Topic 2.6 Redox
Oxidation number
Oxidation number is the number of electrons an atom has gained or lost control of as a result of its bonding.
This definition is necessarily rather vague. What is meant by ‘gained or lost control of’? In the case of ions
comprising a single atom, the meaning is clear. When sodium atoms react, they lose one electron per atom
+
to form sodium ions (Na ) so the oxidation number is +1 (note that charge is written the opposite way as 1+).
Likewise, when fluorine forms ionic compounds each fluorine atom gains one electron to form fluoride ions
-
(F ) so the oxidation number is -1. However, in compounds with covalent bonding, the situation is less clear
and the assignment of oxidation number is based on a series of rules. These rules are based on the concept
of electronegativity, which is a measure of the tendency of an atom to draw the electrons in a covalent bond
to itself. Electronegativity will be covered in more detail later in the course but for now it is enough to know
that electronegativity increases in the order:
reactive metals < unreactive metals < unreactive non-metals < reactive non-metals
In other words, electronegativity increases as you go from the bottom left to the top right of the periodic table.
The rules for assigning oxidation numbers effectively consider the electrons in a covalent bond to belong to
the more electronegative element.
Although this may seem rather arbitrary, the concept of oxidation number is extremely useful when
considering the chemical changes involved in redox reactions.
The rules for assigning oxidation numbers to atoms in a chemical species are as follows:
1 The oxidation number of an element in its elemental state always equals zero,
e.g. oxygen gas, O2 : O=0
2 Some elements always have the same oxidation number when combined with other elements
Group 1 elements are always +1
e.g. in NaCl the Na is +1 (and Cl is -1)
Group 2 elements are always +2
e.g. in MgO the Mg is +2 (and O is -2)
Group 3 elements are usually +3
e.g. in Al2Cl6 the Al is +3
Fluorine is always -1
Oxygen is usually -2 (except when bonded to F or another O)
e.g. in MgO the O is -2; in OF2 the O is +2; and in peroxides, e.g. H 2 O 2 , the O is -1
Chlorine is usually -1 (except when bonded to F or O)
e.g. in NaCl the Cl is -1; in ClF3 the Cl is +3; and in NaClO the Cl is +1
Hydrogen is usually +1 (except when it is combined solely with a less electronegative metal)
e.g. in HCl the H is +1; but in NaH the H is -1
3 The sum of the oxidation numbers in a compound always equals zero,
e.g. magnesium chloride, MgCl2: Mg = +2; Cl = -1; + 2 + (2 x (-1)) = 0
4 The sum of the oxidation numbers in an ion always equals the charge on the ion,
-
e.g. chlorate(I) ion, ClO : Cl = +1; O = -2; +1 + (-2) = -1
01/06/2009
Oxidation number
Oxidation number is the number of electrons an atom has gained or lost control of as a result of its bonding.
This definition is necessarily rather vague. What is meant by ‘gained or lost control of’? In the case of ions
comprising a single atom, the meaning is clear. When sodium atoms react, they lose one electron per atom
+
to form sodium ions (Na ) so the oxidation number is +1 (note that charge is written the opposite way as 1+).
Likewise, when fluorine forms ionic compounds each fluorine atom gains one electron to form fluoride ions
-
(F ) so the oxidation number is -1. However, in compounds with covalent bonding, the situation is less clear
and the assignment of oxidation number is based on a series of rules. These rules are based on the concept
of electronegativity, which is a measure of the tendency of an atom to draw the electrons in a covalent bond
to itself. Electronegativity will be covered in more detail later in the course but for now it is enough to know
that electronegativity increases in the order:
reactive metals < unreactive metals < unreactive non-metals < reactive non-metals
In other words, electronegativity increases as you go from the bottom left to the top right of the periodic table.
The rules for assigning oxidation numbers effectively consider the electrons in a covalent bond to belong to
the more electronegative element.
Although this may seem rather arbitrary, the concept of oxidation number is extremely useful when
considering the chemical changes involved in redox reactions.
The rules for assigning oxidation numbers to atoms in a chemical species are as follows:
1 The oxidation number of an element in its elemental state always equals zero,
e.g. oxygen gas, O2 : O=0
2 Some elements always have the same oxidation number when combined with other elements
Group 1 elements are always +1
e.g. in NaCl the Na is +1 (and Cl is -1)
Group 2 elements are always +2
e.g. in MgO the Mg is +2 (and O is -2)
Group 3 elements are usually +3
e.g. in Al2Cl6 the Al is +3
Fluorine is always -1
Oxygen is usually -2 (except when bonded to F or another O)
e.g. in MgO the O is -2; in OF2 the O is +2; and in peroxides, e.g. H 2 O 2 , the O is -1
Chlorine is usually -1 (except when bonded to F or O)
e.g. in NaCl the Cl is -1; in ClF3 the Cl is +3; and in NaClO the Cl is +1
Hydrogen is usually +1 (except when it is combined solely with a less electronegative metal)
e.g. in HCl the H is +1; but in NaH the H is -1
3 The sum of the oxidation numbers in a compound always equals zero,
e.g. magnesium chloride, MgCl2: Mg = +2; Cl = -1; + 2 + (2 x (-1)) = 0
4 The sum of the oxidation numbers in an ion always equals the charge on the ion,
-
e.g. chlorate(I) ion, ClO : Cl = +1; O = -2; +1 + (-2) = -1
01/06/2009
