Section 7-2

Oxidation Numbers (aka Oxidation States)

• Are used to indicate and assign the general distribution of electrons among the bonded atoms in a molecular compound or ion.

• In an ionic compound, the charges of each ion tell us about the electron distribution

• In a covalent compound, the atoms are sharing electrons. But we can imagine that the atoms inside the molecules are charged, and we call these charges oxidation numbers.

Oxidation Numbers (aka Oxidation States)

• Essentially, oxidation numbers help us to keep track of electrons in reactions.

• They are a book-keeping mechanism to help us track electron distribution – it is helpful to know what is happening with electrons during reactions

They Are Useful

• In naming compounds, in writing formulas, in balancing chemical equations, and in studying certain types of chemical reactions (Ch. 19).

Assigning Oxidation Numbers:General Rule

• Shared electrons are assumed to belong to the more electronegative atom in each bond.

Specific Rules For Assigning Oxidation Numbers

I. Atoms of a pure element have an oxidation number of zero.

Ex:

Na

O2

S8

II. Electronegativity

• The more electronegative element in a binary compound is assigned the number equal to the negative charge it would have as an anion, the less electronegative one is positive (as if it were a cation).

III. Fluorine

• Assigned a value of

-1 in all compounds because it is the most electronegative element.

IV. Oxygen• Assigned a number of -2

in almost all compounds. Exceptions:

In peroxides like H2O2

(when it is -1)

In compounds with the halogens, like OF2 (when it is +2).

V. Hydrogen

• +1 in all compounds with elements that are more electronegative than it is.

• It is -1 when it is combined with metals.

Algebraic Sums

VI. In a neutral compound all oxidation numbers add up to zero.

VII. In a polyatomic ion the sum is equal to the charge of the ion.

VIII. Not For Covalently Bonded Atoms Only

• Monotomic ions have an oxidation number equal to the charge of the ion:

Na+ = +1

Cl − = -1

Because of Rules I-VIII

• It is often possible to assign oxidation numbers when they are not known.

Examples

• Assign oxidation numbers to each atom in the following compounds or ions:

• UF6 U=+6; F=-1

• H2SO4 H=+1; S=+6; O=-2

• ClO3- Cl=+5; O=-2

More Examples

• Assign oxidation numbers to each atom in the following compounds or ions:

• PCl3P=+3; Cl=-1

• PCl5P=+5; Cl=-1

• KH K=+1; H=-1• P4O10 P=+5; O=-2

Using Oxidation Numbers For Formulas and Names

• Tables on p. 205, p. 219, and in the appendix (p. 903), show that both metals and nonmetals can have more than one oxidation number.

Fe = +2 or +3

Carbon -4, +2, +4

Nitrogen -3, +3, +5

Phosphorus -3, +3, +5

Sulfur -2, +4, +6

Chlorine, Bromine, Iodine -1, +1, +3, +5, +7

The Stock System of Nomenclature

• In 7.1 we showed how Roman numerals show the charges of cations. But in reality the system was based on oxidation numbers:

FeO = Fe2+ + O2−

Fe2O3 = 2Fe3+ + 3O2−

Iron (II) oxide

Iron (III) oxide

Used As Alternative To Prefix System

• The international body that governs nomenclature endorses use of the Stock system as being more practical for complicated compounds.

Name the following compounds using the Stock system

• PCl3phosphorus (III) chloride

• PCl5phosphorus (V) chloride

• N2Onitrogen (I) oxide