# How do I know the likely oxidation states of metals, and non-metals?

Nov 21, 2016

Well main group metal ions typically form an ion whose charge is equal to their Group number............i.e. isoelectronic with the last Noble Gas.

#### Explanation:

........And non-metals typically form an ion isoelectronic with the next Noble Gas.

A Group 17 element typically forms an ion with a single negative charge:

$\frac{1}{2} {X}_{2} + {e}^{-} \rightarrow {X}^{-}$ $X = F , C l , B r \ldots \ldots .$

And a Group 16 element typically forms an ion with a double negative charge:

$\frac{1}{2} {O}_{2} \left(g\right) + 2 {e}^{-} \rightarrow {O}^{2 -}$

In each case the element has formed an ion isoelectronic with the next (or last) Noble Gas. And we can go even farther than this, and consider Group 15.

$P \left(s\right) + 3 {e}^{-} \rightarrow {P}^{3 -}$

And we can look at oxidation of the alkali metals (Group I):

$M \left(s\right) \rightarrow {M}^{+} + {e}^{-}$ $M = L i , N a , K , e t c .$

And of the alkaline earths (Group 2):

$M \left(s\right) \rightarrow {M}^{2 +} + 2 {e}^{-}$ $M = C a , B a , S r , e t c .$

The point is that the Group number reflects electronic structure, i.e. the number of electrons present in the valence shell. Group I and Group II metals have 1 and 2 valence electrons respectively. As atomic number increases across a Period, nuclear charge increases accordingly, and non-metals, to the RIGHT of the Period as we face it, tend to be oxidizing, and ADD electrons to their valence shell.