Home Y10 Y11 AS Units A2 Units

# Foundation - Redox Equations to Balance

## Redox Equations - Notes on Assigning Oxidation Numbers

1. Uncombined elements are given the oxidation number zero.

e.g. the oxidation numbers of chlorine in chlorine gas, Cl2, and sodium in sodium metal, Na, are both 0.

2. For a single (i.e. monatomic) ion, the oxidation number of the element is the charge on the ion.

e.g. the oxidation number of chlorine in chloride ions, Cl-, is -1. The oxidation number of sodium in sodium ions, Na+, is +1.

N.B.: The sign should be shown in front of the oxidation number.

3. The sum of the oxidation numbers of the ions in the formula of a compound is zero.

e.g. in sodium chloride, the oxidation number of sodium is +1, of chlorine is -1, and therefore the sum = (+1) + (-1) = 0.

N.B.: This is a very useful rule for working out unknown oxidation numbers.

4. Certain elements almost always have the same oxidation number in their compounds (these, if known, can be used to deduce unknown oxidation numbers for those elements whose oxidation numbers values vary widely).

e.g.

F : always -1 in all its compounds.

O : normally -2 in its compounds except in compounds with fluorine or in peroxides, then it is -1 or even -½ (in superoxides).

H : normally +1 in its compounds, except in a compound with a metal i.e. a metal hydride, e.g. NaH, then it is -1.

Metals which only form one type of ion (or whose valency does not vary) have an oxidation number equal to the charge on the ion (or the valency of the metal).

e.g. sodium is always +1 in its compounds, calcium always +2, potassium always +1, aluminium always +3, etc.

5. For a polyatomic ion (i.e. an ion containing more than one atom) e.g. the sulphate ion, SO42-, the sum of all the oxidation numbers of the elements in the ion is equal to the charge on the ion.

## Redox Equations - Straightforward Examples

 Zn(s) + Fe3+(aq) → Zn2+(aq) + Fe2+(aq)

 Al(s) + Ag+(aq) → Al3+(aq) + Ag(s)

 HNO3(aq) + Sn(s) → NO2(g) + SnO2(s) + H2O(l)

 HNO3(aq) + S(s) → NO2(g) + H2SO4(aq) + H2O(l)

 MnO4-(aq) + H2S(g) + H+(aq) → S(s) + Mn2+(aq) + H2O(l)

 MnO4-(aq) + SO32-(aq) + H+(aq) → SO42-(aq) + Mn2+(aq) + H2O(l)

 MnO4-(aq) + Zn(s) + H+(aq) → Zn2+(aq) + Mn2+(aq) + H2O(l)

## Redox Equations - Harder Examples

 I-(aq) + Cl2(g) → I2(s) + Cl-(aq)

 H+(aq) + Al(s) → H2(g) + Al3+(aq)

 NH3(g) + O2(g) rarr; NO(g) + H2O(g)

 Cr2O72-(aq) + H2S(g) + H+(aq) → S(s) + Cr3+(aq) + H2O(l)

 Cr2O72-(aq) + SO32-(aq) + H+(aq) → SO42-(aq) + Cr3+(aq) + H2O(l)

 Cr2O72-(aq) + Fe2+(aq) + H+(aq) → Fe3+(aq) + Cr3+(aq) + H2O(l)

 MnO4-(aq) + C2O42-(aq) + H+(aq) → CO2(g) + Mn2+(aq) + H2O(l)

 MnO4-(aq) + C2H5OH(aq) + H+(aq) → C2H4O(g) + Mn2+(aq) + H2O(l)

 C2H5OH(aq) + Cr2O72-(aq) + H+(aq) → C2H4O(g) + Cr3+(aq) + H2O(l)

 C2H4O(aq) + Cr2O72-(aq) + H+(aq) → C2H4O2(aq) + Cr3+(aq) + H2O(l)

## Redox Equations - Disproportionation Examples

 Cu+(aq) → Cu(s) + Cu2+(aq)

 Cl2(aq) + -OH(aq) → Cl-(aq) + ClO-(aq) + H2O(l)

 Cl2(aq) + -OH(aq) → Cl-(aq) + ClO3-(aq) + H2O(l)

 HNO2(aq) → HNO3(aq) + NO(g) + H2O(l)

 MnO42-(aq) + H2O(l) → MnO4-(aq) + MnO2(s) + -OH(aq)

 Fe3+(aq) + Fe(s) → Fe2+(aq)

 IO3-(aq) + I-(aq) + H+(aq) → I2(s) + H2O(l)

written by Dr Richard Clarkson : © Saturday, 1 November 1997

Updated : Saturday, 13th October, 2012

mail to: chemistryrules

Created with the aid of,