d and f block elements

INTRODUCTION
Definition: “The elements in which the last differentiating electron enters into the d-orbitals
of the penultimate shell i.e. (n–1) d where n is the last shell are called d-block elements”. A
transition element may also be defined as the element which partially filled d-orbital in their
ground state or most stable oxidation state.
Cu (Z = 29) = 1s2 2s2 2p6 3s2 3p6 3d10 4s1
Cu2+ = 1s2 2s2 2p6 3s2 3p6 3d9 (Cupric ion)
The properties of these elements are intermediate between the properties of s-block and p-
block elements. These elements represent a change or transition in properties from more
electropositive elements (s-block) to less electropositive elements (p-block). Therefore, these
elements are called transition elements.
Position in the periodic table
The position of d-block element has been shown in periodic table as follows:

(1) d-block elements lie in between ‘s’ and ‘p’ block elements, i.e. these elements are located
in the middle of the periodic table.
(2) d-block elements are present in,
Electronic configuration
(1) In d-block elements with increase in atomic number, the d orbitals of penultimate shell
i.e. (n-1) d are gradually filled by electrons. The general electronic configuration of d block
element is, (n-11) d1-10, ns0-2.
(2) Depending upon the d-orbitals of which penultimate shell i.e. n = 4, 5, 6, 7 are filled, four
rows (called series) os ten elements each obtained. They correspond to 3d, 4d, 5d and 6d
subshells.
(3) Energy of ‘(n–1)d’ subshell is slightly greater than ‘ns’ subshell, hence orbital is filled
first then (n – 1) d orbitals.
(4) The general electronic configuration of d-block/d-series elements be shown as follows:

Exceptional configuration of Cr and Cu. The exceptions observed in the first series are in
case of electronic configurations of chromium (Z = 24) and copper (Z = 29). It may be noted
that unlike other elements, chromium and copper have a single electron in the 4s-orbital. This
is due to the gain of additional stability by the atom by either having half-filled configuration
(i.e., containing 5 electrons in the d-sublevel) or completely filled configuration, (i.e.,
containing 10 electrons in the d-sublevel). The 3d-level in case of chromium gets exactly half-
filled with configuration 3d5 4s1 and that in case of copper, it gets completely filled with
configuration 3d10 4s1. This can be explained on the basis of exchange energy. Thus, the
electronic configuration of chromium (Z = 24) and copper (Z = 29) are 1s2 2s2 2p6 3s2 3p6 3d5
4s1 and 1s2 2s2 2p6 3s6 3p6 3d10 4s1 respectively.
Properties
1. Atomic Raddii
(i) The atomic radii of the transition metals lie in-between those of s- and p-block elements.
(ii) Generally, the atomic radii of d-block elements in a series decrease with increase in atomic
number but the decrease in atomic size is small after midway.
Explanation: The atomic radius decreases with the increase in atomic number as the nuclear
charge increases whereas the shielding effect of d-electron is small. After midway, as the
electrons enter the last but one (penultimate) shell, the added d-electron shields (screens) the
outermost electron. Hence, with the increase in the d-electrons, screening effect increases. This
counterbalances the increased nuclear charge due to increase in atomic number. As a result, the
atomic radii remain practically same after chromium. For example, in Fe, the two opposing
tendencies almost counterbalance and there is no change in the size from Mn to Fe.
(iii) At the end of the period, there is a slight increase in the atomic radii.
Explanation: Near the end of series, the increased electron-electron repulsions between added
electrons in the same orbitals are greater than the attractive forces due to the increased nuclear
charge. This results in the expansion of the electron cloud and thus the atomic radius increases.
(iv) The atomic radii increase down the group. This means that the atomic radii of second series
are larger than those of first transition series. But the atomic radii of the second and third
transition series are almost the same. The atomic radii of the elements of the second and third
transition metals are nearly same due to lanthanide contraction (or also called lanthanoid
contraction) discussed later.

Ionic radii
(i) The trend followed by the ionic radii is the same as that followed by atomic radii.
(ii) Ionic radii of transition metals are different in different oxidation states.
(iii) The ionic radii of the transition metals are smaller than those of the representative
elements belonging to the same period.

Metallic character
Except for mercury which is a liquid, all the transition elements have typical metallic structure.
They exhibit all the characteristics of metals. ductile, have high melting and boiling points,
high thermal and electrical conductivity and high tensile strength.
Transition elements have relatively low ionization energies and have one or two electrons in
their outermost energy level (ns1 or ns2). As a result, metallic bonds are formed. Hence, they
behave as metals. Greater the number of unpaired d electrons, stronger is the bonding due to
the overlapping of unpaired electrons between different metal atoms.
Melting Point
Transition metals have high melting points which is due to their strong metallic bond. The
metallic bond. The metallic bonding depends upon the number of unpaired e–. The melting
point first increases (Sc-Cr), reaches a maximum value (Cr) and then decreases (Fe-Zn)
* Tungsten (W) has the highest melting point.
* Mercury (Hg) has the lowest melting point.
* Mn has the lowest melting point in 3d series typical transition elements.

Ionisation energies or Ionisation enthalpies

(i). The first ionisation enthalpies of d-block elements lie between s-block and p-block
elements. They are higher than those of s block elements and are lesser than those of p-block
elements. The ionisation enthalpy gradually increases with increase in atomic number along a
given transition series though some irregularities are observed
Explanation:
(i) The increasing ionization enthalpies are due to increased nuclear charge with
increase in atomic number which reduces the size of the atom making the removal
of outer electron difficult.
(ii) In a given series, the difference in the ionisation enthalpies between any two
successive d-block elements is very much less than the difference in case of
successive s-block or p-block elements.
Explanation
(iii) The addition of d electrons in last but one [(n — 1) or penultimate] shell with
increase in atomic number provides a screening effect and thus shields the outer s
electrons from inward nuclear pull. Thus, the effects of increased nuclear charge
and addition of d electrons tend to oppose each other.
(iv) The first ionization enthalpy of Zn, Cd and Hg are, however, very high because of
the fully filled (n-1) d10 ns2 configuration.
(v) Although second and third ionization enthalpies also, in general, increase along a
period, but the magnitude of increase for the successive elements is much higher.
(vi) The high values of 3rd ionization enthalpies for Cu, Ni and Zn explain why they
show a maximum oxidation state of +2.
(vii) The first ionisation enthalpies of 5d elements are higher as compared to those of 3d
and 4d elements. This is because the weak shielding of nucleus by 4f electrons in
5d elements results in greater effective nuclear charge acting on the outer valence
electrons.
Oxidation state
“The net numerical charge assigned to an atom of an element in its combined state is known as
its Oxidation state or Oxidation number”.
1. With the exception of few elements, most of the d-block elements show more than one
oxidation state i.e. they show variable oxidation states. The elements show variable oxidation
state because of following reasons:
(i) ‘(n-1) d’ and ‘ns’ orbitals in the atoms of d-block elements have almost same energies and
therefore electron can be removed from ‘(n-1)d’ orbitals as easily as ‘s’ orbitals electrons.
(ii) After removing ‘s’ electrons, the remainder is called Kernel of the metal cations. In d-block
elements, the kernel is unstable and therefore it loses one or more electrons from (n – 1)d
electrons. This results in formation of cations with different oxidation states.
2. All transition elements show variable oxidation state except last element in each series.
3. Minimum oxidation state = Total number of electrons in 4s lost.
Maximum oxidation state = (Total number of electrons in 4s + number of unpaired electrons
in 3d lost).
In ‘3d’ series all element contains 2 electrons in ‘4s’ and hence they all give a common
minimum oxidation state of +2. (Except ‘Cr’ and ‘Cu’ where minimum oxidation state is +1.]
The maximum oxidation state is given by Mn i.e. Mn+7 in which two electrons are removed
from 4s and five unpaired electrons are removed from 3d orbitals.
4. The highest oxidation state is shown by Ruthenium (Ru) and Osmium (Os) i.e. +8.
5. Across the period oxidation state increases and it is maximum at the centre and then
decreases even if atomic number increases. The element which shows highest oxidation state
occur in the middle or near the middle of the series and then decreases.
6. Transition metals also show zero oxidation states in metal carbonyl complex. (Nickel
tetracarbonyl). Example: Ni in Ni (CO)4 has zero oxidation state.
7. The bonding in the compounds of lower oxidation state (+2, +3) is mostly ionic and the
bonding in the compounds of higher oxidation state is mostly covalent.
8. The relative stabilities of some oxidation states can be explained on the basis of rule extra
stability, according to which d0, d5 and d10 are stable configurations. For example, the stability
order of some ions is as follows

9. Cu+2 is more stable than Cu+1 even when Cu+1 is 3d10 while Cu+2 is 3d due to high heat of
hydration. Variable oxidation states shown by 3d-series of d-block elements.
Standard electrode potentials (Eo) and chemical reactivity
Thermodynamic stability of the compounds of transition elements can be evaluated in terms of
the magnitude of ionisation enthalpies of the metals — smaller the ionisation enthalpy of the
metal, stabler is its compound. In solution, the stability of the compounds depends upon
electrode potentials rather than ionization enthalpies. Electrode potential values depend upon
factors such as enthalpy of sublimation (or atomisation) of the metal, the ionisation enthalpy
and the hydration enthalpy, i.e.,
(ΔsubH is enthalpy of sublimation/ atomisation)

The total energy, ΔTH, for the process involving sublimation, ionisation and hydration
simultaneously, i.e., for the process, M(s) M+(aq)+e-, will be the sum of the three
types of enthalpies, i.e.,
ΔTH = ΔsubH + ΔiH + ΔhydH
Thus, ΔTH, is the total enthalpy change when solid meal, M is brought in the aqueous
medium in the form of monovalent ion, M+(aq).
Trends in the M2+/M Standard Electrode Potentials

(i) There is no regular trend in the Eo (M2+/M) values. This is because their ionization
enthalpies (I1 + IE2) and sublimation enthalpies do not show any regular trend.
(ii) The general trend towards less negative Eo values along the series is due to the general
increase in the sum of first and second ionization enthalpies.
(iii) Copper shows a unique behaviour in the series as it is the only metal having positive value
for Eo. This explains why is does not liberate H2 gas from acids. It reacts only with the
oxidizing acids (HNO3 and H2 SO4) which are reduced. The reason for positive Eo value for
copper is that the sum of enthalpies of sublimation and ionization is not balanced by
hydration enthalpy.
(iv) The values of Eo for Mn, Ni and Zn are more negative than expected from the general
trend. This is due to greater stability of half-filled d-subshell (d5) in Mn2+, and completely
filled d-subshell (d10) in Zn2+. The exceptional behaviour of Ni towards Eo value from the
regular trend is due to its high negative enthalpy of hydration.

Trends in the M3+/M2+ Standard Electrode Potentials

(i) A very low value for Eo (Sc3+/Sc2+) reflects the stability of Sc3+ ion which has a
noble gas configuration.
(ii) The highest value for Zn is on account of very high stability of Zn2+ ion with d10
configuration. It is difficult to remove an electron from it to change it into +3 state.
(iii) The comparatively high value of Eo (Mn3+/Mn2+) shows that Mn2+ is very stable
which is on account of stable d5 configuration of Mn2+.
(iv) The comparatively low value of Eo (Fe3+/Fe2+) is on account of extra stability of
Fe3+ (d5), i.e., low third ionization enthalpy of Fe.
(v) The comparatively low value for V is on account of the stability of V2+ ion due to
its half-filled t32g configuration
Chemical Reactivity and Eo Values.
The transition metals vary very widely in their chemical reactivity. Some of them are highly
electropositive and dissolve in mineral acids whereas a few of them are ‘noble’, i.e., they do
not react with simple acids. Some results of chemical reactivity of transition metals as related
to their Eo values are given below:
(i) The metals of the first transition series (except copper) are relatively more reactive
than the other series. Thus, they are oxidized by H+ ions though the actual rate is
slow, e.g., Ti and V are passive to dilute non-oxidizing acids at room temperature.
(ii) As already explained, less negative Eo values for M2+/M along the series indicate a
decreasing tendency to form divalent cations.
(iii) More negative Eo values than expected for Mn, Ni and Zn show greater stability for
Mn2+, Ni2+ and Zn2+.
(iv) Eo values for the redox couple M3+/M2+ indicate that Mn3+ and Co3+ ions are the
strongest oxidizing agents in aqueous solution whereas Ti2+, V2+ and Cr2+ are
strongest reducing agents and can liberate hydrogen from a dilute acid, e.g.,
2 Cr2+ (aq) + 2 H+ (aq) 2 Cr3+ (aq) + H2 (g)
Catalytic Property
Most transition elements and their compounds have good catalytic properties because
(a) They possess variable oxidation state
(b) They provide a large surface area for the reactant to be absorbed.

Magnetic Behaviour
When a substance is placed in a magnetic field of strength H, the intensity of the magnetic field
in the substance may be greater than or less than H. If the field in the substance is greater than
H, the substance is a paramagnetic material and attracts line of force. If the field in the
substance is less than H, the substance is diamagnetic. Diamagnetic materials tend to repel lines
of force. The magnetic moment of a substance depends on the strength of magnetic field
generated due to electronic spin, there is a change in electric flux which leads to induction of
magnetic field. The electron is also under orbital angular momentum or orbital spin. It leads to
generation of magnetic moment. In first transition elements series the orbital angular magnetic
moment is insignificant the orbital contribution is quenched by the electric fields of the
surrounding atoms so magnetic moment is equal to the spin magnetic moment only.

𝜇𝑒𝑓𝑓 = √𝑛(𝑛 + 2)𝐵𝑀 n - no. of unpaired electron.

Maximum transition elements and compounds are paramagnetic due to presence of unpaired
electrons.