Redox I
Oxidation number
An oxidation number is a value assigned to an atom either alone or as part of a
compound to represent the number of electrons that the atom has either lost or
gained.
The oxidation state of the following are always fixed when they are in a compound:
Group 1 metals always have an oxidation number of +1.
Group 2 metals always have an oxidation number of +2.
Fluorine always has an oxidation number of -1.
Hydrogen always has an oxidation number of +1, expect for when it is
combined with a metal, where it has an oxidation number of -1.
Oxygen always has an oxidation number of -2, expect for when it is part of a
peroxide compound or combined with fluorine.
The sum of the oxidation numbers of all atoms in a neutrally charged compound is 0.
The sum of the oxidation numbers of all atoms in a polyatomic ion is the charge of
that ion.
In a single atom, the oxidation number is always 0. When two atoms of the same
element are chemically bonded together (e.g. Cl2), the oxidation number of each
atom is 0.
In a simple ion, the oxidation state of the atom in the ion is the charge of the ion.
Oxidation (loss of electrons) of an atom causes the atom’s oxidation number to
increase.
Reduction (gain of electrons) of an atom causes the atom’s oxidation number to
decrease.
Metals generally form positively charged ions by losing electrons, causing their
oxidation number to increase.
Non-metals generally form negatively charged ions by gaining electrons, causing
their oxidation number to decrease.
The name of a compound or ion may require an oxidation number to be indicated as
a Roman numeral. For example, manganese can take a variety of oxidation states.
In MnO4-, it has an oxidation state of +7, and so the ion is referred to as manganate
(VII). Another example is chromium in Cr2O72-, which has an oxidation state of +6,
and so the ion is called dichromate (VI). Iron can also have different oxidation states,
as indicated in FeCl3 – iron (III) chloride, and by FeSO4 – iron (II) sulfate.
Oxidation and Reduction
Oxidising agents oxidise other species, whilst being reduced themselves (gaining
electrons).
Oxidation number
An oxidation number is a value assigned to an atom either alone or as part of a
compound to represent the number of electrons that the atom has either lost or
gained.
The oxidation state of the following are always fixed when they are in a compound:
Group 1 metals always have an oxidation number of +1.
Group 2 metals always have an oxidation number of +2.
Fluorine always has an oxidation number of -1.
Hydrogen always has an oxidation number of +1, expect for when it is
combined with a metal, where it has an oxidation number of -1.
Oxygen always has an oxidation number of -2, expect for when it is part of a
peroxide compound or combined with fluorine.
The sum of the oxidation numbers of all atoms in a neutrally charged compound is 0.
The sum of the oxidation numbers of all atoms in a polyatomic ion is the charge of
that ion.
In a single atom, the oxidation number is always 0. When two atoms of the same
element are chemically bonded together (e.g. Cl2), the oxidation number of each
atom is 0.
In a simple ion, the oxidation state of the atom in the ion is the charge of the ion.
Oxidation (loss of electrons) of an atom causes the atom’s oxidation number to
increase.
Reduction (gain of electrons) of an atom causes the atom’s oxidation number to
decrease.
Metals generally form positively charged ions by losing electrons, causing their
oxidation number to increase.
Non-metals generally form negatively charged ions by gaining electrons, causing
their oxidation number to decrease.
The name of a compound or ion may require an oxidation number to be indicated as
a Roman numeral. For example, manganese can take a variety of oxidation states.
In MnO4-, it has an oxidation state of +7, and so the ion is referred to as manganate
(VII). Another example is chromium in Cr2O72-, which has an oxidation state of +6,
and so the ion is called dichromate (VI). Iron can also have different oxidation states,
as indicated in FeCl3 – iron (III) chloride, and by FeSO4 – iron (II) sulfate.
Oxidation and Reduction
Oxidising agents oxidise other species, whilst being reduced themselves (gaining
electrons).
