d and f block elements NOTES XII

d and f Block Elements
The d-blockof the periodictable containsthe elementsof the groups3-12 inwhichthe d orbitalsare
progressivelyfilledineachof the fourlongperiods.The elementsconstituting the f -block are those
in which the 4 f and 5 f orbitals are progressively filled in the latter two long periods.
There are mainlythree seriesof the transitionmetals,3d series (Sc to Zn), 4d series (Y to Cd) and 5d
series(Lato Hg, omittingCe toLu). The fourth6d serieswhichbeginswithAcisstill incomplete. The
two series of the inner transition metals, (4f and 5f) are known as lanthanoids and actinoids
respectively.
A transitionelementisdefinedasthe one whichhasincompletelyfilleddorbitals in its ground state
or in any one of its oxidation states.
Zinc, cadmium and mercury of group12 have full d10
configurationintheirgroundstate aswell as in
theircommonoxidationstatesandhence,are notregardedastransitionmetals.However,beingthe
end members of the three transition series, their chemistry is studied along with the chemistry of
the transition metals.
THE TRANSITION ELEMENTS (d-BLOCK)
Position in the Position in the Periodic Table Periodic Table
The d–block occupies the large middle section flanked by s– and p– blocks in the periodic table.
Electronic Configurations of the d-Block of the d-Block Elements
In general the electronic configuration of these elements is (n-1)d1–10
ns1–2
.
Half and completelyfilledsetsof orbitalsare relatively more stable. A consequence of this factor is
reflected in the electronic configurations of Cr and Cu in the 3d series. Consider the case of Cr, for
example, which has 3d5
4s1
instead of 3d4
4s2
.
PHYSICAL PROPERTIES
The transition metals (with the exception of Zn, Cd and Hg) are very much hard and have low
volatility.
Their melting and boiling points are high.
• High M.P and B.P - Due to strong metallic bond and the presence of half filled d- orbitals
• Involvement of greater number of electrons from (n-1)d in addition to the ns electrons in
the inter atomic metallic bonding.
• Because of stronger interatomic bonding, transition elements have high M.P and B.P
• In moving along the period from left to right, the M.P of these metals first INCREASES to
MAXIMUM and the DECREASES regularly towards the end of the period.

• Melting points of these metals
rise to a maximum at d5
except
for anomalous values of Mn and
Tc andfall regularly as the atomic
number increases.
TRENDS OF M.P OF 3- d , 4-d AND 5-d
TRANSITION METALS
• The strength of interatomic bonds
in transition elements is roughly related to the number of half filled d- orbitals
• In the beginning the no. of half filled d- orbitals increases till the middle of the period
causing increase in strength of interparticle bonds But thereafter the pairing of
electrons in d – orbitals occurs and the no. of half filled orbitals decreases , which also
cause decrease in M.P
Trends in enthalpies of atomization of transition elements
• 3d < 4d < 5d
• Greater the number of valence electrons, stronger the inter atomic attraction, hence
stronger bonding between atoms resulting in higher enthalpies of atomization.
• metals of the second and third series have greater enthalpies of atomization than the
corresponding elements of the first series
Atomic and ionic radii
 The Atomic/ionic radii first DECREASES till the middle, becomes almost constant
and then INCREASES towards the end of the period.

 New electron enters a d orbital each time the nuclear charge increases by unity,
But 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.
 However the increased nuclear charge is partly cancelled by the increased
screening effect of electrons in the d – orbitals of penultimate shell.
 When the increased nuclear charge and increased Screening effect balance each
other, the atomic radii becomes almost constant.
 Increase in atomic radii towards the end may be attributed to the electron –
electron repulsion.
 In fact the pairing of electrons in d – orbitals occurs after d5
configuration.
 The repulsive interaction between the paired electron causes Increase in
Atomic/ ionic radii
3d< 4d= 5d
• There is 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 4d
• This is due to the intervention of the 4f orbital which must be filled before
the 5d series of elements begin.
There is a steady decrease in atomic radii from La due to the poor shielding of inner
core electrons (4f) is known lanthanoid contraction.
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.

IONISATION ENTHALPY
 Due to an increase in nuclear charge there is an increase in ionisation enthalpy
along each series of the transition elements from left to right.
 Ionisation enthalpies give some guidance concerning the relative stabilities of
oxidation states.
 Although the first ionisation enthalpy, in general, increases, the magnitude of
the increase in the second and third ionisation enthalpies for the successive
elements, in general, is much higher.
 Mostly IE1<IE2 <IE3 in each group
 The increase in IE is primarily due to increase in nuclear charge. As the
transition elements involve the gradual filling of (n-1)d orbitals, the effect of
increase in nuclear charge is partly cancelled by the increase in screening effect.
Consequently, the increase in I.E along the periods of d – block elements is very
small.
 3d < 4d < 5d (in 5d series - ineffective shielding by 4f electrons)
Relation between I.E and Stability of a metal in a given
oxdn state
With the help of I.E, we can predict which of the two metals in a given oxdn state is
thermodynamically more stable. For example,
When a metal M (0) is converted into M (11), the energy required is equal to I1 + I2
Similarly M (IV) = I1 + I2+ I3 + I4
• Ni (0) Ni (II) I1 + I2 =2.49 x 103
kJ mol -1
• Pt (0) Pt (II) I1 + I2 =2.66 x 103
kJ mol -1
• Ni (0) Ni (IV)
I1 + I2+ I3 + I4 =11.299 x 103
kJ mol -1
• Pt (0) Pt (IV)
I1 + I2+ I3 + I4 =9.36 x 103
kJ mol -1
I1 + I2 for Ni (II) is less than I1 + I2 for Pt (II). So Ni (II) is more stable
Similarly Pt (IV) is more stable

OXIDATION STATES
• One of the notable features of a transition element is the great variety of
oxidation states it may show in its compounds
• Stability of a particular oxdn state depends up on nature of the element with
which the transition metals form the compound
• 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.
• Elements in the beginning of the series exhibit fewer oxidation state (have small
no. of electrons in which they lose or contribute for sharing).
• Elements at the end of the series shows fewer oxdn states because they have too
many electrons in d – orbitals. So they have few vacant d – orbitals which can be
involved in bonding.
• Lower oxdn state – Covalent character
• Higher oxdn state – ionic
• Higher oxdn states are more stable for heavier members.
Eg : in group VI, Mo (VI) and W (VI) are more stable than Cr (VI). So Cr (VI) act as
strong oxidizing agent.
• The highest oxdn state - +8 (Ruthenium and Osmium).
• 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)4
and Fe(CO)5, the oxidation state of nickel and iron is zero.

Stable halides of first transition elements
X = F to I, XII = F,
XI = F to Br , X III = Cl to I
The highest oxidation numbers are achieved in TiX4 (tetrahalides), VF5 and CrF6. The
+7 state for Mn is not represented in simple halides but MnO3F is known, and
beyond Mn, no metal has a trihalide except FeX3 and CoF3.
Although V(V)is represented only by VF5, the other halides, however, undergo
hydrolysis to give oxohalides, VOX3. Another feature of fluorides is their instability in
the low oxidation states e.g., VX2 (X = CI, Br or I).
• All Cu(II) halides are known except the iodide. In this case, Cu2+
oxidises I–
to I2
:
2Cu2+
+ 4I-
→ Cu2I2 (s) + I 2
• However, many copper (I) compounds are unstable in aqueous solution and
undergo disproportionation. 2Cu2+
→ Cu2+
+ Cu
• The stability of Cu2+
(aq) rather than Cu+
(aq) is due to the much more negative
ΔhydH0
of Cu2+
(aq) than Cu+
, which more than compensates for the second
ionisation enthalpy of Cu.
• Transition metals also exhibits the highest Oxdn state in their Oxides.
• The ability of Oxygen to stabilize higher oxidation states are much higher than
Fluorine..

• The highest Oxdn state with Fluorine by Mn is +4 in MnF4 while it is + 7 in
Mn2O7.
• Oxygen has the ability to form Multiple bonds with Metal atom.
The oxides of 3 – d transition elements are given below :
• The highest oxidation number in the oxides coincides with the group number and
is attained in Sc2O3 to Mn2O7.
• Beyond Group 7, no higher oxides of Fe above Fe2O3, are known, although
ferrates (VI) (FeO4)2–
, are formed in alkaline media but they readily decompose to
Fe2O3 and O2.
• Besides the oxides, oxocations stabilise V(v) as VO2
+
, V(IV) as VO2+
and Ti(IV) as
TiO2+.
STANDARD ELECTRODE POTENTIAL
• ELECTRODE POTENTIALS ARE THE MEASURE OF THE VALUE OF TOTAL
ENTHALPY CHANGE.
• Electrode Potentials value depends enthalpy of atomization ΔHa & hydration
ΔH hyd
• Lower the std E. P (Eo
red), the more stable is the oxdn state of the metal in
aqueous state.

• The E0
(M2+
/M) value for copper is positive (+0.34V) :high ΔHa and low ΔH
hyd). --- GREATER AMNTOF ENERGY REQUIRED TO TRANSFORM CuINTO
Cu2+
. Due to+ve Eo
, Cu does not liberate hydrogenfromacids.
• The general trend towards less negative Eo
values across the 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 Eo
for Mn, Ni and Zn are more negative
than expected from the trend.
• The stability of the half-filled d sub-shell in Mn2+
and the completely filled d10
configuration in Zn2+
are related to their Eo
values, whereas Eo
for Ni is related to
the highest negative ΔhydHo
.
• The low value for Sc reflects the stability of Sc3+
which has a noble gas
configuration. The highest value for Zn is due to the removal of an electron from

the stable d10
configuration of Zn2+.
The comparatively high value for Mn shows
that Mn2+
(d5
) is particularly stable, whereas comparatively low value for Fe shows
the extra stability of Fe3+
(d5
).
CHEMICAL REACTIVITY
• 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 simple 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.
• The EO
valuesfor M2+/
M indicate a decreasing tendency to form divalentcations
across the series.
• This general trend towards less negative EO
values is related to the increase in the
sum of the first and second ionisation enthalpies.
• It is interesting to note that the EO
values for Mn, Ni and Zn are more negative
than expected from the general trend.
• EO
values for the redox couple M3+/M2+ shows that Mn3+
and Co3+
ions are the
strongest oxidising agents in aqueous solutions. The ions Ti2+,
V2+
and Cr2+
are
strong reducing agents and will liberate hydrogen from a dilute acid,
e.g., 2 Cr2+
(aq) + 2 H+
(aq) → 2 Cr3+
(aq) + H2(g)
MAGNETIC PROPERTIES
• Substances which contain species (Atoms/ions/molecules) with unpared
electrons in their orbitals – PARAMAGNETIC.
• PARAMAGNETIC SUBSTANCES are weakly attracted by the magnetic field.
• Strongly attracted called FERROMAGNETIC.
• Substances which do not contain any unpaired electrons and are repelled my
magnetic field - DIAMAGNETIC.
• Transition metals usually contains unpaired electrons – so it is paramagnetic.

• Paramagnetic behavior increases with increase in unpaired electron.
• Paramagnetism expressed in terms of Magnetic moment., it is related to no. of
unpaired electrons.
• The magnetic moments calculated from the ‘spin-only’ formula and those
derived experimentally.
Magnetic moment µ = √ n(n+2) BM
n- no. of unpaired electrons
BM – Bohr magnetone (unit of M.M)
BM = 9.27x10-21
erg/gauss
• Single unpaired electron has a magnetic moment of 1.73 Bohr magnetons (BM).
Magnetic moment of an electron is due to its spin angular momentum and orbital
angular momentum.
Formation of Coloured Ions
• 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.
• 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.
• Zn 2+
/ Cd 2+
- all d orbitals are fully filled
• Ti 4+
- all d orbitals are vacant

so, no d – d transition occurs. Therefore they do not absorb radiations. So they are
colourless.
Formation of Complex Compounds
• Metal ions bind a number of anions or neutral molecules giving complex
[Fe(CN)6]3–
, [Fe(CN)6]4–
, [Cu(NH3)4]2+
and [PtCl4]2–
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.
Formation of Interstitial Compounds
• When small atoms like H, C or N are trapped inside the crystal lattices of metals
• They are usually non stoichiometric
Example, TiC, Mn4N, Fe3H, VH0.56 and TiH1.7
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.
Alloy Formation
• Because of similar radii and other characteristics of transition metals,
• The alloys so formed are hard and have often high melting points.
• 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),
CATALYTIC ACTIVITY
• 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
DISPROPORTIONATION
• 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 MnVI
O4 2–
+ 4 H+
→ 2 MnVII
O–
4 + MnIV
O2 + 2H2O
Oxides and Oxoanions of Metals
• The elements of first transition series form variety of oxides of different
oxidation states having general formula MO, M2O3, M3O6, MO2, MO3.
• Theses oxides are generally formed by heating the metal with oxygen at high
temperature.
Sc – Sc2O3 Basic
Ti – TiO Basic, Ti2O2 Basic, TiO2 Amphoteric
V – VO Basic, V2O3 Basic, VO2 Ampho, V2O5 Acidic
Cr – CrO Basic, Cr2O3 Ampho, CrO2 Ampho,
CrO3Acidic
Mn – MnO basic, Mn2O3 Basic, Mn3O4 Ampho,
MnO2 Ampho, Mn2O7 Acidic
Fe – FeO Basic, Fe2O3 Amph, Fe3O4 Basic
Co – CoO Basic
Ni – NiO Basic

Cu – Cu2O Basic, CuO Ampho
Zn – ZnO Ampho
In general
Lower oxidation state metal – BASIC
Higher oxidation state metal – ACIDIC
Intermediate oxidation state - AMPHOTERIC
Example
MnO (+2)basic, Mn2O3 (+3)Basic, Mn3O4 (+ 8/3)Ampho,
MnO2 (+4) Ampho, Mn2O7 (+7)Acidic
• The highest oxidation number in the oxides coincides with the group number and
is attained in Sc2O3 to Mn2O7.
• Beyond Group 7, no higher oxides of Fe above Fe2O3, are known, although
ferrates (VI) (FeO4)2–
, are formed in alkaline media but they readily decompose to
Fe2O3 and O2.
• Besides the oxides, oxocations stabilise V(v) as VO2
+
, V(IV) as VO2+
and Ti(IV) as
TiO2+.
• As the oxidation number of a metal increases, ionic character decreases. In the
case of Mn, Mn2O7 is a covalent green oil. Even CrO3 and V2O5 have low melting
points. In these higher oxides, the acidic character is predominant.
Potassium dichromate K2Cr2O7
STEP 1
Dichromates are generally prepared from chromate which in turn are obtained by
the fusion of chromite ore (FeCr2O4) with sodium or potassium carbonate in free
access of air. The reaction with sodium carbonate occurs as follows:
4 FeCr2O4 + 8 Na2CO3 + 7 O2 → 8 Na2CrO4 + 2Fe2O3 + 8 CO2

STEP 2
The yellow solution of sodium chromate is filtered and acidified with sulphuric acid
to give a solution from which orange sodium dichromate, Na2Cr2O7. 2H2O can be
crystallised.
2Na2CrO4 + H2SO4 → Na2Cr2O7 + Na2SO4 + H2O
STEP 3
Conversion of Sodium dichromate in to Potassium dichromate
Na2Cr2O7 + 2 KCl → K2Cr2O7 + 2 NaCl
The oxidation state of chromiumin chromate and dichromate is the
same.
2 CrO4
2–
+ 2H+
→ Cr2O7
2–
+ H2O
Cr2O7
2–
+ 2 OH-
→ 2 CrO4
2–
+ H2O
The chromate ion is tetrahedral whereas the dichromate ion consists of two
tetrahedral sharing one corner with Cr–O–Cr bond angle of 126°.
• Sodium and potassium dichromates are strong oxidising agents
• Potassium dichromate is used as a primary standard in volumetric analysis. In
acidic solution, its oxidising action can be represented as follows:
Cr2O7
2–
+ 14H+
+ 6e–
→ 2Cr3+
+ 7H2O (EV = 1.33V)
• 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–
→ 3I2 + 6 e–
;

 3 H2S → 6H+
+ 3S + 6e–
 3 Sn2+
→ 3Sn4+
+ 6 e–
 6 Fe2+
→ 6Fe3+
+ 6 e–
Cr2O7
2–
+ 14 H+
+ 6 Fe2+
→ 2 Cr3+
+ 6 Fe3+
+ 7 H2O
Potassium permanganate KMnO4
Potassium permanganate is prepared by fusion of MnO2 with an alkali metal
hydroxide and an oxidising agent like KNO3. This produces the dark green K2MnO4
which disproportionates in a neutral or acidic solution to give permanganate.
2MnO2 + 4KOH + O2 → 2K2MnO4 + 2H2O
3MnO4
2–
+ 4H+
→ 2MnO4
–
+ MnO2 + 2H2O
The manganate and permanganate ions are tetrahedral; the green manganate is
paramagnetic with one unpaired electron but the permanganate is diamagnetic.
THE INNER TRANSITION ELEMENTS ( f-BLOCK)
• The elements in which the additional electrons enters (n-2)f orbitals are called
inner transition elements. The valence shell electronic configuration of these
elements can be represented as (n – 2)f0-14
(n – 1)d0-1
ns2
.
• 4f inner transition metals are known as lanthanides because they come
immediately after lanthanum and 5f inner transition metals are known as
actinoids because they come immediately after actinium.

Atomic and ionic sizes: The Lanthanide Contraction
As the atomic number increases, each succeeding element contains one more
electron in the 4f orbital and one proton in the nucleus. The 4f electrons are
ineffective in screening the outer electrons from the nucleus causing imperfect
shielding. As a result, there is a gradual increase in the nucleus attraction for the
outer electrons. Consequently gradual decrease in size occur. This is called
lanthanide contraction.
Consequences of L. C
• There is close resemblance between 4d and 5d transition series.
• Ionization energy of 5d transition series is higher than 3d and 4d transition
series.
• Difficulty in separation of lanthanides
Ionization Enthalpies
• Fairly low I. E
• First ionization enthalpy is around 600 kJ mol-1
, the second about 1200 kJ mol-1
comparable with those of calcium.
• Due to low I. E, lanthanides have high electropositive character
Coloured ions

• Many of the lanthanoid ions are coloured in both solid and in solution due to
f – f transition since they have partially filled f – orbitals.
• Absorption bands are narrow, probably because of the excitation within f
level.
• La3+
and Lu3+
ions do not show any colour due to vacant and fully filled f-
orbitals.
Magnetic properties
• The lanthanoid ions other then the f 0
type (La3+
and Ce3+
) and the f14
type (Yb2+
and Lu3+
) are all paramagnetic. The paramagnetism rises to the maximum in
neodymium.
• Lanthanides have very high magnetic susceptibilities due to their large numbers
of unpaired f-electrons.
Oxidation States
• Predominantly +3 oxidation state.
• +3 oxidation state in La, Gd, Lu are especially stable ( Empty half filled and
Completely filled f – subshell respectively)
• Ce and Tb shows +4 oxdn state ( Ce 4+
- 4fo
& Tb 4+
4f7
)
• Occasionally +2 and +4 ions in solution or in solid compounds are also obtained.
• This irregularity arises mainly from the extra stability of empty, half filled or
filled f subshell.
• The most stable oxidation state of lanthanides is +3. Hence the ions in +2
oxidation state tend to change +3 state by loss of electron acting as reducing
agents whereas those in +4 oxidation state tend to change to +3 oxidation state
by gain of electron acting as a good oxidising agent in aqueous solution.
Question - Why Sm2+
, Eu2+
, and Yb2+
ions in solutions are good reducing agents but
an aqueous solution of Ce4+
is a good oxidizing agent?
Properties
• Silvery white soft metals, tarnish in air rapidly
• Hardness increases with increasing atomic number, samarium being steel hard.
• Good conductor of heat and electricity.
• Promethium - Radioactive

Chemical Properties
• Metal combines with hydrogen when gently heated in the gas.
• The carbides, Ln3C, Ln2C3 and LnC2 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 and hydroxides, M2O3 and M(OH)3, basic like alkaline earth
metal oxides and hydroxides.
The Actinides
• All isotopes are radioactive, with only 232
Th, 235
U, 238
U and 244
Pu having long half-
lives.
• Only Th and U occur naturally-both are more abundant in the earth’s crust than
tin.
• The others must be made by nuclear processes.
• The dominant oxidation state of actinides is +3. Actinides also exhibit an
oxidation state of +4. Some actinides such as uranium, neptunium and
plutonium also exhibit an oxidation state of +6.
• The actinides show actinide contraction (like lanthanide contraction) due to
poor shielding of the nuclear charge by 5f electrons.
• All the actinides are radioactive. Actinides are radioactive in nature.

Actinoide Contraction
• The size of atoms / M3+
ions decreases regularly along actinoid seris. The steady
decrease in ionic/ atomic radii with increase in atomic number is called
Actinoide Contraction.
• The contraction is greater from element to element in this series – due to poor
shielding effect by 5 f electron.
Magnetic properties
• Paramagnetic behaviour
• Magnetic properties are more complex than those of lanthanoids.
M.P and B.P
 High M.P and B.P
 Do not follow regular gradation of M.P or B.P with increase in atomic number
IONISATION ENTHALPY
• Low I.E. so electropositiity is High
COLOUR
• Generally coloured
• Colour depends up on the number of 5 f electrons
• The ions containing 5 f o
and 5 f 7
are colouress
Eg –
U 3+
(5 f 3
) – Red NP 3+
(5 f 4
) – Bluish

Assign reasons for the following:
(i) The enthalpies of atomisation of transition elements are high.
(ii) The transition metals and many of their compounds act as good catalysts.
(iii) From element to element the actinoid contraction is greater than the lanthanoid
contraction.
(iv) The E° value for the Mn 3+ / Mn 2+ couple is much more positive than that of Cr 3+ / Cr 2+.
(v) Scandium (Z = 21) does not exhibit variable oxidation states and yet it is regarded as a
transition element.
OR
a. What may be the possible oxidation states of the transition metals with the following d
electronic configurations in the ground state of their atoms: 3d 3 4s 2 , 3d5 4s 2 and 3d 6 4s
2 . Indicate relative stability of oxidation states in each case.
b. Write steps involved in the preparation of (i) Na 2CrO4 from chromite ore and (ii)
K2MnO4 from pyrolusite ore.
HINTS
(i) This is because transition metals have strong metallic bonds as they have a large number
of unpaired electrons.
(ii) The catalytic activity of transition metals is attributed to the following reasons:
(a) Because of their variable oxidation states transition metals form unstable
intermediate compounds and provide a new path with lower activation energy for the
reaction.
(b) In some cases, the transition metal provides a suitable large surface area with free
valencies on which reactants are adsorbed.
(iii) This is due to poorer shielding by 5f electrons in actinoids than that by 4f electron in the
lanthanoids.
(iv) This is because half filled d-subshell (3d5 ) in Mn 2+ is more stable.
(v) This is because scandium has partially filled d orbitals in the ground state (3d1 4s 2 ).
OR
(a)
Electronic Configuration Element Possible O.S. More stable O.S.
3d 2 4s 2 Vanadium + 2, + 3, + 4, + 5 + 5
3d5 4s 2 Manganese + 2, + 3, + 4, + 5, + 6, + 7 + 2, + 7
3d 6 4s 2 Iron + 2, + 3, + 4, + 6 + 2, + 3
(b)
(i) Chromite ore is fused with sodium carbonate in excess of air.
4FeCr2O4 + 8Na2 CO3 + 7O2 8Na2 CrO4 + 2Fe2O3 + 8CO2
Chromiteore Sod. Chromate
(ii) Pyrolusite ore (MnO2 ) is fused with KOH in the presence of O2 or oxidising agent such
as KNO3
2MnO2 + 4KOH + O2 2 K2 MnO4 + 2H2O
Pyrolusite ore Potassium maganate
QUESTION
Assign a reason for each of the following:
(i) The third ionization energy of Mn (Z = 25) is higher than that of either Cr (Z = 24) or Fe
(Z = 26).
(ii) Simple copper (I) salts are not stable in aqueous solutions.
HINTS
(I)This is because Mn 2+ is more stable as it has exactly half filled configuration 3d5 4s 0

(ii) Cu 2+(aq) is much more stable than Cu +(aq). This is because, although second ionization
enthalpy of copper is large but for Cu 2+(aq) is much more negative than that of Cu +(aq) and
therefore, it more than compensates for the second ionisation enthalpy of copper. Therefore,
Cu + ion aqueous solution undergoes disproportionation.
2Cu + (aq) Cu 2+ (aq) + Cu(s)
QUESYION
Discuss the relative stability in aqueous solutions of +2 oxidation state among the elements :
Cr, Mn, Fe and Co. How would you justify this situation? (At. Nos. Cr = 24, Mn = 25, Fe =
26, Co = 27)?
HINTS
On the basis of electrochemical series the standard electrode potential shows the following
order E° Mn 2+ /Mn < E° Cr 2+ /Cr < E° Fe2+ /Fe < E° Co2+ /Co
Therefore, Co2+ gets easily reduced to metallic cobalt while it is difficult to reduce Mn 2+.
Hence Mn 2+ will be most stable and the increasing stability order will be Co2+ < Fe2+ < Cr 2+
< Mn 2+
.

d and f block elements NOTES XII

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d and f block elements NOTES XII