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In p-block elements, higher oxidation states are less stable down the group due to the inert pair effect. This is not the case for transition metals.

Why do heavier transition metals show higher oxidation states than lighter ones? Is the inert pair effect not valid for transition metals also?

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The inert pair effect is based on the fact that main group elements’ oxidation states depend on s and p orbitals (and only them). When going down the periodic table, the energy difference between s and p orbitals changes leading to some elements losing their valence s electrons more easily than others.

Transition metals’ chemistry happens in the d orbitals primarily — unless you count the copper and zinc groups wherein a significant part of the chemistry is in fact only s orbital chemistry. The d orbitals — at first approximation of the free ion — are degenerate, i.e. they all have the same energy. Thus, these electrons are typically much more accessable. Furthermore, going down the periodic table increases the number of electrons counted as core electrons meaning that the outermost valene electrons experience a weaker effective nuclear attraction. It is therefore easier (i.e. requires less energy) to remove valence electrons and higher oxidation states are much more accessible.

If it weren’t for the inert pair effect, this would also be visible in main group chemistry — except it is, if you compare the relative stabilities of iodate and chlorate.

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The sum of the total ionaisation energies for achieving the highest oxidation state in a particular group is less for the heavier elements than for the lighter elements.Therefor stability of higher oxidation states increases down the group.

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    $\begingroup$ So what makes the total ionization energy drop? $\endgroup$ – Oscar Lanzi Mar 25 '18 at 14:19

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