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Slide 1

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Chapter 10
Redox reactions…read more

Slide 2

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Rules for assigning
oxidation numbers
1. The oxidation number of an uncombined element is 0.
2. A simple monatomic ion has an oxidation number equal
to it's charge
3. The oxidation number of group 1 metals is always +1
and of group 2 metals is always +2 *without exception*
4. Fluorine is always -1
hydrogen is always +1 except in metal hydrides
oxygen is always -2 except in F2O and peroxides (e.g.
5. The sum of the oxidation numbers in a molecule adds
up to 0
A polyatomic ion adds up to the charge on the ion.
The oxidation number is the charge on an atom of an
element in a compound calculated assuming that all
the atoms in the compound are simple monatomic
The more electronegative element is given an oxidation
number of -1 per bond
When an element is oxidised its oxidation number
Electronegativity increases across a period and decreases
down a group.
Oxidation numbers should be whole numbers and are
always either + or -.…read more

Slide 3

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Balancing Redox half
1. Work out formulae of the species before
and after the change ­ balance if required
2. Work out oxidation state of the element
before and after the change
3. Add electrons to one side of the equation
so that the oxidation states balance
4. If the charges on the species (ions and
electrons) on either side of the equation
do not balance then add sufficient H+
ions to one of the sides to balance the
5. If the equation still doesn't balance, add
sufficient water molecules to one side…read more

Slide 4

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Combining half equations
A combination of two ionic half equations, one
involving oxidation and the other reduction
1. Write out two half equations
2. Multiply the equations so that the number of
electrons in each is the same
3. Add the two equations and cancel out the
electrons on either side
4. If necessary, cancel any other species which
appear on both sides…read more


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