D AND F BLOCK ELEMENTS

The electronic configurations of d-block elements have, in general, the following characteristics:
 an inner core of electrons with noble gas configuration i.e., ns2np6.
 (n – 1) d orbitals are filled progressively.
 Most of these have 2 electrons in the outermost i.e., ns-subshell.
Some of the elements (e.g., Cr, Cu, Nb, Mo, Tc, Ru, Rh, Ag, Pt, Au, Rg etc.) have only one electron in ns subshell
Pd has no electron in the ns-subshell. Pd (Z = 46): [Kr] 4d10 5s0
 In La (Z = 57) one electron goes to 5d-orbital before filling of 4f-orbital (an exception from Aufbau order).
 GENERAL CHARACTERISTICS OF D-BLOCK ELEMENTS
The general characteristics of d-block elements are:
1. Nearly all the transition elements have typical metallic properties such have high tensile strength, ductility,
malleability, high thermal and electrical conductivity and metallic lustre.
2. Except mercury which is liquid at room temperature, other transition elements have typical metallic
structures.
3. They have high melting and boiling points and have higher heats of vaporisation than non-transition
elements.
4. The transition elements have very high densities as compared to the metals of groups I and II (s-block).
5. The first ionisation energies of d-block elements are higher than those of s-block elements but are lesser
than those of p-block elements.
6. They are electropositive in nature.
7. Most of them form coloured compounds.
8. They have good tendency to form complexes.
9. They exhibit several oxidation states.
10. Their compounds are generally paramagnetic in nature.
11. They form alloys with other metals.
12. They form interstitial compounds with elements such as hydrogen, boron, carbon, nitrogen, etc.
13. Most of the transition metals such as Mn, Ni, Co, Cr, V, Pt, etc. and their compounds have been used as good
catalysts.
 1. The atomic radii of elements of a particular series
decrease with increase in atomic number but this
decrease in atomic radii becomes small after midway.
Explanation. The atomic radius decreases in a period
in the beginning, because with increase in atomic
number, the nuclear charge goes on increasing
2 . At the end of each period, there is a slight increase in
the atomic radii.
Explanation. Towards the end of the series, there are
increased electron-electron repulsions between the
added electrons in the same orbitals which exceed the
attractive forces due to increased nuclear charge.
3 . The ionic radii decrease with increase in oxidation state. For the same oxidation state, the ionic radii
generally decrease with increase in nuclear charge.

4 .Metallic Character and Enthalpy of Atomization All the transition elements are metals. They exhibit all the
characteristics of metals. They all have high density. except for Mn, Zn, Cd and Hg, all the transition metals have
one or more common metallic structures such as bcc (body centred cubic), hcp (hexagonal close packed) or ccp
(cubic close packed) structures at normal temperatures.
5. Density All these metals have high density. Within a period, the densities vary inversely with the atomic radii.
As we move in a period, the densities increase (as the radii decrease). Osmium and iridium have very high
density among these elements.

6. Melting points of these metals rise to a maximum value and then

decrease with increase in atomic number. However, manganese and
technetium metals have abnormally low melting points.
Explanation. The high melting and boiling points of these metals are due
to strong metallic bonds between the atoms of these elements.
Tungsten(W) has the highest melting point among the d-block elements.
7. Ionization enthalpy increases from left to right but quite slowly among d-block elements.
• zinc has high ionization enthalpy because electron has to be removed from 4s orbital of stable (3d10 4s2)
• Chromium and copper have exceptionally high ionization enthalpy values than those of their neighbours.
These exceptions are attributed to the extra stability of half filled and completely filled set of d-orbitals in
chromium (3d5) and copper (3d10) respectively.
• The value of second ionization enthalpy for zinc is correspondingly low because the ionisation involves the
removal of an electron resulting stable 3d10 configuration.
• The trend in third ionisation enthalpies is not complicated
by the 4s orbital factor and shows high values for Mn2+
and Zn2+ because of stable 3d5 and 3d10 electronic
configurations.
• Fe is very small because loss of third electron results in
stable 3d5 configuration
• The high values for copper, nickel and zinc indicate that
it is difficult to obtain oxidation state more than two (+2)
for these elements.
8. OXIDATION STATE

*Oxidation states within brackets are unstable.

** Most common oxidation states are in bold type.
• The highest oxidation states are found in compounds of fluorides and oxides because fluorine and oxygen
are most electronegative elements.
• The lowest oxidation states of Cr (3d54s1) and Cu (3d104s1) are +1 while for others, it is +2 (3d1–104s2).
• Except scandium, the most common oxidation state of the first row transition elements is +2 which arises
due to loss of two 4s-electrons.
• The elements which show the greatest number of oxidation states occur in or near the middle of the series.
For example, in the first transition series, manganese exhibits all the oxidation states from +2 to +7. T
• The small number of oxidation states at the extreme left hand side end (Sc, Ti) is due to lesser number of
electrons to lose or share.
• On the other hand, at the extreme right hand side end (Cu, Zn), it is due to large number of d-electrons
• Re exibits more number of oxidation state, OS +8 highest oxidation state
• The highest oxidation state is equal to the sum of the outer s(ns) and (n – 1) d-electrons. (e.g., TiIVO2
VVO+ 2’ CrVIO4 2−, MnVII O4 −, etc). For the remaining five elements, the minimum oxidation state is
given by the electrons in s-orbital while the maximum oxidation state is not related to their electronic
configurations, (e.g, Fe (II) and (III), Co (II) and (III), Ni (II), Cu (I) and (II), Zn(II). The highest oxidation
state shown by any transition metal is +8.
• In the +2 and +3 oxidation states, the bonds formed are mostly ionic.
• Transition metals also form compounds in low oxidation states such as +1 and 0 or negative. The
common examples are [Ni(CO)4], [Fe(CO)5] in which nickel and iron are in zero oxidation state.
8. Standard Electrode Potentials The magnitude of ionization enthalpy gives the amount of energy
required to remove electrons to form a particular oxidation state of the metal in a compound. Thus,
the value of ionisation enthalpies gives information regarding the thermodynamic stability of the
transition metal compounds in different oxidation states.
• Smaller the ionisation enthalpy of the metal, the stabler is its compound.
• nickel (II) compounds are thermodynamically more stable than platinum (II) compounds.
• platinum (IV) compounds are relatively more stable than nickel (IV) compounds.
• The smaller the value of total energy change for a particular oxidation state in aqueous solution,
greater will be the stability of that oxidation state.
• The lower the electrode potential i.e., more negative the standard reduction potential of the
electrode, the more stable is the oxidation state of the transition metal in the aqueous solution.
• Zinc has low enthalpy of atomisation and fairly large hydration energy. But it has also low
electrode potential (– 0.76 V) because of its very high ionisation enthalpy (IE1 + IE2).
• The transition elements react with halogens at high temperatures to form transition
metal halides. These reactions have very high activation energies, therefore, higher
temperatures are required to start the reaction. But once the reaction starts, the heat
of reaction is sufficient to continue the reaction. The halogens react in the order:
Order of reactivity: F2 > Cl2 > Br2 > I2 The stable halides of 3d series of transition
metals
• If E° value is more positive (or less negative), it means M3+ ions can be readily
reduced to M2+ ions.
• M3+ ions are less stable than M2+ ions in aqueous solution.
• M2+ ions are stable and cannot readily lose electrons to form M3+ ions.
• If E° value is less positive (or more negative), it means M3+ ions cannot be readily
reduced to M2+ ions.
• M3+ ions are more stable than M2+ ions.
• M2+ ions can readily lose electrons to form M3+ ions in solution.
• copper has positive reduction potential, E (0.34 V) and this shows that copper is least
reactive metal out of the first transition series.
(i) The low value of scandium reflects the stability of Sc3+ which has a noble gas configuration.
(ii) The comparatively high value for Mn shows that Mn2+ (d5 configuration) is particularly stable.
(iii) On the other hand, comparatively low value for Fe shows the extra stability of Fe3+ (d5
configuration).
(iv) The comparatively low value of V is related to the stability of V2+ (due to half filled t2g 3 energy level
of 3d orbitals in octahedral crystal field splitting
(i) In general, the elements of first transition series tend to exist in low oxidation states.
Chromium to zinc form stable difluorides and the other chlorides are also known.
(ii) Since fluorine is the most electronegative element, the transition metals show highest
oxidation states with fluorine. For example, CrF6 and VF5.
(iii) The highest oxidation states are found in TiX4 (tetrahalides, X = F, Cl, Br and I), VF5 and
CrF6.
(iv) The +7 oxidation state for Mn is not shown by simple halides. However, MnO3F is known
in which the oxidation state of Mn is +7.
(v) After Mn, the tendency to show higher oxidation states with halogens are uncommon.
Iron and cobalt form trihalides FeX3 (X = F, Cl or Br) and CoF3.
(vi) The tendency of fluorine to stablise the highest oxidation state is due to either higher
lattice enthalpy as in case of CoF3 or higher bond enthalpy due to higher covalent bonds
e.g., VF5 and CrF6.
(vii) V(V) is shown by VF5 only. However, the other halides undergo hydrolysis to form
oxohalides, VOX3.
(viii)Fluorides are relatively unstable in their low oxidation states. For example, vanadium
forms only VX2 (X = Cl, Br or I) and copper can form CuX (X = Cl, I).
(IX) Catalytic properties: Many of the transition metals and their compounds, particularly
oxides act as catalysts for a number of chemical reactions. Iron, cobalt, nickel, platinum,
chromium, manganese and their compounds are commonly used catalysts. All transitional
metals show multiple oxidation states and have large surface area so all metals work as a
catalyst.
(x) Magnetic properties: On the basis of the behaviour of substances in magnetic field, they
are of two types: (i) Diamagnetic, (ii) Paramagnetic. Diamagnetic substances have paired
electrons only. e.g., Zn has only paired electrons. In paramagnetic substances, it is
necessary to have at least one unpaired electron. Paramagnetism increases with the
increase in number of unpaired electrons. B.M. = Bohr magneton (unit of magnetic
moment) Diamagnetic and paramagnetic substances are repelled and attracted in the
magnetic field respectively (Magnetic properties of transition elements).
• Formation of Coloured Ions Most of the compounds of transition metals are coloured in
the solid form or solution form. This is in contrast to the compounds of s- and p-block
elements which are usually white.
 FeCr2O4

NaOH OR Na2CO3 O2

Na2CrO4+ Fe2O3 + H2O

H2SO4

Na2Cr2O7

KCL

K2Cr2O7
+ 3Sn4+
POTTASSIUM PERMANGANATE – KMNO4