Physicochemical Properties : Metallic Character

All the transition elements are metals, since the number of electrons in the outermost shell is very small, being equal to 2. They are hard, malleable and ductile. They exhibit all the three types of structures: face centred cubic (fcc), hexagonal close packed (hcp) and body centred cubic (bcc).

Metals of VIII and IB Groups are softer and more ductile than other metals. It appears that covalent and metallic bonding both exist in the atoms of transition metals. The presence of unfilled d-orbitals favours covalent bonding. These metals are good conductors of heat and electricity.

Melting and Boiling Points

The transition elements have very high melting and boiling points as compared to those of s and p block elements. Zn , Cd and Hg have relatively low values. The reason for these low values is that these metals have completely filled d-orbitals with no unpaired electron that may be available for covalent bonding amongst the atoms of these metals. The formation of covalent bonding occurs in the rest of the d-block elements on account of the presence of unfilled d-orbitals.

Although melting and boiling points show no definite trends in the three transition series, the metals having the highest melting and boiling points are towards the middle of each transition series.

Atomic (Covalent) and Ionic Radii

It will be seen that atomic and ionic radii values decrease generally, on moving from left to right in the period. This is due to the fact that an increase in the nuclear charge tends to attract the electron cloud inwards. The atomic radii for the elements from Cr to Cu are, however, very close to one another.

This is due to the fact that simultaneous addition of electron to 3d-level exercise the reverse effect by screening the outer 4s-electrons from the inward pull of the nucleus. As a result of these two opposing effects, the atomic radii do not alter much on moving from Cr to Cu.

The ionic radii of M2+ and M3+ ions follow the same trends as their atomic radii. The radii of M2+ ions, although somewhat smaller than that of Ca2+ ion (0.99 A° ) are comparable with it. Thus MO oxides of the transition element should be similar to CaO in many ways, although somewhat less basic and less soluble in water.

Similarly the hydration energies of M2+ ions should be similar to but somewhat greater than that of Ca2+ ion. This is borne out by facts, since the hydration energy of Ca2+ ion is 395 kcal and the observed values of hydration energies for the elements Ti2+ … Cu2+ are between 446 kcal and 597 kcal.

The radii of M3+ ions are slightly greater than that of Ga3+ ion (0.62 A°). Thus M2O3 oxides of transition elements should be similar to but slightly less acidic (more basic) than Ga2O3 and the hydration energies of M3+ ions should be less than 1124 kcal which is the hydration energy of Ga3+ ion. The observed values of hydration energies for the series Sc3+ … Fe3+ are between 947 kcal and 1072 kcal.

Ionisation Potentials

The first ionisation potentials of transitional elements lie between the values of those of s- and p-block elements. The first ionisation potentials of all the transition elements lie between 5 to 10 electron volts. In case of transition elements the addition of the extra electron in the (n-1) d level provides a screening effect which shields the outer ns electrons from the inward pull of positive nucleus on the outer ns electrons.
Thus the effects of the increasing nuclear charge and the shielding effect created due to the expansion of (n-1)d orbital oppose each other. On account of these counter effects, the ionisation potentials increase rather slowly on moving in a period of the first transition series.

First ionisation potentials

It is evident that the values for the first four 3d block elements (Sc, Ti, V and Cr) differ only slightly from one another. Similarly the values for Fe, Co, Ni and Cu also are fairly close to one another. The value of II for Zn is considerably higher. This is due to the extra-stability of 3d10 level which is completely filled in Zn atom.

Second ionisation potentials

The second ionisation potentials are seen to increase more or less regularly with the increase of atomic number. The value of III for Cr and Cu are higher than those of their neighbours. This is due to the fact that the electronic configurations of Cr+ and Cu+ ions have extra stable 3d5 and 3d10 levels.

There is a sudden fall in the values of ionisation potentials in going from II B (Zn-group elements) to IIIA sub-group. This sudden fall is explained on the basis that in case of IIIA group elements the electron to be removed is from a 4p-orbital which is incompletely filled, while in case of the II B group elements, the electron to be removed is from 4s-orbital which is completely filled. Thus more energy will be required to remove an electron from a filled 4s-orbital in comparison to that used to remove an electron from a 4p-orbital which is incompletely filled.

Electropositive character of transitional elements as compared to that of alkali metals and alkaline earth metals.

The values of first ionisation potentials of transition elements in most cases lie between those of s-and p-block elements. Thus the transition elements are less electropositive than the elements of I A and II A groups. Thus, although the transitional elements do form ionic compounds, yet they do not form ionic compounds so readily as the alkali and alkaline earth metals do. Also, unlike the alkali and alkaline earth metals, the transitional elements also have a tendency to form the covalent compounds under certain conditions.
Generally the compounds in which the transitional elements show a smaller valency are ionic, while those in which a higher valency is exhibited are covalent in character.

Oxidation States

One of the most important property that distinguishes transition elements from s-and p-block elements is that they show variable oxidation states. s-and p-block elements have oxidation states either equal to their group number G or equal to (8-G). The transition elements on the other hand exhibit variable oxidation states.

This unique property is due to the fact that the energy levels of 3d, 4d and 5d orbitals are fairly close to those of 4s, 5s and 6s orbitals respectively and, therefore, in addition to ns electrons and variable number of (n-1) d electrons are also lost in getting various oxidation states.

(i) Minimum oxidation state: All the transition elements with the exception of Cr, Cu, Ag, Au and Hg which have a minimum oxidation state of +1 exhibit a minimum oxidation state of +2. In most cases this +2 oxidation state arises due to the loss of two s-electrons.

(ii) Maximum oxidation state: Each of the elements in groups III B to VII B can show the maximum oxidation state equal to its group number. Thus, Cr in group VIB shows a maximum oxidation state of +6 in Cr2O72− ion. Most of the elements in VIII group show a maximum oxidation state equal to + 6. However, Ru and Os have a maximum oxidation state equal to +8 which is the highest oxidation state shown by any element.

(iii) Relative stability of various oxidation states: The relative stabilities of various oxidation states of 3d-series elements can be correlated with the extra stability of 3d0, 3d5 and 3d10 configurations to some extent. Thus Ti4+ (3d0) is more stable than Ti3+ (3d1) and similarly Mn2+ (3d5) is more stable than Mn4+ (3d4). It may, however, be pointed out that such a generalisation for the relative stability of various oxidation states of 4d and 5d series elements is often rather difficult to realise.

The higher oxidation state of 4d and 5d series elements are generally more stable than those of the elements of 3d series, e.g., Mo, Tc (4d series elements) and W, Re (5d-series elements) form the oxyanions: MoVIO42− , TcVIIO42−, WVIO42−, ReVIIO4 which are stable and in which the transition elements concerned show their maximum oxidation states. The corresponding oxyanions of Cr and Mn namely CrVIO42− and MnVIIO4 are strong oxidising agents.

Furthermore, the highest oxidation states of second and third row elements are encountered in compounds containing the more electronegative elements viz. F, O and Cl.

Thus for the compounds RuVIIIO4, OsVIIIO4, WVICl6 and PtVIF6 there are no analogs being formed by the first row elements. The lower oxidation states particularly +2 and +3 are important in the chemistry of aquated and complex ions of the 3d-series (i.e. first row) elements but these ions are not very important in the chemistry of second (i.e. 4d series) and third (5d-series) row elements. In short it may be said that in going down a sub-group the stability of the higher oxidation states increases while that of lower oxidation states decreases.

(iv) Formation of ionic and covalent compounds: Transition elements cannot form ionic compounds in higher oxidation states because the loss of more than three electrons is prevented by the higher attractive force exerted (on the electrons) by the nucleus. Higher oxidation states of these metals are not formed by the actual loss of electrons but due to the formation of new hybrid orbitals involving (n-1)d, ns and np orbitals.

The transition elements form ionic bonds in the lower oxidation states and the ionic character of the bond decreases with the increase in the oxidation state. With this decrease in the ionic character of the bond the basic character of the oxides decreases, e.g. MnO (oxidation states of Mn = +2) is basic, MnO2 (Mn = +4) is amphoteric and MnO3 (Mn = +6) is acidic

Colour

Ionic and covalent compounds of transition elements are usually markedly coloured, in contrast to compounds of the s and p block elements which are often white and are generally not strongly coloured. Colour is associated with incomplete electron shells and the ability to promote an electron from one energy level to other. Exactly the right amount of energy to do this is obtained by absorbing the light of a particular wave length.

Illustration : Why Zn+2 salts are white while Ni2+ salts are blue

Solution: Zn+2 has completely filled d-orbitals (3d10) while Ni2+ has incompletely filled d-orbitals (3d8)

In the transition elements, d-electrons are promoted to a higher energy level within the d-subshell. This corresponds to a fairly small energy difference, and so light is absorbed in the visible region. If red light is absorbed then the transmitted light contains an excess of the other colours of the spectrum-particularly blue, so that the compound appears blue, for example Cu2+

Exercise : [Ti(H2O)6]3+ is coloured while [Sc(H2O)6]3+ is colourless. Explain

Complex Formation :

The transition elements have an unparallel tendency to form coordination compounds with Lewis Base, i.e., with groups which are able to donate an electron pair.

These groups are called ligands. A ligand may be a neutral molecule such as NH3 or ion such as Clor CN etc.

Co3+ + 6(NH3) —> [Co(NH3)6]3+

Fe2+ + 6CN —> [Fe(CN)6]4-

The reason transition elements are good in forming complexes are:

(i) Small size and high effective nuclear charge

(ii) Availability of low lying vacant d−orbitals which can accept lone pair of electrons donated by a ligand.

Catalytic properties

Transition metals and their compounds are known to act as good catalyst due to the following reasons:

i) Due to the their variable oxidation state, they form unstable intermediate compounds and provide a new path with lower activation energy for the reaction (Intermediate compound formation theory)

ii) In some cases the finely divided metals or their compounds provide a large surface area for adsorption and the adsorbed reactants react faster due to the closer contact(Adsorption theory)

1. TiCl3 Used as Ziegler – Natta catalyst

2. V2O5 Converts SO2 to SO3 in the contact process for making H2SO4

3. MnO2 Used as a catalyst to decompose KClO3 to give O2

4. Fe Used in Haber – Bosch process for making NH3

5. FeCl3 Production of CCl4 from CS2 and Cl2

6. FeSO4 & H2O2 Fenton’s reagent

7. PdCl2 Wacker process for the following conversion

C2H2 + H2O + PdCl2 –> CH3CHO + 2HCl + Pd

8. Pd For hydrogenation (Phenol –> Cyclohexanone)

9. Pt/PtO Adams catalyst used for reduction

10. Pt SO2 –> SO3 contact process

11. Pt Cleaning car exhaust fumes

12. Cu In manufacture of (CH3)2SiCl2

13. Cu/V Oxidation of cyclohexanol

14. CuCl2 Deacon process for making Cl2 from HCl

15. Ni Raney nickel

Magnetic Properties

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 paramagnetic. Thus paramagnetic materials attract lines of force and if it is free to move, a paramagnetic material will move from a weaker to a stronger part of the field. Paramagnetism arises as a result of unpaired electron in the atom. If the field in the substance is less than H, the substance is diamagnetic. They tend to repel lines of force and move from a strong to weaker part of a magnetic field. In diamagnetic substance, electrons are paired up.

Exercise 2: Why does Mn (II) show maximum paramagnetic character amongst the bivalent ions of the first transition series?

It should be noted that Fe,Co and Ni are ferromagnetic. Ferromagnetic materials may be regarded as special case of paramagnetism in which the moments of individual atoms become aligned and all point in the same direction. It is also possible to get antiferromagnetism by pairing the moments of adjacent atoms which point in opposite directions.
It occurs in salts of Fe3+ , Mn2+ etc.

Paramagnetism is expressed in terms of magnetic moment, which is related to the number of unpaired electrons as follows


n = number of unpaired electrons

B.M. = Bohr Magneton, unit of magnetic moment

More the magnetic moment more is the paramagnetic behaviour

Illustration 2: Calculate the magnetic moment of V3+

Solution: The electronic configuration of V3+ is [Ar] 4s°3d2 .

In the d-orbitals there are 2 unpaired electrons


μ = √2(2+2) = √8 = 2.73 B.M

Illustration 3: Calculate the magnetic moments of Fe2+ and Fe3+

Solution: In Fe2+ there are 4 unpaired electrons.

μ = √4(4+2) = √24 = 4.89 B.M

In Fe3+ there are 5 unpaired electrons.

μ = √5(5+2) = √35 = 5.91 B.M

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