`=>` Nearly all the transition elements display typical metallic properties such as :

(i) high tensile strength

(ii) ductility malleability

(iii) high thermal and electrical conductivity

(iv) metallic lustre

● With the exceptions of `color{red}(Zn)`, `color{red}(Cd)`, `color{red}(Hg)` and `color{red}(Mn)`, they have one or more typical metallic structures at normal temperatures.

`=>` The transition metals (with the exception of `color{red}(Zn)`, `color{red}(Cd)` and `color{red}(Hg)`) are very much hard and have low volatility.

`=>` Their melting and boiling points are high. Fig. 8.1 depicts the melting points of the `color{red}(3d)`, `color{red}(4d)` and `color{red}(5d)` transition metals.

● The high melting points of these metals are attributed to the involvement of greater number of electrons from `color{red}((n-1)d)` in addition to the `color{red}(ns)` electrons in the interatomic metallic bonding.

● In any row the melting points of these metals rise to a maximum at `color{red}(d^5)` except for anomalous values of `color{red}(Mn)` and `color{red}(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 `color{red}(d)` orbital is particularly 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.

● 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.

`=>` 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 `color{red}(d)`-orbital each time the nuclear charge increases by unity. And we know that the shielding effect of a `color{red}(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 (`color{red}(3d)`) to the second (`color{red}(4d)`) series of the elements but the radii of the third (`color{red}(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 `color{red}(4f)` orbitals which must be filled before the `color{red}(5d)` series of elements begin.

● The filling of `color{red}(4f)` before `color{red}(5d)` orbital results in a regular decrease in atomic radii called `color{green}(text(Lanthanoid contraction))` which essentially compensates for the expected increase in atomic size with increasing atomic number.

● The net result of the lanthanoid contraction is that the second and the third `color{red}(d)` series exhibit similar radii (e.g., `color{red}(Zr)` `160` pm, `color{red}(Hf)` `159` pm) and have very similar physical and chemical properties.

● The factor responsible for the lanthanoid contraction is attributed to the imperfect shielding of one electron by another in the same set of orbitals.

● However, the shielding of one `color{red}(4f)` electron by another is less than that of one `color{red}(d)` electron by another, and as the nuclear charge increases along the series, there is fairly regular decrease in the size of the entire `color{red}(4f^n)` 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 (`color{red}(Z = 22)`) to copper (`color{red}(Z = 29)`) the significant increase in the density may be noted .

`=>` Due to an increase in nuclear charge which accompanies the filling of the inner `color{red}(d)`-orbitals, there is an increase in ionisation enthalpy along each series of the transition elements from left to right. However, many small variations occur.

`=>` Table 8.2 gives the values for the first three ionisation enthalpies of the first row elements.

● These values show that the successive enthalpies of these elements do not increase as steeply as in the main group elements.

● Although, the first ionisation enthalpy increases, the magnitude of the increase in the second and third ionisation enthalpies for the successive elements, in general, is much higher.

`=>` The irregular trend in the first ionisation enthalpy of the `color{red}(3d)` metals, though of little chemical significance, can be accounted for by considering that the removal of one electron alters the relative energies of `color{red}(4s)` and `color{red}(3d)` orbitals. So, the unipositive ions have `d^n` configurations with no `color{red}(4s)` electrons.

● There is thus, a reorganisation energy accompanying ionisation with some gains in exchange energy as the number of electrons increases and from the transference of `color{red}(s)` electrons into `color{red}(d)` orbitals.

`=>` There is the generally expected increasing trend in the values as the effective nuclear charge increases.

● However, the value of `color{red}(Cr)` is lower because of the absence of any change in the `color{red}(d)` configuration and the value for `color{red}(Zn)` higher because it represents an ionisation from the `color{red}(4s)` level.

`=>` The lowest common oxidation state of these metals is `+2`.

● To form the `color{red}(M^(2+))` ions from the gaseous atoms, the sum of the first and second ionisation energies is required in addition to the enthalpy of atomisation for each element.

● The dominant term is the second ionisation enthalpy which shows unusually high values for `color{red}(Cr)` and `color{red}(Cu)` where the `color{red}(d^5)` and `color{red}(d^10)` configurations of the `color{red}(M^+)` ions are disrupted, with considerable loss of exchange energy.

● The value for `color{red}(Zn)` is correspondingly low as the ionisation consists of the removal of an electron which allows the production of the stable `color{red}(d^(10))` configuration.

`=>` The trend in the third ionisation enthalpies is not complicated by the `color{red}(4s)` orbital factor and shows the greater difficulty of removing an electron from the `color{red}(d^5)` (`color{red}(Mn^(2+))`) and `color{red}(d^(10))` (`color{red}(Zn^(2+))`) ions superimposed upon the general increasing trend.

● In general, the third ionisation enthalpies are quite high and there is a marked break between the values for `color{red}(Mn^(2+))` and `color{red}(Fe^(2+))`.

`=>` Also, the high values for copper, nickel and zinc indicate why it is difficult to obtain oxidation state greater than two for these elements.