filled in the ground state as well as in

their common oxidation states. Therefore, they are not regarded as

transition elements.

The d orbitals of the transition elements protrude to the periphery of

an atom more than the other orbitals (i.e., s and p), hence, they are more

influenced by the surroundings as well as affect the atoms or molecules

surrounding them. In some respects, ions of a given d

configuration

(n = 1 – 9) have similar magnetic and electronic properties. With partly

filled d orbitals these elements exhibit certain characteristic properties

such as display of a variety of oxidation states, formation of coloured

ions and entering into complex formation with a variety of ligands.

The transition metals and their compounds also exhibit catalytic

property and paramagnetic behaviour. All these characteristics have

been discussed in detail later in this Unit.

There are greater similarities in the properties of the transition

elements of a horizontal row in contrast to the non-transition elements.

However, some group similarities also exist. We shall first study the

general characteristics and their trends in the horizontal rows

(particularly 3d row) and then consider some group similarities.

2nd Series

Y Zr Nb Mo Tc Ru Rh Pd Ag Cd

Z 39 40 41 42 43 44 45 46 47

5s 2 2 1 1 1 1 1 0 1 2

4d 1 2 4 5 6 7 8 10 10 10

3rd Series

La Hf Ta W Re Os Ir Pt Au Hg

Z 57 72

6s 2 2 2 2 2 2 2 1 1 2

5d 1 2 3 4 5 6 7 9 10 10

Ac Rf Db Sg Bh Hs Mt Ds Rg Cn

Z 89 104 105 106 107 108 109 110 111

112

7s 2 2 2 2 2 2 2 2 1 2

6d 1 2 3 4 5 6 7 8 10 10

4th Series

On what ground can you say that scandium (Z = 21) is a transition

element but zinc (Z = 30) is not?

On the basis of incompletely filled 3d orbitals in case of scandium atom

in its ground state (3d

), it is regarded as a transition element. On the

other hand, zinc atom has completely filled d orbitals (3d

) in its

ground state as well as in its oxidised state, hence it is not regarded

as a transition element.

Example 8.1Example 8.1

Example 8.1

SolutionSolution

Solution

2020-21

218Chemistry

M.p./10

Atomic number

Intext QuestionIntext Question

Intext Question

8.1 Silver atom has completely filled d orbitals (4d

) in its ground state.

How can you say that it is a transition element?

We will discuss the properties of elements of first transition series

only in the following sections.

8.3.1 Physical Properties

Nearly all the transition elements display typical metallic properties

such as high tensile strength, ductility, malleability, high thermal and

electrical conductivity and metallic lustre. With the exceptions of Zn,

Cd, Hg and Mn, they have one or more typical metallic structures at

normal temperatures.

8.38.3

8.3

GeneralGeneral

General

Properties ofProperties of

Properties of

the Transitionthe Transition

the Transition

Elements

ElementsElements

Elements

(d-Block)(d-Block)

(d-Block)

(bcc = body centred cubic; hcp = hexagonal close packed;

ccp = cubic close packed; X = a typical metal structure).

Fig. 8.1: Trends in melting points of

transition elements

The transition metals (with the exception

of Zn, Cd and Hg) are very hard and have low

volatility. Their melting and boiling points are

high. Fig. 8.1 depicts the melting points of

transition metals belonging to 3d, 4d and 5d

series. The high melting points of these metals

are attributed to the involvement of greater

number of electrons from (n-1)d in addition to

the ns electrons in the interatomic metallic

bonding. In any row the melting points of these

metals rise to a maximum at d

except for

anomalous values of Mn and Tc and fall

regularly as the atomic number increases.

They have high enthalpies of atomisation which

are shown in Fig. 8.2. The maxima at about

the middle of each series indicate that one

unpaired electron per d orbital is particularly

Lattice Structures of Transition Metals

Sc Ti V Cr Mn Fe Co Ni Cu Zn

hcp hcp bcc bcc X bcc ccp ccp

ccp X

(bcc) (bcc) (bcc, ccp) (hcp) (hcp) (hcp)

Y Zr Nb Mo Tc Ru Rh Pd Ag Cd

hcp hcp bcc bcc hcp hcp ccp ccp ccp X

(bcc) (bcc) (hcp)

La Hf Ta W Re Os Ir Pt Au Hg

hcp

hcp bcc bcc hcp hcp ccp ccp ccp X

(ccp,bcc) (bcc)

2020-21

219 The d- and f- Block Elements

favourable for strong interatomic interaction. In general, greater the

number of valence electrons, stronger is the resultant bonding. Since

the enthalpy of atomisation is an important factor in determining the

standard electrode potential of a metal, metals with very high enthalpy

of atomisation (i.e., very high boiling point) tend to be noble in their

reactions (see later for electrode potentials).

Another generalisation that may be drawn from Fig. 8.2 is that the

metals of the second and third series have greater enthalpies of

atomisation than the corresponding elements of the first series; this is an

important factor in accounting for the occurrence of much more frequent

metal – metal bonding in compounds of the heavy transition metals.

Fig. 8.2

Trends in enthalpies

of atomisation of

transition elements

In general, ions of the same charge in a given series show progressive

decrease in radius with increasing atomic number. This is because the

new electron enters a d orbital each time the nuclear charge increases

by unity. It may be recalled that the shielding effect of a d electron is

not that effective, hence the net electrostatic attraction between the

nuclear charge and the outermost electron increases and the ionic

radius decreases. The same trend is observed in the atomic radii of a

given series. However, the variation within a series is quite small. An

interesting point emerges when atomic sizes of one series are compared

with those of the corresponding elements in the other series. The curves

in Fig. 8.3 show an increase from the first (3d) to the second (4d) series

of the elements but the radii of the third (5d) series are virtually the

same as those of the corresponding members of the second series. This

phenomenon is associated with the intervention of the 4f orbitals which

must be filled before the 5d series of elements begin. The filling of 4f

before 5d orbital results in a regular decrease in atomic radii called

Lanthanoid contraction which essentially compensates for the expected

8.3.2 Variation in

Atomic and

Ionic Sizes

Transition

Metals

/kJ

mol

–1

2020-21

220Chemistry

increase in atomic size with increasing atomic number. The net result

of the lanthanoid contraction is that the second and the third d series

exhibit similar radii (e.g., Zr 160 pm, Hf 159 pm) and have very similar

physical and chemical properties much more than that expected on

the basis of usual family relationship.

The factor responsible for the lanthanoid

contraction is somewhat similar to that observed

in an ordinary transition series and is attributed

to similar cause, i.e., the imperfect shielding of

one electron by another in the same set of orbitals.

However, the shielding of one 4f electron by

another is less than that of one d electron by

another, and as the nuclear charge increases

along the series, there is fairly regular decrease

in the size of the entire 4f

orbitals.

The decrease in metallic radius coupled with

increase in atomic mass results in a general

increase in the density of these elements. Thus,

from titanium (Z = 22) to copper (Z = 29) the

significant increase in the density may be noted

(Table 8.2).

Sc Ti V Cr Mn Fe Co Ni Cu Zn

Y Zr Nb Mo Tc Ru Rh

Pd Ag Cd

La Hf Ta W Re Os Ir Pt

Au Hg

Radius/nm

Fig. 8.3: Trends in atomic radii of

transition elements

Atomic number 21 22 23 24 25 26 27 28 29 30

Electronic configuration

M 3d

[Ar] 3d

–

Enthalpy of atomisation, ∆

/kJ mol

–1

326 473 515 397 281 416 425 430 339

126

Ionisation enthalpy/

∆

∆∆

∆

/kJ mol

–1

∆

I 631

656 650 653 717 762 758 736 745 906

∆

II 1235 1309 1414 1592 1509 1561 1644 1752 1958 1734

∆

III 2393 2657 2833 2990 3260 2962 3243 3402 3556 3837

Metallic/ionic M 164 147 135 129 137 126 125 125 128 137

radii/pm M

– – 79 82 82 77 74 70 73 75

73 67 64 62 65 65 61 60 – –

Standard

electrode M

/M – –1.63 –1.18 –0.90 –1.18 –0.44 –0.28 –0.25 +0.34 -0.76

potential E

/V M

– –0.37 –0.26 –0.41 +1.57 +0.77

+1.97 – – –

Density/g cm

–3

3.43 4.1 6.07 7.19 7.21 7.8 8.7 8.9 8.9 7.1

Element Sc Ti V Cr Mn Fe Co Ni Cu Zn

Table 8.2: Electronic Configurations and some other Properties of

the First Series of Transition Elements

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221 The d- and f- Block Elements

Why do the transition elements exhibit higher enthalpies of

atomisation?

Because of large number of unpaired electrons in their atoms they

have stronger interatomic interaction and hence stronger bonding

between atoms resulting in higher enthalpies of atomisation.

Example 8.2Example 8.2

Example 8.2

SolutionSolution

Solution

There is an increase in ionisation enthalpy along each series of the

transition elements from left to right due to an increase in nuclear

charge which accompanies the filling of the inner d orbitals. Table

8.2 gives the values of the first three ionisation enthalpies of the first

series of transition elements. These values show that the successive

enthalpies of these elements do not increase as steeply as in the case

of non-transition elements. The variation in ionisation enthalpy along

a series of transition elements is much less in comparison to the variation

along a period of non-transition elements. The first ionisation enthalpy,

in general, increases, but the magnitude of the increase in the second

and third ionisation enthalpies for the successive elements, is much

higher along a series.

The irregular trend in the first ionisation enthalpy of the metals of

3d series, though of little chemical significance, can be accounted for

by considering that the removal of one electron alters the relative energies

of 4s and 3d orbitals. You have learnt that when d-block elements form

ions, ns electrons are lost before (n – 1) d electrons. As we move along

the period in 3d series, we see that nuclear charge increases from

scandium to zinc but electrons are added to the orbital of inner subshell,

i.e., 3d orbitals. These 3d electrons shield the 4s electrons from the

increasing nuclear charge somewhat more effectively than the outer

shell electrons can shield one another. Therefore, the atomic radii

decrease less rapidly. Thus, ionization energies increase only slightly

along the 3d series. The doubly or more highly charged ions have d

configurations with no 4s electrons. A general trend of increasing values

of second ionisation enthalpy is expected as the effective nuclear charge

increases because one d electron does not shield another electron from

the influence of nuclear charge because d-orbitals differ in direction.

However, the trend of steady increase in second and third ionisation

enthalpy breaks for the formation of Mn

and Fe

respectively. In both

the cases, ions have d

configuration. Similar breaks occur at

corresponding elements in the later transition series.

The interpretation of variation in ionisation enthalpy for an electronic

configuration d

is as follows:

The three terms responsible for the value of ionisation enthalpy are

attraction of each electron towards nucleus, repulsion between the

8.3.3 Ionisation

Enthalpies

Intext QuestionIntext Question

Intext Question

8.2 In the series Sc (Z = 21) to Zn (Z = 30), the enthalpy of atomisation

of zinc is the lowest, i.e., 126 kJ mol

–1

. Why?

2020-21

222Chemistry

electrons and the exchange energy. Exchange energy is responsible for

the stabilisation of energy state. Exchange energy is approximately

proportional to the total number of possible pairs of parallel spins in

the degenerate orbitals. When several electrons occupy a set of

degenerate orbitals, the lowest energy state corresponds to the maximum

possible extent of single occupation of orbital and parallel spins (Hunds

rule). The loss of exchange energy increases the stability. As the stability

increases, the ionisation becomes more difficult. There is no loss of

exchange energy at d

configuration. Mn

has 3d

configuration and

configuration of Cr

is d

, therefore, ionisation enthalpy of Mn

is lower

than Cr

. In the same way, Fe

has d

configuration and Mn

has 3d

configuration. Hence, ionisation enthalpy of Fe

is lower than the Mn

In other words, we can say that the third ionisation enthalpy of Fe is

lower than that of Mn.

The lowest common oxidation state of these metals is +2. To

form the M

ions from the gaseous atoms, the sum of the first and

second ionisation enthalpy is required in addition to the enthalpy of

atomisation. The dominant term is the second ionisation enthalpy

which shows unusually high values for Cr and Cu where M

ions

have the d

and d

configurations respectively. The value for Zn is

correspondingly low as the ionisation causes the removal of 1s

electron which results in the formation of stable d

configuration.

The trend in the third ionisation enthalpies is not complicated by

the 4s orbital factor and shows the greater difficulty of removing an

electron from the d

(Mn

) and d

(Zn

) ions. In general, the third

ionisation enthalpies are quite high. Also the high values for third

ionisation enthalpies of copper, nickel and zinc indicate why it is

difficult to obtain oxidation state greater than two for these elements.

Although ionisation enthalpies give some guidance concerning the

relative stabilities of oxidation states, this problem is very complex and

not amenable to ready generalisation.

One of the notable features of a transition elements is the great variety

of oxidation states these may show in their compounds. Table 8.3 lists

the common oxidation states of the first row transition elements.

Sc Ti V Cr Mn Fe Co Ni Cu Zn

+2 +2 +2 +2 +2 +2 +2 +1 +2

+3 +3 +3 +3 +3 +3

+3 +3 +2

+4 +4 +4 +4 +4 +4 +4

+5 +5 +5

+6 +6 +6

Table 8.3: Oxidation States of the first row Transition Metals

(the most common ones are in bold types)

8.3.4 Oxidation

States

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223 The d- and f- Block Elements

The elements which give the greatest number of oxidation states

occur in or near the middle of the series. Manganese, for example,

exhibits all the oxidation states from +2 to +7. The lesser number of

oxidation states at the extreme ends stems from either too few electrons

to lose or share (Sc, Ti) or too many d electrons (hence fewer orbitals

available in which to share electrons with others) for higher valence

(Cu, Zn). Thus, early in the series scandium(II) is virtually unknown

and titanium (IV) is more stable than Ti(III) or Ti(II). At the other end,

the only oxidation state of zinc is +2 (no d electrons are involved). The

maximum oxidation states of reasonable stability correspond in value

to the sum of the s and d electrons upto manganese (Ti

, V

2–

, Mn

VII

–

) followed by a rather abrupt decrease in stability of

higher oxidation states, so that the typical species to follow are Fe

II,III

, Ni

, Cu

I,II

, Zn

The variability of oxidation states, a characteristic of transition

elements, arises out of incomplete filling of d orbitals in such a way

that their oxidation states differ from each other by unity, e.g., V

, V

III

, V

. This is in contrast with the variability of oxidation states of non

transition elements where oxidation states normally differ by a unit

of two.

An interesting feature in the variability of oxidation states of the d–

block elements is noticed among the groups (groups 4 through 10).

Although in the p–block the lower oxidation states are favoured by the

heavier members (due to inert pair effect), the opposite is true in the

groups of d-block. For example, in group 6, Mo(VI) and W(VI) are

found to be more stable than Cr(VI). Thus Cr(VI) in the form of dichromate

in acidic medium is a strong oxidising agent, whereas MoO

and WO

are not.

Low oxidation states are found when a complex compound has

ligands capable of π-acceptor character in addition to the σ-bonding.

For example, in Ni(CO)

and Fe(CO)

, the oxidation state of nickel and

iron is zero.

Name a transition element which does not exhibit variable

oxidation states.

Scandium (Z = 21) does not exhibit variable oxidation states.

Example 8.3Example 8.3

Example 8.3

SolutionSolution

Solution

Intext QuestionIntext Question

Intext Question

8.3 Which of the 3d series of the transition metals exhibits the

largest number of oxidation states and why?

2020-21

224Chemistry

Table 8.4 contains the thermochemical parameters related to the

transformation of the solid metal atoms to M

ions in solution and their

standard electrode potentials. The observed values of E

and those

calculated using the data of Table 8.4 are compared in Fig. 8.4.

The unique behaviour of Cu, having a positive E

, accounts for its

inability to liberate H

from acids. Only oxidising acids (nitric and hot

concentrated sulphuric) react with Cu, the acids being reduced. The

high energy to transform Cu(s) to Cu

(aq) is not balanced by its hydration

enthalpy. The general trend towards less negative E

values across the

8.3.5 Trends in the

Standard

Electrode

Potentials

Intext QuestionIntext Question

Intext Question

8.4 The E

/M) value for copper is positive (+0.34V). What is possible

reason for this? (Hint: consider its high ∆

and low ∆

hyd

)

Why is Cr

reducing and Mn

oxidising when both have d

configuration?

is reducing as its configuration changes from d

to d

, the latter

having a half-filled t

level (see Unit 9) . On the other hand, the change

from Mn

to Mn

results in the half-filled (d

) configuration which has

extra stability.

Example 8.4Example 8.4

Example 8.4

SolutionSolution

Solution

Fig. 8.4: Observed and calculated values for the standard

electrode potentials

→ M°) of the elements Ti to Zn

series is related to the general increase in the sum of the first and second

ionisation enthalpies. It is interesting to note that the value of E

for Mn,

Ni and Zn are more negative than expected from the trend.

2020-21

225 The d- and f- Block Elements

Element (M)

∆∆

∆

(M)

∆∆

∆

∆∆

∆

∆∆

∆

hyd

) E

Ti 469 656 1309 -1866 -1.63

V 515 650 1414 -1895 -1.18

Cr 398

653 1592 -1925 -0.90

Mn 279

717 1509 -1862 -1.18

Fe 418

762 1561 -1998 -0.44

Co 427

758 1644 -2079 -0.28

Ni 431

736 1752 -2121 -0.25

Cu 339

745 1958 -2121 0.34

Zn 130

906 1734 -2059 -0.76

Table 8.4: Thermochemical data (kJ mol

-1

) for the first row Transition

Elements and the Standard Electrode Potentials for the

Reduction of M

to M.

The stability of the half-filled d sub-shell in Mn

and the completely

filled d

configuration in Zn

are related to their E

values, whereas

for Ni is related to the highest negative ∆

hyd

An examination of the E

) values (Table 8.2) shows the varying

trends. The low value for Sc reflects the stability of Sc

which has a

noble gas configuration. The highest value for Zn is due to the removal

of an electron from the stable d

configuration of Zn

. The

comparatively high value for Mn shows that Mn

) is particularly

stable, whereas comparatively low value for Fe shows the extra stability

of Fe

). The comparatively low value for V is related to the stability

of V

(half-filled t

level, Unit 9).

Table 8.5 shows the stable halides of the 3d series of transition metals.

The highest oxidation numbers are achieved in TiX

(tetrahalides), VF

and CrF

. The +7 state for Mn is not represented in simple halides but

MnO

F is known, and beyond Mn no metal has a trihalide except FeX

and CoF

. The ability of fluorine to stabilise the highest oxidation state is

due to either higher lattice energy as in the case of CoF

, or higher bond

enthalpy terms for the higher covalent compounds, e.g., VF

and CrF

Although V

is represented only by VF

, the other halides, however,

undergo hydrolysis to give oxohalides, VOX

. Another feature of fluorides

is their instability in the low oxidation states e.g., VX

(X = CI, Br or I)

8.3.6 Trends in

the M

Standard

Electrode

Potentials

8.3.7 Trends in

Stability of

Higher

Oxidation

States

+ 6 CrF

+ 5 VF

CrF

+ 4 TiX

CrX

MnF

+ 3 TiX

CrX

MnF

FeX

CoF

+ 2 TiX

III

CrX

MnX

FeX

CoX

NiX

CuX

ZnX

+ 1 CuX

III

Oxidation

Number

Table 8.5: Formulas of Halides of 3d Metals

Key: X = F → I; X

= F → Br; X

= F, CI; X

III

= CI → I

2020-21

226Chemistry

and the same applies to CuX. On the other hand, all Cu

halides are

known except the iodide. In this case, Cu

oxidises I

–

to I

( )

2 2 2

2Cu 4I Cu I I

+ −

+ → +

However, many copper (I) compounds are unstable in aqueous

solution and undergo disproportionation.

2Cu

→ Cu

+ Cu

The stability of Cu

(aq) rather than Cu

(aq) is due to the much

more negative ∆

hyd

of Cu

(aq) than Cu

, which more than

compensates for the second ionisation enthalpy of Cu.

The ability of oxygen to stabilise the highest oxidation state is

demonstrated in the oxides. The highest oxidation number in the oxides

(Table 8.6) coincides with the group number and is attained in Sc

to Mn

. Beyond Group 7, no higher oxides of Fe above Fe

, are

known, although ferrates (VI)(FeO

)

2–

, are formed in alkaline media but

they readily decompose to Fe

and O

. Besides the oxides, oxocations

stabilise V

as VO

, V

as VO

and Ti

as TiO

. The ability of oxygen

to stabilise these high oxidation states exceeds that of fluorine. Thus

the highest Mn fluoride is MnF

whereas the highest oxide is Mn

The ability of oxygen to form multiple bonds to metals explains its

superiority. In the covalent oxide Mn

, each Mn is tetrahedrally

surrounded by O’s including a Mn–O–Mn bridge. The tetrahedral [MO

]

ions are known for V

, Cr

, Mn

and Mn

VII

+ 7 Mn

+ 6 CrO

+ 5 V

+ 4 TiO

CrO

MnO

+ 3 Sc

+ 2 TiO VO (CrO) MnO FeO CoO NiO CuO ZnO

+ 1 Cu

Table 8.6: Oxides of 3d Metals

* mixed oxides

Groups

3 4 5 6 7 8 9 10 11

Oxidation

Number

Intext QuestionIntext Question

Intext Question

8.5 How would you account for the irregular variation of ionisation

enthalpies (first and second) in the first series of the transition elements?

Example 8.5Example 8.5

Example 8.5

How would you account for the increasing oxidising power in the

series VO

< Cr

2–

< MnO

–

This is due to the increasing stability of the lower species to which they

are reduced.

SolutionSolution

Solution

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227 The d- and f- Block Elements

For the first row transition metals the E

values are:

V Cr Mn Fe Co Ni Cu

/M) –1.18 – 0.91 –1.18 – 0.44 – 0.28 – 0.25 +0.34

Explain the irregularity in the above values.

The E

/M) values are not regular which can be explained from

the irregular variation of ionisation enthalpies (

i i

∆ + ∆

1 2

H H

) and also

the sublimation enthalpies which are relatively much less for

manganese and vanadium.

Why is the E

value for the Mn

/Mn

couple much more positive

than that for Cr

/Cr

or Fe

/Fe

? Explain.

Much larger third ionisation energy of Mn (where the required change

is d

to d

) is mainly responsible for this. This also explains why the

+3 state of Mn is of little importance.

8.3.9 Magnetic

Properties

Transition metals vary widely in their chemical reactivity. Many of

them are sufficiently electropositive to dissolve in mineral acids, although

a few are ‘noble’—that is, they are unaffected by single acids.

The metals of the first series with the exception of copper are relatively

more reactive and are oxidised by 1M H

, though the actual rate at

which these metals react with oxidising agents like hydrogen ion (H

) is

sometimes slow. For example, titanium and vanadium, in practice, are

passive to dilute non oxidising acids at room temperature. The E

values

for M

/M (Table 8.2) indicate a decreasing tendency to form divalent

cations across the series. This general trend towards less negative E

values is related to the increase in the sum of the first and second

ionisation enthalpies. It is interesting to note that the E

values for Mn,

Ni and Zn are more negative than expected from the general trend.

Whereas the stabilities of half-filled d subshell (d

) in Mn

and completely

filled d subshell (d

) in zinc are related to their E

values; for nickel, E

value is related to the highest negative enthalpy of hydration.

An examination of the E

values for the redox couple M

(Table

8.2) shows that Mn

and Co

ions are the strongest oxidising agents

in aqueous solutions. The ions Ti

, V

and Cr

are strong reducing

agents and will liberate hydrogen from a dilute acid, e.g.,

2 Cr

(aq) + 2 H

(aq) → 2 Cr

(aq) + H

(g)

8.3.8 Chemical

Reactivity

and E

Values

Example 8.6Example 8.6

Example 8.6

Intext QuestionsIntext Questions

Intext Questions

8.6 Why is the highest oxidation state of a metal exhibited in its oxide or

fluoride only?

8.7 Which is a stronger reducing agent Cr

or Fe

and why ?

When a magnetic field is applied to substances, mainly two types of

magnetic behaviour are observed: diamagnetism and paramagnetism

(Unit 1). Diamagnetic substances are repelled by the applied field while

the paramagnetic substances are attracted. Substances which are

SolutionSolution

Solution

Example 8.7Example 8.7

Example 8.7

SolutionSolution

Solution

2020-21

228Chemistry

attracted very strongly are said to be ferromagnetic. In fact,

ferromagnetism is an extreme form of paramagnetism. Many of the

transition metal ions are paramagnetic.

Paramagnetism arises from the presence of unpaired electrons, each

such electron having a magnetic moment associated with its spin angular

momentum and orbital angular momentum. For the compounds of the

first series of transition metals, the contribution of the orbital angular

momentum is effectively quenched and hence is of no significance. For

these, the magnetic moment is determined by the number of unpaired

electrons and is calculated by using the ‘spin-only’ formula, i.e.,

( )

n 2

µ =

where n is the number of unpaired electrons and µ is the magnetic

moment in units of Bohr magneton (BM). A single unpaired electron

has a magnetic moment of 1.73 Bohr magnetons (BM).

The magnetic moment increases with the increasing number of

unpaired electrons. Thus, the observed magnetic moment gives a useful

indication about the number of unpaired electrons present in the atom,

molecule or ion. The magnetic moments calculated from the ‘spin-only’

formula and those derived experimentally for some ions of the first row

transition elements are given in Table 8.7. The experimental data are

mainly for hydrated ions in solution or in the solid state.

0 0 0

1 1.73 1.75

2 2.84 2.76

3 3.87 3.86

4 4.90 4.80

5 5.92 5.96

4 4.90 5.3 – 5.5

3 3.87 4.4 – 5.2

2 2.84 2.9 – 3, 4

1 1.73 1.8 – 2.2

0 0

Ion Configuration Unpaired

electron(s)

Magnetic moment

Calculated Observed

Table 8.7: Calculated and Observed Magnetic Moments (BM)

Calculate the magnetic moment of a divalent ion in aqueous solution

if its atomic number is 25.

With atomic number 25, the divalent ion in aqueous solution will have

configuration (five unpaired electrons). The magnetic moment, µ is

( )

5 5.92 BM

5 2

µ = =

Example 8.8Example 8.8

Example 8.8

SolutionSolution

Solution

2020-21

229 The d- and f- Block Elements

colourless

purple

blue

green

violet

blue

pink

yellow

green

bluepink

green

blue

colourless

Configuration Example Colour

Table 8.8: Colours of Some of the First Row (aquated)

Transition Metal Ions

8.3.11 Formation

of Complex

Compounds

Intext QuestionIntext Question

Intext Question

8.8 Calculate the ‘spin only’ magnetic moment of M

(aq)

ion (Z = 27).

When an electron from a lower energy d orbital is excited to a higher

energy d orbital, the energy of excitation corresponds to the frequency

of light absorbed (Unit 9). This frequency generally lies in the visible

region. The colour observed corresponds to the complementary colour

of the light absorbed. The

frequency of the light

absorbed is determined by

the nature of the ligand.

In aqueous solutions

where water molecules are

the ligands, the colours

of the ions observed are

listed in Table 8.8. A few

coloured solutions of

d–block elements are

illustrated in Fig. 8.5.

8.3.10 Formation

of Coloured

Ions

Fig. 8.5: Colours of some of the first row

transition metal ions in aqueous solutions. From

left to right: V

,Mn

,Fe

,Co

,Ni

and Cu

Complex compounds are those in which the metal ions bind a number

of anions or neutral molecules giving complex species with

characteristic properties. A few examples are: [Fe(CN)

]

3–

, [Fe(CN)

]

4–

[Cu(NH

)

]

and [PtCl

]

2–

. (The chemistry of complex compounds is

2020-21

230Chemistry

dealt with in detail in Unit 9). The transition metals form a large

number of complex compounds. This is due to the comparatively

smaller sizes of the metal ions, their high ionic charges and the

availability of d orbitals for bond formation.

The transition metals and their compounds are known for their catalytic

activity. This activity is ascribed to their ability to adopt multiple

oxidation states and to form complexes. Vanadium(V) oxide (in Contact

Process), finely divided iron (in Haber’s Process), and nickel (in Catalytic

Hydrogenation) are some of the examples. Catalysts at a solid surface

involve the formation of bonds between reactant molecules and atoms

of the surface of the catalyst (first row transition metals utilise 3d and

4s electrons for bonding). This has the effect of increasing the

concentration of the reactants at the catalyst surface and also weakening

of the bonds in the reacting molecules (the activation energy is lowering).

Also because the transition metal ions can change their oxidation states,

they become more effective as catalysts. For example, iron(III) catalyses

the reaction between iodide and persulphate ions.

2 I

–

+ S

2–

→ I

+ 2 SO

2–

An explanation of this catalytic action can be given as:

2 Fe

+ 2 I

–

→ 2 Fe

+ I

2 Fe

+ S

2–

→ 2 Fe

+ 2SO

2–

Interstitial compounds are those which are formed when small atoms

like H, C or N are trapped inside the crystal lattices of metals. They are

usually non stoichiometric and are neither typically ionic nor covalent,

for example, TiC, Mn

N, Fe

H, VH

and TiH

1.7

, etc. The formulas

quoted do not, of course, correspond to any normal oxidation state of

the metal. Because of the nature of their composition, these compounds

are referred to as interstitial compounds. The principal physical and

chemical characteristics of these compounds are as follows:

(i) They have high melting points, higher than those of pure metals.

(ii) They are very hard, some borides approach diamond in hardness.

(iii) They retain metallic conductivity.

(iv) They are chemically inert.

An alloy is a blend of metals prepared by mixing the components.

Alloys may be homogeneous solid solutions in which the atoms of one

metal are distributed randomly among the atoms of the other. Such

alloys are formed by atoms with metallic radii that are within about 15

percent of each other. Because of similar radii and other characteristics

of transition metals, alloys are readily formed by these metals. The

alloys so formed are hard and have often high melting points. The best

known are ferrous alloys: chromium, vanadium, tungsten, molybdenum

and manganese are used for the production of a variety of steels and

stainless steel. Alloys of transition metals with non transition metals

such as brass (copper-zinc) and bronze (copper-tin), are also of

considerable industrial importance.

8.3.12 Catalytic

Properties

8.3.13 Formation

Interstitial

Compounds

8.3.14 Alloy

Formation

2020-21

231 The d- and f- Block Elements

Intext QuestionIntext Question

Intext Question

8.9 Explain why Cu

ion is not stable in aqueous solutions?

8.4

Some

Important

Compounds ofCompounds of

Compounds of

Transition

Elements

ElementsElements

Elements

What is meant by ‘disproportionation’ of an oxidation state? Give an

example.

When a particular oxidation state becomes less stable relative to other

oxidation states, one lower, one higher, it is said to undergo disproportionation.

For example, manganese (VI) becomes unstable relative to manganese(VII) and

manganese (IV) in acidic solution.

3 Mn

2–

+ 4 H

→ 2 Mn

VII

–

+ Mn

+ 2H

Example 8.9Example 8.9

Example 8.9

SolutionSolution

Solution

8.4.1 Oxides and Oxoanions of Metals

These oxides are generally formed by the reaction of metals with

oxygen at high temperatures. All the metals except scandium form

MO oxides which are ionic. The highest oxidation number in the

oxides, coincides with the group number and is attained in Sc

. Beyond group 7, no higher oxides of iron above Fe

are

known. Besides the oxides, the oxocations stabilise V

as VO

, V

and Ti

as TiO

As the oxidation number of a metal increases, ionic character

decreases. In the case of Mn, Mn

is a covalent green oil. Even CrO

and V

have low melting points. In these higher oxides, the acidic

character is predominant.

Thus, Mn

gives HMnO

and CrO

gives H

CrO

and H

is, however, amphoteric though mainly acidic and it gives VO

3–

well as VO

salts. In vanadium there is gradual change from the basic

to less basic V

and to amphoteric V

. V

dissolves in acids

to give VO

salts. Similarly, V

reacts with alkalies as well as acids

to give

−

and

respectively. The well characterised CrO is basic

but Cr

is amphoteric.

Potassium dichromate K

Potassium dichromate is a very important chemical used in leather

industry and as an oxidant for preparation of many azo compounds.

Dichromates are generally prepared from chromate, which in turn are

obtained by the fusion of chromite ore (FeCr

) with sodium or

potassium carbonate in free access of air. The reaction with sodium

carbonate occurs as follows:

4 FeCr

+ 8 Na

+ 7 O

→ 8 Na

CrO

+ 2 Fe

+ 8 CO

The yellow solution of sodium chromate is filtered and acidified

with sulphuric acid to give a solution from which orange sodium

dichromate, Na

O can be crystallised.

2Na

CrO

+ 2 H

→ Na

+ 2 Na

+ H

2020-21

232Chemistry

Sodium dichromate is more soluble than potassium dichromate.

The latter is therefore, prepared by treating the solution of sodium

dichromate with potassium chloride.

+ 2 KCl → K

+ 2 NaCl

Orange crystals of potassium dichromate crystallise out. The

chromates and dichromates are interconvertible in aqueous solution

depending upon pH of the solution. The oxidation state of chromium

in chromate and dichromate is the same.

2 CrO

2–

+ 2H

→ Cr

2–

+ H

2–

+ 2 OH

→ 2 CrO

2–

+ H

The structures of

chromate ion, CrO

2–

and

the dichromate ion, Cr

2–

are shown below. The

chromate ion is tetrahedral

whereas the dichromate ion

consists of two tetrahedra

sharing one corner with

Cr–O–Cr bond angle of 126°.

Sodium and potassium dichromates are strong oxidising agents;

the sodium salt has a greater solubility in water and is extensively

used as an oxidising agent in organic chemistry. Potassium dichromate

is used as a primary standard in volumetric analysis. In acidic solution,

its oxidising action can be represented as follows:

2–

+ 14H

+ 6e

–

→ 2Cr

+ 7H

O (E

= 1.33V)

Thus, acidified potassium dichromate will oxidise iodides to iodine,

sulphides to sulphur, tin(II) to tin(IV) and iron(II) salts to iron(III). The

half-reactions are noted below:

6 I

–

→ 3I

+ 6 e

–

; 3 Sn

→ 3Sn

+ 6 e

–

3 H

S → 6H

+ 3S + 6e

–

; 6 Fe

→ 6Fe

+ 6 e

–

The full ionic equation may be obtained by adding the half-reaction for

potassium dichromate to the half-reaction for the reducing agent, for e.g.,

2–

+ 14 H

+ 6 Fe

→ 2 Cr

+ 6 Fe

+ 7 H

Potassium permanganate KMnO

Potassium permanganate is prepared by fusion of MnO

with an alkali

metal hydroxide and an oxidising agent like KNO

. This produces the

dark green K

MnO

which disproportionates in a neutral or acidic

solution to give permanganate.

2MnO

+ 4KOH + O

→ 2K

MnO

+ 2H

3MnO

2–

+ 4H

→ 2MnO

–

+ MnO

+ 2H

Commercially it is prepared by the alkaline oxidative fusion of MnO

followed by the electrolytic oxidation of manganate (Vl).

F d

used with KOH, oxidise

with air or KNO

2 4

MnO MnO ;

manganate ion

−

→

4 4

lectrolytic oxidation in

alkaline solution

MnO MnO

manganate permanganate ion

− −

→

2020-21

233 The d- and f- Block Elements

In the laboratory, a manganese (II) ion salt is oxidised by

peroxodisulphate to permanganate.

2Mn

+ 5S

2–

+ 8H

O → 2MnO

–

+ 10SO

2–

+ 16H

Potassium permanganate forms dark purple (almost black) crystals which

are isostructural with those of KClO

. The salt is not very soluble in water

(6.4 g/100 g of water at 293 K), but when heated it decomposes at 513 K.

2KMnO

→ K

MnO

+ MnO

+ O

It has two physical properties of considerable interest: its intense colour

and its diamagnetism along with temperature-dependent weak

paramagnetism. These can be explained by the use of molecular orbital

theory which is beyond the present scope.

The manganate and permanganate ions are tetrahedral; the π-

bonding takes place by overlap of p orbitals of oxygen with d orbitals

of manganese. The green manganate is paramagnetic because of one

unpaired electron but the permanganate is diamagnetic due to the

absence of unpaired electron.

Acidified permanganate solution oxidises oxalates to carbon dioxide,

iron(II) to iron(III), nitrites to nitrates and iodides to free iodine.

The half-reactions of reductants are:

COO

–

COO

–

10 CO

10e

–

5 Fe

→ 5 Fe

+ 5e

–

5NO

–

+ 5H

O → 5NO

–

+ 10H

+ l0e

–

10I

–

→ 5I

+ 10e

–

The full reaction can be written by adding the half-reaction for

KMnO

to the half-reaction of the reducing agent, balancing wherever

necessary.

If we represent the reduction of permanganate to manganate,

manganese dioxide and manganese(II) salt by half-reactions,

MnO

–

+ e

–

→ MnO

2–

= + 0.56 V)

MnO

–

+ 4H

+ 3e

–

→ MnO

+ 2H

O (E

= + 1.69 V)

MnO

–

+ 8H

+ 5e

–

→ Mn

+ 4H

O (E

= + 1.52 V)

We can very well see that the hydrogen ion concentration of the

solution plays an important part in influencing the reaction. Although

many reactions can be understood by consideration of redox potential,

kinetics of the reaction is also an important factor. Permanganate at

] = 1 should oxidise water but in practice the reaction is extremely slow

unless either manganese(ll) ions are present or the temperature is raised.

A few important oxidising reactions of KMnO

are given below:

1. In acid solutions:

(a) Iodine is liberated from potassium iodide :

10I

–

+ 2MnO

–

+ 16H

→ 2Mn

+ 8H

O + 5I

(b) Fe

ion (green) is converted to Fe

(yellow):

5Fe

+ MnO

–

+ 8H

→ Mn

+ 4H

O + 5Fe

2020-21

234Chemistry

2–

+ 2MnO

–

+ 16H

——> 2Mn

+ 8H

O + 10CO

(d) Hydrogen sulphide is oxidised, sulphur being precipitated:

S —> 2H

+ S

2–

+ 2MnO

–

+ 16H

——> 2Mn

+ 8H

O + 5S

(e) Sulphurous acid or sulphite is oxidised to a sulphate or

sulphuric acid:

5SO

2–

+ 2MnO

–

+ 6H

——> 2Mn

+ 3H

O + 5SO

2–

(f) Nitrite is oxidised to nitrate:

5NO

–

+ 2MnO

–

+ 6H

——> 2Mn

+ 5NO

–

+ 3H

2. In neutral or faintly alkaline solutions:

(a) A notable reaction is the oxidation of iodide to iodate:

2MnO

–

+ H

O + I

–

——> 2MnO

+ 2OH

–

+ IO

–

(b) Thiosulphate is oxidised almost quantitatively to sulphate:

8MnO

–

+ 3S

2–

+ H

O ——> 8MnO

+ 6SO

2–

+ 2OH

–

; the presence of zinc sulphate

or zinc oxide catalyses the oxidation:

2MnO

–

+ 3Mn

+ 2H

O ——> 5MnO

+ 4H

Note: Permanganate titrations in presence of hydrochloric acid are

unsatisfactory since hydrochloric acid is oxidised to chlorine.

UsesUses

Uses: Besides its use in analytical chemistry, potassium permanganate is

used as a favourite oxidant in preparative organic chemistry. Its uses for the

bleaching of wool, cotton, silk and other textile fibres and for the decolourisation

of oils are also dependent on its strong oxidising power.

THE INNER TRANSITION ELEMENTS (f-BLOCK)

The f-block consists of the two series, lanthanoids (the fourteen elements

following lanthanum) and actinoids (the fourteen elements following

actinium). Because lanthanum closely resembles the lanthanoids, it is

usually included in any discussion of the lanthanoids for which the

general symbol Ln is often used. Similarly, a discussion of the actinoids

includes actinium besides the fourteen elements constituting the series.

The lanthanoids resemble one another more closely than do the members

of ordinary transition elements in any series. They have only one stable

oxidation state and their chemistry provides an excellent opportunity to

examine the effect of small changes in size and nuclear charge along a

series of otherwise similar elements. The chemistry of the actinoids is, on

the other hand, much more complicated. The complication arises partly

owing to the occurrence of a wide range of oxidation states in these

elements and partly because their radioactivity creates special problems

in their study; the two series will be considered separately here.

The names, symbols, electronic configurations of atomic and some

ionic states and atomic and ionic radii of lanthanum and lanthanoids

(for which the general symbol Ln is used) are given in Table 8.9.

8.58.5

8.5

TheThe

The

LanthanoidsLanthanoids

Lanthanoids

2020-21

235 The d- and f- Block Elements

110

100

61 6559 63 67 69 71

Ionic radii/pm

Atomic number

8.5.1 Electronic

Configurations

8.5.2 Atomic and

Ionic Sizes

It may be noted that atoms of these elements have electronic

configuration with 6s

common but with variable occupancy of 4f level

(Table 8.9). However, the electronic configurations of all the tripositive

ions (the most stable oxidation state of all the lanthanoids) are of the

form 4f

(n = 1 to 14 with increasing atomic number).

The overall decrease in atomic and ionic radii from lanthanum to

lutetium (the lanthanoid contraction) is a unique feature in the

chemistry of the lanthanoids. It has far reaching

consequences in the chemistry of the third

transition series of the elements. The decrease

in atomic radii (derived from the structures of

metals) is not quite regular as it is regular in

ions (Fig. 8.6). This contraction is, of

course, similar to that observed in an ordinary

transition series and is attributed to the same

cause, the imperfect shielding of one electron

by another in the same sub-shell. However, the

shielding of one 4 f electron by another is less

than one d electron by another with the increase

in nuclear charge along the series. There is

fairly regular decrease in the sizes with

increasing atomic number.

The cumulative effect of the contraction of

the lanthanoid series, known as lanthanoid

contraction, causes the radii of the members

of the third transition series to be very similar

to those of the corresponding members of the

second series. The almost identical radii of Zr

(160 pm) and Hf (159 pm), a consequence of

the lanthanoid contraction, account for their

occurrence together in nature and for the

difficulty faced in their separation.

In the lanthanoids, La(II) and Ln(III) compounds are predominant

species. However, occasionally +2 and +4 ions in solution or in solid

compounds are also obtained. This irregularity (as in ionisation

enthalpies) arises mainly from the extra stability of empty, half-filled

or filled f subshell. Thus, the formation of Ce

is favoured by its

noble gas configuration, but it is a strong oxidant reverting to the

common +3 state. The E

value for Ce

/ Ce

is + 1.74 V which

suggests that it can oxidise water. However, the reaction rate is very

slow and hence Ce(IV) is a good analytical reagent. Pr, Nd, Tb and Dy

also exhibit +4 state but only in oxides, MO

. Eu

is formed by losing

the two s electrons and its f

configuration accounts for the formation

of this ion. However, Eu

is a strong reducing agent changing to the

common +3 state. Similarly Yb

which has f

configuration is a

reductant. Tb

has half-filled f-orbitals and is an oxidant. The

behaviour of samarium is very much like europium, exhibiting both

+2 and +3 oxidation states.

8.5.3 Oxidation

States

Fig. 8.6: Trends in ionic radii of lanthanoids

2020-21

236Chemistry

Electronic configurations* Radii/pm

Atomic Name Symbol Ln Ln

Ln Ln

Number

57 Lanthanum La 5d

187 106

58 Cerium Ce 4f

183

103

59 Praseodymium Pr 4f

182

101

60 Neodymium Nd 4f

181 99

61 Promethium Pm 4f

181 98

62 Samarium Sm 4f

180 96

63 Europium Eu 4f

199 95

64 Gadolinium Gd 4f

180 94

65 Terbium Tb 4f

178 92

66 Dysprosium Dy 4f

177 91

67 Holmium Ho 4f

176 89

68 Erbium Er 4f

175 88

69 Thulium Tm 4f

174 87

70 Ytterbium Yb 4f

173 86

71 Lutetium Lu 4f

–

Table 8.9: Electronic Configurations and Radii of Lanthanum and Lanthanoids

* Only electrons outside [Xe] core are indicated

All the lanthanoids are silvery white soft metals and tarnish rapidly in air.

The hardness increases with increasing atomic number, samarium being

steel hard. Their melting points range between 1000 to 1200 K but

samarium melts at 1623 K. They have typical metallic structure and are

good conductors of heat and electricity. Density and other properties

change smoothly except for Eu and Yb and occasionally for Sm and Tm.

Many trivalent lanthanoid ions are coloured both in the solid state

and in aqueous solutions. Colour of these ions may be attributed to

the presence of f electrons. Neither La

nor Lu

ion shows any colour

but the rest do so. However, absorption bands are narrow, probably

because of the excitation within f level. The lanthanoid ions other

than the f

type (La

and Ce

) and the f

type (Yb

and Lu

) are

all paramagnetic.

The first ionisation enthalpies of the lanthanoids are around

600 kJ mol

–1

, the second about 1200 kJ mol

–1

comparable with those

of calcium. A detailed discussion of the variation of the third ionisation

enthalpies indicates that the exchange enthalpy considerations (as in

3d orbitals of the first transition series), appear to impart a certain

degree of stability to empty, half-filled and completely filled orbitals

f level. This is indicated from the abnormally low value of the third

ionisation enthalpy of lanthanum, gadolinium and lutetium.

In their chemical behaviour, in general, the earlier members of the series

are quite reactive similar to calcium but, with increasing atomic number,

they behave more like aluminium. Values for E

for the half-reaction:

(aq) + 3e

–

→ Ln(s)

8.5.4 General

Characteristics

2020-21

237 The d- and f- Block Elements

LnC

with C

2773 K

NLn

heated with N

with H

Ln O

2 3

with acids

rns in O

hea

ted w

ith

with halogens

LnX

Ln(OH)

+ H

Ln S

8.68.6

8.6

The ActinoidsThe Actinoids

The Actinoids

are in the range of –2.2 to –2.4 V

except for Eu for which the value is

–2.0 V. This is, of course, a small

variation. The metals combine with

hydrogen when gently heated in the

gas. The carbides, Ln

C, Ln

and LnC

are formed when the metals are heated

with carbon. They liberate hydrogen

from dilute acids and burn in halogens

to form halides. They form oxides M

and hydroxides M(OH)

. The

hydroxides are definite compounds, not

just hydrated oxides. They are basic

like alkaline earth metal oxides and

hydroxides. Their general reactions are

depicted in Fig. 8.7.

The best single use of the

lanthanoids is for the production of alloy steels for plates and pipes. A

well known alloy is mischmetall which consists of a lanthanoid metal

(~ 95%) and iron (~ 5%) and traces of S, C, Ca and Al. A good deal of

mischmetall is used in Mg-based alloy to produce bullets, shell and

lighter flint. Mixed oxides of lanthanoids are employed as catalysts in

petroleum cracking. Some individual Ln oxides are used as phosphors

in television screens and similar fluorescing surfaces.

The actinoids include the fourteen elements from Th to Lr. The names,

symbols and some properties of these elements are given in Table 8.10.

Table 8.10: Some Properties of Actinium and Actinoids

Electronic conifigurations* Radii/pm

Atomic Name Symbol M M

Number

89 Actinium Ac 6d

111

90 Thorium Th 6d

91 Protactinium Pa 5f

92 Uranium U 5f

103 93

93 Neptunium Np 5f

101 92

94 Plutonium Pu 5f

100 90

95 Americium Am 5f

99 89

96 Curium Cm 5f

99 88

97 Berkelium Bk 5f

98 87

98 Californium Cf 5f

98 86

99 Einstenium Es 5f

– –

100 Fermium Fm 5f

– –

101 Mendelevium Md 5f

– –

102 Nobelium No 5f

– –

103 Lawrencium Lr 5f

–

Fig 8.7: Chemical reactions of the lanthanoids.

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238Chemistry

The actinoids are radioactive elements and the earlier members have

relatively long half-lives, the latter ones have half-life values ranging from

a day to 3 minutes for lawrencium (Z =103). The latter members could be

prepared only in nanogram quantities. These facts render their study

more difficult.

All the actinoids are believed to have the electronic configuration of 7s

and variable occupancy of the 5f and 6d subshells. The fourteen electrons

are formally added to 5f, though not in thorium (Z = 90) but from Pa

onwards the 5f orbitals are complete at element 103. The irregularities in

the electronic configurations of the actinoids, like those in the lanthanoids

are related to the stabilities of the f

, f

and f

occupancies of the 5f

orbitals. Thus, the configurations of Am and Cm are [Rn] 5f

and

[Rn] 5f

. Although the 5f orbitals resemble the 4f orbitals in their

angular part of the wave-function, they are not as buried as 4f orbitals

and hence 5f electrons can participate in bonding to a far greater extent.

The general trend in lanthanoids is observable in the actinoids as well.

There is a gradual decrease in the size of atoms or M

ions across the

series. This may be referred to as the actinoid contraction (like lanthanoid

contraction). The contraction is, however, greater from element to element

in this series resulting from poor shielding by 5f electrons.

There is a greater range of oxidation states, which is in part attributed to

the fact that the 5f, 6d and 7s levels are of comparable energies. The

known oxidation states of actinoids are listed in Table 8.11.

The actinoids show in general +3 oxidation state. The elements, in the

first half of the series frequently exhibit higher oxidation states. For example,

the maximum oxidation state increases from +4 in Th to +5, +6 and +7

respectively in Pa, U and Np but decreases in succeeding elements (Table

8.11). The actinoids resemble the lanthanoids in having more compounds

in +3 state than in the +4 state. However, +3 and +4 ions tend to hydrolyse.

Because the distribution of oxidation states among the actinoids is so

uneven and so different for the former and later elements, it is unsatisfactory

to review their chemistry in terms of oxidation states.

8.6.1 Electronic

Configurations

8.6.2 Ionic Sizes

8.6.3 Oxidation

States

The actinoid metals are all silvery in appearance but display

a variety of structures. The structural variability is obtained

due to irregularities in metallic radii which are far greater

than in lanthanoids.

8.6.4 General

Characteristics

and Comparison

with Lanthanoids

Ac Th Pa U Np Pu Am Cm Bk Cf Es Fm Md No Lr

3 3 3 3 3 3 3 3 3

3 3 3 3 3

4 4 4 4 4 4 4 4

5 5 5 5 5

6 6 6 6

7 7

Table 8.11: Oxidation States of Actinium and Actinoids

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239 The d- and f- Block Elements

The actinoids are highly reactive metals, especially when finely divided.

The action of boiling water on them, for example, gives a mixture of oxide

and hydride and combination with most non metals takes place at

moderate temperatures. Hydrochloric acid attacks all metals but most are

slightly affected by nitric acid owing to the formation of protective oxide

layers; alkalies have no action.

The magnetic properties of the actinoids are more complex than those

of the lanthanoids. Although the variation in the magnetic susceptibility

of the actinoids with the number of unpaired 5 f electrons is roughly

parallel to the corresponding results for the lanthanoids, the latter have

higher values.

It is evident from the behaviour of the actinoids that the ionisation

enthalpies of the early actinoids, though not accurately known, but are

lower than for the early lanthanoids. This is quite reasonable since it is to

be expected that when 5f orbitals are beginning to be occupied, they will

penetrate less into the inner core of electrons. The 5f electrons, will therefore,

be more effectively shielded from the nuclear charge than the 4f electrons

of the corresponding lanthanoids. Because the outer electrons are less

firmly held, they are available for bonding in the actinoids.

A comparison of the actinoids with the lanthanoids, with respect to

different characteristics as discussed above, reveals that behaviour similar

to that of the lanthanoids is not evident until the second half of the

actinoid series. However, even the early actinoids resemble the lanthanoids

in showing close similarities with each other and in gradual variation in

properties which do not entail change in oxidation state. The lanthanoid

and actinoid contractions, have extended effects on the sizes, and

therefore, the properties of the elements succeeding them in their

respective periods. The lanthanoid contraction is more important because

the chemistry of elements succeeding the actinoids are much less known

at the present time.

8.78.7

8.7

SomeSome

Some

ApplicationsApplications

Applications

of d- andof d- and

of d- and

f-Block

f-Blockf-Block

f-Block

ElementsElements

Elements

Iron and steels are the most important construction materials. Their

production is based on the reduction of iron oxides, the removal of

impurities and the addition of carbon and alloying metals such as Cr, Mn

and Ni. Some compounds are manufactured for special purposes such as

TiO for the pigment industry and MnO

for use in dry battery cells. The

battery industry also requires Zn and Ni/Cd. The elements of Group 11

are still worthy of being called the coinage metals, although Ag and Au

Name a member of the lanthanoid series which is well known

to exhibit +4 oxidation state.

Cerium (Z = 58)

Example 8.10Example 8.10

Example 8.10

SolutionSolution

Solution

Intext QuestionIntext Question

Intext Question

8.10 Actinoid contraction is greater from element to element than

lanthanoid contraction. Why?

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240Chemistry

are restricted to collection items and the contemporary UK ‘copper’ coins

are copper-coated steel. The ‘silver’ UK coins are a Cu/Ni alloy. Many of

the metals and/or their compounds are essential catalysts in the chemical

industry. V

catalyses the oxidation of SO

in the manufacture of

sulphuric acid. TiCl

with A1(CH

)

forms the basis of the Ziegler catalysts

used to manufacture polyethylene (polythene). Iron catalysts are used in

the Haber process for the production of ammonia from N

mixtures.

Nickel catalysts enable the hydrogenation of fats to proceed. In the Wacker

process the oxidation of ethyne to ethanal is catalysed by PdCl

. Nickel

complexes are useful in the polymerisation of alkynes and other organic

compounds such as benzene. The photographic industry relies on the

special light-sensitive properties of AgBr.

The d-block consisting of Groups 3-12 occupies the large middle section of the periodic

table. In these elements the inner d orbitals are progressively filled. The f-block is placed

outside at the bottom of the periodic table and in the elements of this block, 4f and

5f orbitals are progressively filled.

Corresponding to the filling of 3d, 4d and 5d orbitals, three series of transition

elements are well recognised. All the transition elements exhibit typical metallic properties

such as –high tensile strength, ductility, malleability, thermal and electrical conductivity

and metallic character. Their melting and boiling points are high which are attributed

to the involvement of (n–1)d electrons resulting into strong interatomic bonding. In

many of these properties, the maxima occur at about the middle of each series which

indicates that one unpaired electron per d orbital is particularly a favourable configuration

for strong interatomic interaction.

Successive ionisation enthalpies do not increase as steeply as in the main group

elements with increasing atomic number. Hence, the loss of variable number of electrons

from (n–1)d orbitals is not energetically unfavourable. The involvement of (n–1)d electrons

in the behaviour of transition elements impart certain distinct characteristics to these

elements. Thus, in addition to variable oxidation states, they exhibit paramagnetic

behaviour, catalytic properties and tendency for the formation of coloured ions, interstitial

compounds and complexes.

The transition elements vary widely in their chemical behaviour. Many of them are

sufficiently electropositive to dissolve in mineral acids, although a few are ‘noble’. Of the

first series, with the exception of copper, all the metals are relatively reactive.

The transition metals react with a number of non-metals like oxygen, nitrogen,

sulphur and halogens to form binary compounds. The first series transition metal oxides

are generally formed from the reaction of metals with oxygen at high temperatures. These

oxides dissolve in acids and bases to form oxometallic salts. Potassium dichromate and

potassium permanganate are common examples. Potassium dichromate is prepared from

the chromite ore by fusion with alkali in presence of air and acidifying the extract.

Pyrolusite ore (MnO

) is used for the preparation of potassium permanganate. Both the

dichromate and the permanganate ions are strong oxidising agents.

The two series of inner transition elements, lanthanoids and actinoids constitute

the f-block of the periodic table. With the successive filling of the inner orbitals, 4f, there

is a gradual decrease in the atomic and ionic sizes of these metals along the series

(lanthanoid contraction). This has far reaching consequences in the chemistry of the

elements succeeding them. Lanthanum and all the lanthanoids are rather soft white

metals. They react easily with water to give solutions giving +3 ions. The principal

oxidation state is +3, although +4 and +2 oxidation states are also exhibited by some

SummarySummary

Summary

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241 The d- and f- Block Elements

occasionally. The chemistry of the actinoids is more complex in view of their ability to

exist in different oxidation states. Furthermore, many of the actinoid elements are radioactive

which make the study of these elements rather difficult.

There are many useful applications of the d- and f-block elements and their

compounds, notable among them being in varieties of steels, catalysts, complexes,

organic syntheses, etc.

8.1 Write down the electronic configuration of:

(i) Cr

(iii) Cu

(v) Co

+ (vii) Mn

(ii) Pm

(iv) Ce

(vi) Lu

(viii) Th

8.2 Why are Mn

compounds more stable than Fe

towards oxidation to their

+3 state?

8.3 Explain briefly how +2 state becomes more and more stable in the first half

of the first row transition elements with increasing atomic number?

8.4 To what extent do the electronic configurations decide the stability of

oxidation states in the first series of the transition elements? Illustrate

your answer with examples.

8.5 What may be the stable oxidation state of the transition element with the

following d electron configurations in the ground state of their atoms : 3d

, 3d

and 3d

8.6 Name the oxometal anions of the first series of the transition metals in

which the metal exhibits the oxidation state equal to its group number.

8.7 What is lanthanoid contraction? What are the consequences of lanthanoid

contraction?

8.8 What are the characteristics of the transition elements and why are they

called transition elements? Which of the d-block elements may not be

regarded as the transition elements?

8.9 In what way is the electronic configuration of the transition elements different

from that of the non transition elements?

8.10 What are the different oxidation states exhibited by the lanthanoids?

8.11 Explain giving reasons:

(i) Transition metals and many of their compounds show paramagnetic

behaviour.

(ii) The enthalpies of atomisation of the transition metals are high.

(iii) The transition metals generally form coloured compounds.

(iv) Transition metals and their many compounds act as good catalyst.

8.12 What are interstitial compounds? Why are such compounds well known for

transition metals?

8.13 How is the variability in oxidation states of transition metals different from

that of the non transition metals? Illustrate with examples.

8.14 Describe the preparation of potassium dichromate from iron chromite ore.

What is the effect of increasing pH on a solution of potassium dichromate?

8.15 Describe the oxidising action of potassium dichromate and write the ionic

equations for its reaction with:

(i) iodide (ii) iron(II) solution and (iii) H

Exercises

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242Chemistry

8.16 Describe the preparation of potassium permanganate. How does the acidified

permanganate solution react with (i) iron(II) ions (ii) SO

and (iii) oxalic acid?

Write the ionic equations for the reactions.

8.17 For M

/M and M

systems the E

values for some metals are as follows:

/Cr -0.9V Cr

/Cr

-0.4 V

/Mn -1.2V Mn

/Mn

+1.5 V

/Fe -0.4V Fe

/Fe

+0.8 V

Use this data to comment upon:

(i) the stability of Fe

in acid solution as compared to that of Cr

or Mn

and

(ii) the ease with which iron can be oxidised as compared to a similar process

for either chromium or manganese metal.

8.18 Predict which of the following will be coloured in aqueous solution? Ti

, V

, Sc

, Mn

, Fe

and Co

. Give reasons for each.

8.19 Compare the stability of +2 oxidation state for the elements of the first

transition series.

8.20 Compare the chemistry of actinoids with that of the lanthanoids with special

reference to:

(i) electronic configuration (iii) oxidation state

(ii) atomic and ionic sizes and (iv) chemical reactivity.

8.21 How would you account for the following:

(i) Of the d

species, Cr

is strongly reducing while manganese(III)

is strongly oxidising.

(ii) Cobalt(II) is stable in aqueous solution but in the presence of

complexing reagents it is easily oxidised.

(iii) The d

configuration is very unstable in ions.

8.22 What is meant by ‘disproportionation’? Give two examples of disproportionation

reaction in aqueous solution.

8.23 Which metal in the first series of transition metals exhibits +1 oxidation

state most frequently and why?

8.24 Calculate the number of unpaired electrons in the following gaseous ions: Mn

, V

and Ti

. Which one of these is the most stable in aqueous solution?

8.25 Give examples and suggest reasons for the following features of the transition

metal chemistry:

(i) The lowest oxide of transition metal is basic, the highest is

amphoteric/acidic.

(ii) A transition metal exhibits highest oxidation state in oxides

and fluorides.

(iii) The highest oxidation state is exhibited in oxoanions of a metal.

8.26 Indicate the steps in the preparation of:

(i) K

from chromite ore. (ii) KMnO

from pyrolusite ore.

8.27 What are alloys? Name an important alloy which contains some of the

lanthanoid metals. Mention its uses.

8.28 What are inner transition elements? Decide which of the following atomic

numbers are the atomic numbers of the inner transition elements : 29, 59,

74, 95, 102, 104.

8.29 The chemistry of the actinoid elements is not so smooth as that of the

lanthanoids. Justify this statement by giving some examples from the

oxidation state of these elements.

8.30 Which is the last element in the series of the actinoids? Write the electronic

configuration of this element. Comment on the possible oxidation state of

this element.

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243 The d- and f- Block Elements

Answers to Some Intext Questions

8.1 Silver (Z = 47) can exhibit +2 oxidation state wherein it will have

incompletely filled d-orbitals (4d), hence a transition element.

8.2 In the formation of metallic bonds, no eletrons from 3d-orbitals are involved

in case of zinc, while in all other metals of the 3d series, electrons from

the d-orbitals are always involved in the formation of metallic bonds.

8.3 Manganese (Z = 25), as its atom has the maximum number of unpaired

electrons.

8.5 Irregular variation of ionisation enthalpies is mainly attributed to varying

degree of stability of different 3d-configurations (e.g., d

, d

are

exceptionally stable).

8.6 Because of small size and high electronegativity oxygen or fluorine can

oxidise the metal to its highest oxidation state.

8.7 Cr

is stronger reducing agent than Fe

Reason: d

→

occurs in case of Cr

to Cr

But d

→

occurs in case of Fe

to Fe

In a medium (like water) d

is more stable as compared to d

(see CFSE)

8.9 Cu

in aqueous solution underoes disproportionation, i.e.,

2Cu

(aq)

→

(aq) + Cu(s)

The E

value for this is favourable.

8.10 The 5f electrons are more effectively shielded from nuclear charge. In other

words the 5f electrons themselves provide poor shielding from element to

element in the series.

8.31 Use Hund’s rule to derive the electronic configuration of Ce

ion, and calculate

its magnetic moment on the basis of ‘spin-only’ formula.

8.32 Name the members of the lanthanoid series which exhibit +4 oxidation states

and those which exhibit +2 oxidation states. Try to correlate this type of

behaviour with the electronic configurations of these elements.

8.33 Compare the chemistry of the actinoids with that of lanthanoids with reference to:

(i) electronic configuration (ii) oxidation states and (iii) chemical reactivity.

8.34 Write the electronic configurations of the elements with the atomic numbers

61, 91, 101, and 109.

8.35 Compare the general characteristics of the first series of the transition metals

with those of the second and third series metals in the respective vertical

columns. Give special emphasis on the following points:

(i) electronic configurations (ii) oxidation states (iii) ionisation enthalpies

and (iv) atomic sizes.

8.36 Write down the number of 3d electrons in each of the following ions: Ti

, V

, Mn

, Fe

, Co

, Ni

and Cu

. Indicate how would you expect the five

3d orbitals to be occupied for these hydrated ions (octahedral).

8.37 Comment on the statement that elements of the first transition series possess

many properties different from those of heavier transition elements.

8.38 What can be inferred from the magnetic moment values of the following complex

species ?

Example Magnetic Moment (BM)

[Mn(CN)

) 2.2

[Fe(H

]

5.3

[MnCl

] 5.9

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