1. Redox – Part 2 a. demonstrate an understanding of:
iv. how oxidation number is a useful concept in terms of the classification of reactions as redox and as disproportionation
b. write ionic half-equations and use them to construct full ionic equations.

2. Task: Match up the formulas with their
correct names

3. Naming Compounds
Some elements can have more than one oxidation number.
So to avoid confusion when naming the compounds they form, its helpful to use these numbers in their names.

4. Naming Compounds
What are the missing oxidation numbers here for Manganese and Chromium, respectively? ? ?

5. Task Answer – sometimes, instead of using the systematic name, scientists still use the traditional names for some substances.

6. Redox and Oxidation Numbers
An increase in oxidation number means oxidation has occurred.
While a decrease in oxidation number means reduction has occurred.
E.g.
What would be
the oxidation
numbers for
each of the
species in this
reaction? +1 -2 +1 -1
Therefore,
chlorine has been reduced, and
sulphur has been oxidised.

7. Redox and Oxidation Numbers
Another example is the Thermite Reaction (click here to see it in action!!)
But what is oxidised and what is reduced? +3 -2 +3 -2
Therefore,
iron has been reduced, and
aluminium has been oxidised.

8. Disproportionation
Is a type of reaction in which a substances both oxidised and reduced in the same reaction. E.g. the breakdown of hydrogen peroxide:

9. Disproportionation – another example
The reaction of copper(I) oxide with dilute sulphuric acid:
Study the equation carefully – what disproportionates?
The answer is copper (Cu) which gets oxidised from +1 to +2 and also reduced from +1 to zero, in the same reaction.
Other examples of such reactions include reactions of chlorine with water, and chlorine with hot sodium hydroxide solution.

10. Questions

11. Answers

12. Oxidation States Some elements show a range of oxidation states.
One colourful example is the reduction of vanadium (V) to vanadium (II) through sucessive oxidation number.

13. Vanadium – its different oxidation states
Ammonium Vanadate (V) (white solid) can be added to dilute HCl to give the orange Dioxovanadium (V) ion:
VO3- + 2H+
VO2+ + H2O +5 +5
If granulated zinc is now added it reduces the vanadium over a period of several minutes and gives several colour changes:
Blue [VO(H2O)5]2+ Vanadium (IV)
Green [VCl2(H2O)4] + Vanadium (III)
Violet [V(H2O)6]2+ Vanadium (II)

14. Common Oxidising Agents
What is an oxidising agent?
A species that reacts by oxidising something else and getting reduced itself.
So, as we will see, electrons in the half equations are always on the left-hand side of the equation, as oxidising agents gain electrons.

15. More Oxidising Agents

16. More Oxidising Agents

17. Common Reducing Agents
What is an reducing agent?
A species that reacts by reducing something else and getting oxidised itself.
So, this time, electrons in the half equations are always on the right-hand side of the equation, as reducing agents loose electrons.

18. More Reducing Agents

19. More Reducing Agents

20. Writing balanced equations from ionic half equations - Worked Example…
Firstly, write down the
half equations you
need.

21. Worked Example…..cont…. X5

22. Worked Example…..cont….
Notice, that in adding the two equations together
The electrons are cancelled out and not
Included in the final equation

23. Again, start with the half equations
Worked Example 2…
Again, start with the half equations
Have a go
before you turn to
the next slide for
the answer

25. Again, start with the half equations
Worked Example 3…
Again, start with the half equations

26. ? X Or

28. Ionic to Molecular Equation
To turn an ionic equation to a molecular equation, you would simply add the spectator ions.
Eg. Reaction of potassium manganate (VII) solution with iron (II) sulphate solution in the presence of dilute sulphuric acid:
Ionic equation – which you work out
Add in spectator ions by using clues
In question
molecular equation

29. Questions

30. Answers