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Table 8.5 shows the stable halides of the 3d series of transition metals. The highest oxidation numbers are achieved in TiX4 (tetrahalides), VF5 and CrF6 . The +7 state for Mn is not represented in simple halides but MnO3F is known, and beyond Mn no metal has a trihalide except FeX3 and CoF3 . The ability of fluorine to stabilise the highest oxidation state is due to either higher lattice energy as in the case of CoF3 , or higher bond enthalpy terms for the higher covalent compounds, e.g., VF5 and CrF6. Although V+5 is represented only by VF5, the other halides, however, undergo hydrolysis to give oxohalides, VOX3 . Another feature of fluorides is their instability in the low oxidation states e.g., VX2 (X = CI, Br or I)and the same applies to CuX. On the other hand, all CuII halides are known except the iodide. In this case, Cu2+ oxidises I– to I2: ( ) 2 2Cu 4I Cu I I 2 2 2 s + − + → + However, many copper (I) compounds are unstable in aqueous solution and undergo disproportionation. 2Cu+ → Cu2+ + Cu The stability of Cu2+ (aq) rather than Cu+ (aq) is due to the much more negative ∆hydH V of Cu2+ (aq) than Cu+ , which more than compensates for the second ionisation enthalpy of Cu. The ability of oxygen to stabilise the highest oxidation state is demonstrated in the oxides. The highest oxidation number in the oxides (Table 8.6) coincides with the group number and is attained in Sc2O3 to Mn2O7 . Beyond Group 7, no higher oxides of Fe above Fe2O3 , are known, although ferrates (VI)(FeO4 ) 2–, are formed in alkaline media but they readily decompose to Fe2O3 and O2 . Besides the oxides, oxocations stabilise Vv as VO2 + , VIV as VO2+ and Ti IV as TiO2+. The ability of oxygen to stabilise these high oxidation states exceeds that of fluorine. Thus the highest Mn fluoride is MnF4 whereas the highest oxide is Mn2O7 . The ability of oxygen to form multiple bonds to metals explains its superiority. In the covalent oxide Mn2O7 , each Mn is tetrahedrally surrounded by O’s including a Mn–O–Mn bridge. The tetrahedral [MO4 ] nions are known for VV , CrVl, MnV , MnVl and MnVII .

chemical properties :

Transition metals vary widely in their chemical reactivity. Many of them are sufficiently electropositive to dissolve in mineral acids, although a few are ‘noble’—that is, they are unaffected by single acids. The metals of the first series with the exception of copper are relatively more reactive and are oxidised by 1M H+ , though the actual rate at which these metals react with oxidising agents like hydrogen ion (H+ ) is sometimes slow. For example, titanium and vanadium, in practice, are passive to dilute non oxidising acids at room temperature. The E V values for M2+/M (Table 8.2) indicate a decreasing tendency to form divalent cations across the series. This general trend towards less negative E V values is related to the increase in the sum of the first and second ionisation enthalpies. It is interesting to note that the E V values for Mn, Ni and Zn are more negative than expected from the general trend. Whereas the stabilities of half-filled d subshell (d 5 ) in Mn2+ and completely filled d subshell (d 10) in zinc are related to their E e values; for nickel, E o value is related to the highest negative enthalpy of hydration. An examination of the E V values for the redox couple M3+/M2+ (Table 8.2) shows that Mn3+ and Co3+ ions are the strongest oxidising agents in aqueous solutions. The ions Ti 2+, V2+ and Cr2+ are strong reducing agents and will liberate hydrogen from a dilute acid, e.g., 2 Cr2+(aq) + 2 H+ (aq) → 2 Cr3+(aq) + H2 (g)

magnetic properties:

When a magnetic field is applied to substances, mainly two types of magnetic behaviour are observed: diamagnetism and paramagnetism (Unit 1). Diamagnetic substances are repelled by the applied field while the paramagnetic substances are attracted. Substances which areattracted very strongly are said to be ferromagnetic. In fact, ferromagnetism is an extreme form of paramagnetism. Many of the transition metal ions are paramagnetic. Paramagnetism arises from the presence of unpaired electrons, each such electron having a magnetic moment associated with its spin angular momentum and orbital angular momentum. For the compounds of the first series of transition metals, the contribution of the orbital angular momentum is effectively quenched and hence is of no significance. For these, the magnetic moment is determined by the number of unpaired electrons and is calculated by using the ‘spin-only’ formula, i.e., µ = n( ) n 2 + where n is the number of unpaired electrons and µ is the magnetic moment in units of Bohr magneton (BM). A single unpaired electron has a magnetic moment of 1.73 Bohr magnetons (BM). The magnetic moment increases with the increasing number of unpaired electrons. Thus, the observed magnetic moment gives a useful indication about the number of unpaired electrons present in the atom, molecule or ion. The magnetic moments calculated from the ‘spin-only’ formula and those derived experimentally for some ions of the first row transition elements are given in Table 8.7. The experimental data are mainly for hydrated ions in solution or in the solid state.

formation of coloured ions

When an electron from a lower energy d orbital is excited to a higher energy d orbital, the energy of excitation corresponds to the frequency of light absorbed (Unit 9). This frequency generally lies in the visible region. The colour observed corresponds to the complementary colour of the light absorbed. The frequency of the light absorbed is determined by the nature of the ligand. In aqueous solutions where water molecules are the ligands, the colours of the ions observed are listed in Table 8.8. A few coloured solutions of d–block elements are illustrated in Fig. 8.5

Formation of Complex Compounds Intext Question 8.8 Calculate the ‘spin only’ magnetic moment of M2+ (aq) ion (Z = 27). When an electron from a lower energy d orbital is excited to a higher energy d orbital, the energy of excitation corresponds to the frequency of light absorbed (Unit 9). This frequency generally lies in the visible region. The colour observed corresponds to the complementary colour of the light absorbed. The frequency of the light absorbed is determined by the nature of the ligand. In aqueous solutions where water molecules are the ligands, the colours of the ions observed are listed in Table 8.8. A few coloured solutions of d–block elements are illustrated in Fig. 8.5. 8.3.10 Formation of Coloured Ions Fig. 8.5: Colours of some of the first row transition metal ions in aqueous solutions. From left to right: V4+,V3+,Mn2+,Fe3+,Co2+,Ni2+and Cu2+ . Complex compounds are those in which the metal ions bind a number of anions or neutral molecules giving complex species with characteristic properties. A few examples are: [Fe(CN)6 ] 3–, [Fe(CN)6 ] 4– , [Cu(NH3 )4 ] 2+ and [PtCl4 ] 2–. (The chemistry of complex compounds is crucial for understanding the behavior of transition metals in various applications, including catalysis, materials science, and biological systems.

catalytic properties

The transition metals and their compounds are known for their catalytic activity. This activity is ascribed to their ability to adopt multiple oxidation states and to form complexes. Vanadium(V) oxide (in Contact Process), finely divided iron (in Haber’s Process), and nickel (in Catalytic Hydrogenation) are some of the examples. Catalysts at a solid surface involve the formation of bonds between reactant molecules and atoms of the surface of the catalyst (first row transition metals utilise 3d and 4s electrons for bonding). This has the effect of increasing the concentration of the reactants at the catalyst surface and also weakening of the bonds in the reacting molecules (the activation energy is lowering). Also because the transition metal ions can change their oxidation states, they become more effective as catalysts. For example, iron(III) catalyses the reaction between iodide and persulphate ions. 2 I– + S2O8 2– → I2 + 2 SO4 2– An explanation of this catalytic action can be given as: 2 Fe3+ + 2 I– → 2 Fe2+ + I2 2 Fe2+ + S2O8 2– → 2 Fe3+ + 2SO4 2–

formation of interstitial compounds: Interstitial compounds are those which are formed when small atoms like H, C or N are trapped inside the crystal lattices of metals. They are usually non stoichiometric and are neither typically ionic nor covalent, for example, TiC, Mn4N, Fe3H, VH0 .56 and TiH1.7, etc. The formulas quoted do not, of course, correspond to any normal oxidation state of the metal. Because of the nature of their composition, these compounds are referred to as interstitial compounds. The principal physical and chemical characteristics of these compounds are as follows: (i) They have high melting points, higher than those of pure metals. (ii) They are very hard, some borides approach diamond in hardness. (iii) They retain metallic conductivity. (iv) They are chemically inert.

alloy formation: An alloy is a blend of metals prepared by mixing the components. Alloys may be homogeneous solid solutions in which the atoms of one metal are distributed randomly among the atoms of the other. Such alloys are formed by atoms with metallic radii that are within about 15 percent of each other. Because of similar radii and other characteristics of transition metals, alloys are readily formed by these metals. The alloys so formed are hard and have often high melting points. The best known are ferrous alloys: chromium, vanadium, tungsten, molybdenum and manganese are used for the production of a variety of steels and stainless steel. Alloys of transition metals with non transition metals such as brass (copper-zinc) and bronze (copper-tin), are also of considerable industrial importance.