oxidation, reduction and redox reactions - chemistry

In terms of electrons, what does oxidation and reduction mean?

Oxidation = the loss of electrons

Reduction = the gain of electrons

How do we know if a reaction is a redox reaction?

A reaction is a redox reaction when both reduction and oxidation occur at the same time

If an elements oxidation state has increased (more positive), what does this mean?

It has been oxidised.

If an elements oxidation state has decreased (more negative), what does this mean?

It has been reduced.

What does the term reducing agent mean?

Reducing agents donate electrons (lose) electrons and are oxidised themselves.

A reducing agent causes another substance to be reduced (causing another substance to gain electrons).

State the oxidation numbers/states of uncombined elements, ions, group 1 elements, group 2 elements, aluminium, hydrogen, chlorine, fluorine and oxygen.

1) Uncombined elements - always 0

2) ions - oxidation state is the same as the charge on the ion.

3) group 1 - always +1

4) group 2 - always +2

5) aluminium - always +3

6) hydrogen - +1 except in hydrides where it is -1

7) chlorine - -1 (except if in a compound with F and O - it would have a positive value)

8) Fluorine - always -1

9) Oygen - -2 (except in peroxides where it is -1).

State the oxidation numbers/states of uncombined elements, ions, group 1 elements, group 2 elements, aluminium, hydrogen, chlorine, fluorine and oxygen.

1) Uncombined elements - always 0

2) ions - oxidation state is the same as the charge on the ion.

3) group 1 - always +1

4) group 2 - always +2

5) aluminium - always +3

6) hydrogen - +1 except in hydrides where it is -1

7) chlorine - -1 (except if in a compound with F and O - it would have a positive value)

8) Fluorine - always -1

9) Oygen - -2 (except in peroxides where it is -1).

What are the oxidation states of transposition metals?

They have variable oxidation states.

What are the rules for balancing half-equations?

1) Write down the species before and after the reaction.

2) Balance any atoms apart from oxygen and hydrogen.

3) Balance any oxygens with H2O

4) Balance any hydrogens with H+ ions.

5) Balance charges with electrons.

Write the stages of balancing the half equation involving the conversion of Fe2+ to Fe3+.

The Fe ions are already balanced (one on either side) so no need to balance them.

There are no oxygens or hydrogens so there is no need to use H+ and H2O to balance the hydrogens and the oxygens.

However the charges are different on either side so an electron must be added to the Fe 3+ side to make both the charges +2.

Write the stages of balancing the half equation of Mno4- to Mn2+.

The atoms (other than hydrogen and oxygen) are balanced.

The 4 oxygens on the left side need to be balanced out by adding 4 H2O molecules on the right hand side.

As there are now 8 hydrogens on the right hand side, we need to add 8H+ ions to the left hand side.

Finally, balancing the charges - the left hand side has a charge of +7 and the right hand side has a charge of +2 therefore 5 electrons (e-) are added to the left hand side so both sides have equal charge.

What can happen with the past two had equations?

You can combine two Hal equations to form a full ionic equation.

Before you combine the two half equations, what needs to happen?

The electrons in both half equations need to be balanced, therefore the whole of the Fe equation needs to be multiplied by 5 as there is only 1 electron in that equation but 5 in the Mn equation.

You then cancel out the electrons and combine the two half equations- there should be no electrons in the final equation.

What is the full ionic equation also known as?

The full ionic equation shows both reduction and oxidation and is therefore a redox reaction.