Oxidation Numbers

How do you calculate oxidation numbers?

Oxidation numbers are used to track the charges of atoms in a molecule such that their sum equals the entire charge of the molecule.

What is the role of electronegativities when predicting the oxidation numbers of an atom?

Atoms that are more electronegative, i.e. the upper right corner of the periodic table, will tend to have negative oxidation numbers, while atoms that are less electronegative, i.e. the lower left corner of the periodic table, will have positive oxidation numbers. Intuitively, it makes sense: an atom with a high electronegativity will want electrons and be negative (and therefore serves as an oxidizing agent - an electron remover).

What are the general oxidation numbers for the groups across the periodic table?

First off I use the term ``general'' because oxidation numbers can vary greatly, but these are the general numbers assigned to the groups across the periodic table:

Secondly, an atom can either donate or receive electrons depending upon if the atom is pairing up with an atom that is more or less electronegative, respectively:

Table 39.1: Oxidation numbers for Groups across the periodic table.

	Group:
	I	II	III	IV	V	VI	VII	VIII
Oxidation number of the atom if it is gaining $e^{-}$				-4	-3	-2	-1	0
Oxidation number of the atom if it is losing $e^{-}$	+1	+2	+3	+4	+5	+5	+7	0

What makes for a good oxidizing agent and reducing agent?

Again, electron affinity comes into play. Atoms or molecules with large affinities for electrons (top right-hand corner of the periodic table) are good oxidizing agents (electron removers), while atoms in the lower left-hand corner with lower affinities for electrons are good reducing agents (electron donors).

Therefore, fluorine makes an excellent oxidizing agent as do the oxygen species (O$_{2}$ and O$_{3}$) and chlorine. In regards to molecules that make good oxidizing agents, common ones are permanganate (MnO$_{4}^{-}$), chromate (CrO$_{4}^{2-}$), and dichromate (Cr$_{2}$O$_{7}^{2-}$).

For common reducing agents, look towards Group I atoms such as sodium (Na) and other metals such as zinc (Zn), aluminum (Al) and calcium (Ca).

From the example above (35.1), what are the oxidation numbers of the reactants and products?

Complete Reaction:

$2 MnO_{4 \ (aq)}^{-} + 5 H_{2}C_{2}O_{4 \ (aq)} + 6 H^{+}_{(aq)} \rightarrow 10 CO_{2 \ (g)} + 2 Mn^{2+}_{(aq)} + 8 H_{2}O_{(l)}$

Reactants being oxidized:

$H_{2}C_{2}O_{4} \rightarrow CO_{2}$

H$_{2}$	C$_{2}$	O$_{4}$		C	O$_{2}$
+1	+3	-2		+4	-2
x 2	x 2	x 4	Net charge:	x 1	x 2	Net charge:
+2	+6	-8	= 0	+4	-4	= 0

Reactants being reduced:

$MnO_{4} \rightarrow Mn^{2+}$

Mn	O$_{4}$		Mn$^{2+}$
+7	-2		+2
x 1	x 4	Net charge:	x 1	Net charge:
+7	-8	= -1	+2	= 0