Transition elements having partly filled d-orbitals exhibit several interesting properties. For example, they exhibit variable oxidation states, form coloured complexes with different anions and neutral molecules and show paramagnetic behaviour. Transition metals and their compounds also possess catalytic properties. Some of the important properties of transition metals are discussed below:
All transition elements are metallic in nature. Transition elements exhibit typical metallic properties such· as lustre, high tensile strength, ductility, malleability and high thermal and electrical conductivity. They show gradual decrease in · ·, electropositive character in moving from left to right. The . metallic bonds in transition metals are quite strong. This is due to the greater effective nuclear charge and the large number of valence electrons. Due to the presence of strong metallic bonds, the transition metals are hard, possess high densities and high energies of Atomisation. Atomisation energies of the first transition series are represented graphically in Fig. 19.1.
It may be observed that atomisation energies exhibit the maxima at about the middle of the series. It indicates that interatomic interactions become stronger with increase in half filled d-orbitals.
Most of the transition elements have densities higher than 5 g cm-3, the highest being that of iridium (22.6 g cm-3) . Scandium has the least density among transition metals.
MELTING AND BOILING POINTS
The melting and boiling points of transition elements are generally very high. This is due to strong metallic bond and the presence of half-filled d-orbitals in them. Due to these half-filled orbitals, some covalent bonds also exist between atoms of transition elements. Because of stronger interatomic bonding, transition elements have high melting and boiling points.
In moving along the period from left to right, the melting points of these metals first increase to maximum and then decrease regularly towards the end of the period. In any row the melting points of these metals rise to a maximum at.
The ionization energies of transition elements are higher than those of s-block elements but lower than p-block elements. In a particular transition series, ionization energy increases gradually as we move from left to right However, the relative difference of ionization energy values of any two consecutive d-block elements of particular period is much smaller than those of s- and p-block elements. Reason. The increase in ionization energy 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 ionization energy along the period of d-block elements is very small.
Table 19.2. Some Physical Properties of the First Row Transition Elements
The ionization energies of 3d-transition series are given in Table 19.2 and graphically represented in Fig. 19.2.
Fig. 19.2. Ionization energies of first transition series.
Significance of lm1hafn Energy Values
The magnitudes of ionization energies give some indication of the energy required to raise the metal to a particular oxidation state. From the knowledge of values of ionization energies of the metals it is possible to rationalize the relative stabilities of various oxidation states. For example. the sums of first two and first four ionization energies of nickel and platinum are given in Table 19.3.
Table 19.3. Ionization Energies of Nickel and Platinum
Since sum of the first two ionization energies is less for nickel, therefore, Ni(II) compounds are thermodynamically more stable than Pt(II) compounds. On the other hand, Pt(IV) compounds are more stable than Ni(IV) compounds because sum of first four ionization energies is less for platinum. K2PtC16 is well known compound of platinum with +4 oxidation state. The corresponding nickel compound does not exist In addition to ionization energy, the other factors that determine the stability of a particular state are the atomisation energy of the metal and the lattice energy or the solvation energy.
ATOMIC AND IONIC RADII
The atomic and ionic radii of transition elements are smaller than those of s-block elements and larger than those of p-block elements. Among the elements of the particular transition series. as the atomic number increases, the atomic radii first -decrease till the middle, become almost constant and then increase towards the end of the period. For example, the atomic radii of first transition series decrease from Sc to Cr. remain almost constant till Cu and then increase towards the end.
Table 19.4. Atomisation Energies, Ionization Energies and
Standard Electrode Potentials for Elements of First Transition Series
Table 19.5. Atomic and Ionic Radii of Elements of First Transition Series
Reason. The decrease in size in the beginning is attributed to the increase in nuclear charge. 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 become almost constant Increase in atomic radii towards the end may be attributed to the electron-electron repulsions. In fact, the pairing of electrons in d-orbitals occurs after d5 configuration. The repulsive interactions between the paired electrons in d-orbitals become very dominant towards the end of the period and cause the expansion of electron cloud and thus, resulting in increased atomic size. The ionic radii also follow the similar trend.
All transition metals exhibit a great variety of oxidation states. The oxidation states of first row transition elements are listed in Table 19.6. The less common and unstable oxidation states are given in the parentheses. The stability of a particular oxidation state depends upon the nature of the element with which the transition metal forms the compound. The highest oxidation states are found in compounds of fluorine and oxygen. This is due to the high electronegativity values and small size of fluorine and oxygen.
The variable oxidation states of transition elements are due to the participation of ns and (n -1) d-electrons in bonding. The lower oxidation state is generally, exhibited when
ns-electrons participate in bonding and higher oxidation states are shown when ns as well as (n- 1) d-electrons take part in bonding. It may be noted the oxidation states of transition elements differ from each other by unity whereas oxidation states of non-transition elements generally differ by two.
Some noteworthy features of oxidation states of the transition elements are:
1. In each group, the highest oxidation state increases with increase in atomic number, reaches a maximum in the middle and then starts decreasing. For example, for the first transition series the maximum oxidation state is shown by manganese.
Table 19.6. Different Oxidation States of Transition Metals
Note. Less common and unstable oxidation states are given in parentheses.
2. The elements which exhibit the maximum number of oxidation states occur either in or near the middle of the series. For example, in the first transition series manganese exhibits maximum number of oxidation states ( + 2 to + 7).
3. The elements in the beginning of the series exhibit fewer oxidation states because they have small number of electrons which they can lose or contribute for sharing. The elements at the end of the series exhibit fewer oxidation states because they have too many d-electrons and hence have fewer vacant d-orbitals which can be involved in bonding.
4. For the elements of first transition series (except scandium) + 2 oxidation state is the most common oxidation state. This oxidation state arises due to the loss of 4s-electrons.
5. The transition elements in lower oxidation states ( + 2 and + 3) generally form ionic bonds. In higher oxidation states, the bonds formed are essentially covalent. For example, in tetraoxochromate(VI) ion (CrO42-). the bonds formed between chromium and oxygen are covalent.
6. Some transition metals also show oxidation state of zero in their compounds. [Ni(CO)4] and [Fe(CO)5] are common examples.
In contrast to the representative elements, transition elements form many coordination complexes. Their tendency to form complexes is attributed to the following reasons:
1. Small size and high charge density of the ions of transition metals.
2. Presence of vacant orbitals of appropriate energy which can accept lone pairs of electrons donated by other groups (ligands).
Some examples of coordination complexes are:
(i) [AgH3)2] Cl (ii) K4[Fe(CN)6]
Coordination complexes have been discussed in detail in Section 19.4.
The compounds of transition elements are usually coloured both in solid state and in aqueous solution. The colour of these complexes is due to absorption of some radiation from visible light, which is used in promoting an electron from one of the d-orbitals to another. This can be explained as under:
The d-orbitals in the transition elements do not have same energy in their complexes. Under the influence of the ligands attached, the d-orbitals split into two sets of orbitals having slightly different energies. In the transition elements, which have partly filled d-orbitals, the transition of electron can take place from one of the lower d-orbitals to some higher d-orbital within the same subshell. The energy required for this transition falls in the visible region. So when white light falls on these complexes they absorb a particular colour from the radiation for the promotion of electron and the remaining colours are emitted. The colour of the complex is due to this emitted radiation. For example, copper(IT) salts are bluish green due to absorption of red light. Ti3+ salts appear purple due to absorption of yellow light.
The energy difference between the two sets of d-orbitats in the central atom of the complex depends on the nature of ligands and the structure of the complex ion. As a result different complexes of the same metal ion, with different ligands, may have different colours. For example, [COC14f is blue in colour whereas [CO(H2O)6] 2+ is pink[Fe(H2O)6 ]2+ is green in colour whereas [Fe(CN)6]4- is yellow.
Zn2+ and Ti4+ salts are white because they do not absorb any radiation in the visible region. In these compounds, d-d transitions are not possible because in Zn2+ all the d-orbitals are fully filled whereas in Ti4+ all the d-orbitals are vacant. The colours of some transition metal ions in aqueous solutions are given in Table 19.7.
Table 19.7. The Colours of Some Transition Metal Ions in Aqueous Solution
The substances, which contain some species (atoms, ions or molecules) with unpaired electrons in their orbitals, behave as paramagnetic substances. Such substances are weakly attracted by magnetic field. On the other hand, the substances whose constituent particles do not contain any unpaired electrons are repelled by magnetic field and are called diamagne1ic.
The transition metal ions generally contain one or more unpaired electrons in them and hence their complexes are generally paramagnetic. The paramagnetic character increases with increase in number of unpaired electrons.
Many transition metals and their compounds are known to act as catalysts. For example, finely divided iron acts as catalyst in the manufacture of ammonia by Haber Process. vanadium pentoxide (V2O5) or platinum act as catalyst for the oxidation of SO2 to SO3 in Contact Process, ferrous sulphate and hydrogen peroxide (Fenton’s reagent) are used for the oxidation of alcohols to aldehydes.
The catalytic activity of transition metals is attributed to the following reasons:
l. Because of their variable oxidation states transition metals sometimes form unstable intermediau compounds and provide a new path with lower activation energy for the reaction.
For example, v p5 catalyses the oxidation of SO2 to SO3.
The catalytic action of V2O5 can be understood a5 follows:
During the conversion of SO2 to SO3, V2O5 adsorbs SO: molecule on its surface and gives oxygen to it to form SO, and V2O 4. V2O 4 then reacts with oxygen to form V2O5
Similarly, iron(III) catalyses the reaction between iodide and persulphate ions.
The catalytic action of iron(III) in this reaction is explained as follows
2Fe3+ + 21 à 2Fe2+ + I2
2Fe2+ + S2O82- à 2Fe3+ + 2SO4
2. In some cases transition metals provide a suitable surface of the reaction to take place. The reactant are adsorbed on the surface of the catalyst where reaction occurs.
Adsorption results in increased concentration of reactants at the surface and also weakens the bonds between atoms in the reactant molecules.
Catalytic Properties of Transition Elements and their Compounds
Some first row metals and their compounds used as catalysts are given below:
Catalyst Process Catalysed
TiC14 Used as the Ziegler-Natta catalyst in the
polymerization of ethene and propene
V2O 5 Used as catalyst during conversion of S02 to S03
in the Contact process for the manufacture of
MnO2 Used as catalyst to decompose KC103 to produce
2KC1O3 à 2KCI + 3O2
Fe Iron in the presence of a promotor act as catalyst in
Haber process for the manufacture of ammonia
FeCJ3 Used as catalyst in the production of CC14 from
cs2 and cl2
Co2(CO)8 Oxo process for conversion of alkenes to alkanals
Ni Hydrogenation of vegetable oils
CuCl2 Used as catalyst in the manufacture of chlorine from
HCI (Deacon process).
The catalytic activity of transition metal compounds can be demonstrated by the following activity.