`=>` One of the notable features of a transition element is the great variety of oxidation states it may show in its compounds.

`=>` Table 8.3 lists the common oxidation states of the first row transition elements.

`=>` The elements which give the greatest number of oxidation states occur in or near the middle of the series.

● Manganese, for example, exhibits all the oxidation states from `+2` to `+7`.

`=>` The lesser number of oxidation states at the extreme ends stems from either too few electrons to lose or share (`color{red}(Sc)`, `color{red}(Ti)`) or too many `color{red}(d)` electrons (hence fewer orbitals available in which to share electrons with others) for higher valence (`color{red}(Cu)`, `color{red}(Zn)`).

● Thus, early in the series scandium(II) is virtually unknown and titanium (IV) is more stable than `color{red}(Ti )`(`III`) or `color{red}(Ti)`(`II`).

● At the other end, the only oxidation state of zinc is `+2` (no `color{red}(d)` electrons are involved).

● The maximum oxidation states of reasonable stability correspond in value to the sum of the `s` and `d` electrons upto manganese (`color{red}(Ti^(IV)O_2)`, `color{red}(V^(V)O_2^+)`, `color{red}(Cr^(VI)O_4^(2-))`, `color{red}(Mn^(VII)O_4^-)`) followed by a rather abrupt decrease in stability of higher oxidation states, so that the typical species to follow are `color{red}(Fe^(II,III))`, `color{red}(Co^(II,III))`, `color{red}(Ni^(II))`, `color{red}(Cu^(I,II))`, `color{red}(Zn^(II))`.

`=>` The variability of oxidation states, a characteristic of transition elements, arises out of incomplete filling of `d` orbitals in such a way that their oxidation states differ from each other by unity, e.g., `color{red}(V^(II))`, `color{red}(V^(III))`, `color{red}(V^(IV))`, `color{red}(V^V)`.

● This is not in agreement with the variability of oxidation states of non transition elements where oxidation states normally differ by a unit of two.

`=>` An interesting feature in the variability of oxidation states of the `color{red}(d)`–block elements is noticed among the groups (groups `4` through `10`).

`color{red}(text(Note )) : ` In the `color{red}(p)`–block the lower oxidation states are favoured by the heavier members (due to inert pair effect), the opposite is true in the groups of `color{red}(d)`-block.

● For example, in group 6, `color{red}(Mo(VI))` and `color{red}(W(VI))` are found to be more stable than `color{red}(Cr(VI))`. Thus `color{red}(Cr(VI))` in the form of dichromate in acidic medium is a strong oxidising agent, whereas `color{red}(MoO_3)` and `color{red}(WO_3)` are not.

`=>` Low oxidation states are found when a complex compound has ligands capable of `color{red}(π)`-acceptor character in addition to the `color{red}(σ)`-bonding. For example, in `color{red}(Ni(CO)_4)` and `color{red}(Fe(CO)_5)`, the oxidation state of nickel and iron is zero.

`=>` Table 8.4 contains the thermochemical parameters related to the transformation of the solid metal atoms to `color{red}(M^(2+))` ions in solution and their standard electrode potentials.

`=>` The observed values of `color{red}(E^⊖)` and those calculated using the data of Table 8.4 are compared in Fig. 8.4.

`=>` The unique behaviour of `color{red}(Cu)`, having a positive `color{red}(E^⊖)`, accounts for its inability to liberate `color{red}(H_2)` from acids.

`=>` Only oxidising acids (nitric and hot concentrated sulphuric) react with `color{red}(Cu)`, the acids being reduced.

`=>` The high energy to transform `color{red}(Cu(s))` to `color{red}(Cu^(2+))`(aq) is not balanced by its hydration enthalpy.

`=>` The general trend towards less negative `color{red}(E^⊖)` values across the series is related to the general increase in the sum of the first and second ionisation enthalpies.

`color{red}(text(Note ))` : The value of `color{red}(E^(⊖))` for `color{red}(Mn)`, `color{red}(Ni)` and `color{red}(Zn)` are more negative than expected from the trend.

`=>` An examination of the `color{red}(E (M^(3+) | M^(2+) ))` values (Table 8.2) shows the varying trends.

`=>` The low value for `color{red}(Sc)` reflects the stability of `color{red}(Sc^(3+))` which has a noble gas configuration.

`=>` The highest value for `color{red}(Zn)` is due to the removal of an electron from the stable `color{red}(d^(10))` configuration of `color{red}(Zn^(2+))`.

`=>` The comparatively high value for `color{red}(Mn)` shows that `color{red}(Mn^(2+) (d^5))` is particularly stable, whereas comparatively low value for `color{red}(Fe)` shows the extra stability of `color{red}(Fe^(3+) (d^5))`.

`=>` The comparatively low value for `color{red}(V)` is related to the stability of `color{red}(V^(2+))` (half-filled `color{red}(t_(2g))` level).

`=>` Table 8.5 shows the stable halides of the `color{red}(3d)` series of transition metals.

`=>` The highest oxidation numbers are achieved in `color{red}(TiX_4)` (tetrahalides), `color{red}(VF_5)` and `color{red}(CrF_6)`.

`=>` The `+7` state for `color{red}(Mn)` is not represented in simple halides but `color{red}(MnO_3F)` is known, and beyond `color{red}(Mn)` no metal has a trihalide except `color{red}(FeX_3)` and `color{red}(CoF_3)`.

`=>` The ability of fluorine to stabilise the highest oxidation state is due to either higher lattice energy as in the case of `color{red}(CoF_3)`, or higher bond enthalpy terms for the higher covalent compounds, e.g., `color{red}(VF_5)` and `color{red}(CrF_6)`.

● `color{red}(V^V)` is represented only by `color{red}(VF_5)`, the other halides, however, undergo hydrolysis to give oxohalides, `color{red}(VOX_3)`.

● Another feature of fluorides is their instability in the low oxidation states e.g., `color{red}(VX_2)` (`color{red}(X = CI)`, `color{red}(Br)` or `color{red}(I)`) and the same applies to `color{red}(CuX)`.

`=>` All `color{red}(Cu^I)` halides are known except the iodide. In this case, `color{red}(Cu^(2+))` oxidises `color{red}(I^-)` to `color{red}(I_2)` :

`color{red}(2Cu^(2-) + 4I^(-) -> Cu_2 I_2 (s) + I_2)`

`=>` But, many copper (I) compounds are unstable in aqueous solution and undergo disproportionation.

`color{red}(2Cu^(+) → Cu^(2+) + Cu)`

● The stability of `color{red}(Cu^(2+))` (aq) rather than `color{red}(Cu^(+))` (aq) is due to the much more negative `color{red}(Δ_(hyd)H^⊖)` of `color{red}(Cu^(2+))` (aq) than `color{red}(Cu^(+))`, which more than compensates for the second ionisation enthalpy of `color{red}(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 `color{red}(Sc_2O_3)` to `color{red}(Mn_2O_7)`.

● Beyond Group 7, no higher oxides of `color{red}(Fe)` above `color{red}(Fe_2O_3)`, are known, although ferrates (VI) `color{red}((FeO_4)^(2–))`, are formed in alkaline media but they readily decompose to `color{red}(Fe_2O_3)` and `color{red}(O_2)`.

● Besides the oxides, oxocations stabilise `color{red}(V^v)` as `color{red}(VO_2^+)`, `color{red}(V^(IV))` as `color{red}(VO^(2+))` and `color{red}(Ti^(IV))` as `color{red}(TiO^(2+))`.

`=>` The ability of oxygen to stabilise these high oxidation states exceeds that of fluorine. Thus, the highest `color{red}(Mn)` fluoride is `color{red}(MnF_4)` whereas the highest oxide is `color{red}(Mn_2O_7)`.

`=>` The ability of oxygen to form multiple bonds to metals explains its superiority.

`=>` In the covalent oxide `color{red}(Mn_2O_7)`, each `color{red}(Mn)` is tetrahedrally surrounded by `color{red}(O’s)` including a `color{red}(Mn–O–Mn)` bridge.

`=>` The tetrahedral `color{red}([MO_4]^n)` ions are known for `color{red}(V^V)`, `color{red}(Cr^(Vl))`, `color{red}(Mn^(V))`, `color{red}(Mn^(Vl))` and `color{red}(Mn^(VII))`.