1. Oxidation and Reduction
Definitions of oxidation and reduction
Oxidation numbers
Redox equations

2. Oxidation - reduction Oxidation is loss of electrons
Reduction is gain of electrons
Oxidation is always accompanied by reduction
The total number of electrons is kept constant
Oxidizing agents oxidize and are themselves reduced
Reducing agents reduce and are themselves oxidized

3. Follow the electrons

4. Oxidation numbers
Oxidation number is the number of electrons gained or lost by the element in making a compound
Metals are typically considered more 'cation-like' and would possess positive oxidation numbers, while nonmetals are considered more 'anion-like' and would possess negative oxidation numbers.

5. Predicting oxidation numbers
Oxidation number of atoms in element is zero in all cases
Oxidation number of element in monatomic ion is equal to the charge
sum of the oxidation numbers in a compound is zero
sum of oxidation numbers in polyatomic ion is equal to the charge
F has oxidation number –1
H has oxidn no. +1; except in metal hydrides where it is –1
Oxygen is usually –2. Except:
O is –1 in hydrogen peroxide, and other peroxides
O is –1/2 in superoxides KO2
In OF2 O is +2

6. Position of element in periodic table determines oxidation number
G1A is +1
G2A is +2
G3A is +3 (some rare exceptions)
G5A are –3 in compounds with metals, H or with NH4+. Exceptions are in compounds to the right; in which case use rules 3 and 4.
G6A below O are –2 in binary compounds with metals, H or NH4+. When they are combined with O or with a lighter halogen, use rules 3 and 4.
G7A elements are –1 in binary compounds with metals, H or NH4+ or with a heavier halogen. When combined with O or a lighter halogen, use rules 3 and 4.

7. Identifying reagents
Those elements that tend to give up electrons (metals) are typically categorized as reducing agents and those that tend to accept electrons (nonmetals) are referred to as oxidizing agents.

8. Iron can reduce Cu2+ to Cu
The iron nail reduces the Cu2+ ions and becomes coated with metallic Cu. At the same time, the intensity of the blue color diminishes due to loss of Cu2+ ions from solution.

9. Any element can be both an oxidizer and reducer depending on relative positions in the activity series
Fe reduced Cu2+, but Cu can reduce Ag+ (lower activity
Fe2+ is reduced by Zn

10. More active metals are strongly reducing

11. Predicting results of displacement reactions
In this reaction the element metal A displaces the ion metal B from its compound
This will only occur if A lies above B in the activity series
Displacement reaction exercises

12. Nuggets of redox processes
Where there is oxidation there is always reduction
Oxidizing agent
Reducing agent
Is itself reduced
Is itself oxidized
Gains electrons
Loses electrons
Causes oxidation
Causes reduction

13. Identify redox by change in oxidation numbers: follow the flow of electrons
Reducing agent increases its oxidation number
Oxidizing agent decreases its oxidation number

14. In reaction with metals, nonmetals are always oxidizers
Reactions of elements are always redox
The nonmetal gains electrons, becomes a negative ion
The metal loses electrons, becomes a positive ion
Identification is harder when there are no elements involved: oxidation numbers must be used

15. Balancing redox equations: systematic methods
Oxidation number method – tracking changes in the oxidation numbers
Half-reaction method – tracking changes in the flow of electrons
Same principles, different emphasis
Use is a matter of choice, but familiarity with both is important

16. Oxidation number method
What goes up must come down…
Sum of the changes in oxidation numbers in any process is zero

17. Six habits of the redox equation balancer

18. Walking through the steps
STEP 1: write the unbalanced net ionic equation
In an acid solution, permanganate is reduced by bromide ion to give Mn2+ ion and bromine

19. STEP 2: Balance the equation for elements other than O and H*
We need to double the bromide ions on the L.H.S. to balance the equation
*O and H can remain unbalanced because we will top up with water and hydronium ions later

20. STEP 3: Assign oxidation numbers
Use the rules of oxidation numbers
Element is zero
Monatomic ion: oxidation number is same as charge
Oxide is -2 +2 -1 +7

21. STEP 4: Identify oxidized and reduced
Mn is reduced from +7 to +2
Net gain of 5 electrons
Br is oxidized from -1 to 0
Net loss of 1 electron +2 -1 +7

22. STEP 5: Balance the oxidized and reduced species
For every Mn reduced (decrease in oxidation number of 5), need five Br- oxidized (increase in oxidation number of 1)
Equation becomes
Redox is now complete but material balance is not

23. STEP 6: Material balance with H2O and H+
Strategy: add H2O to the side that lacks for O and add H+ (the reaction is in acid solution) to the other side
In basic solution we add OH- and H2O instead of H2O and H+ respectively
Test equation for both atoms and charges

24. The Half-Reaction method
Any redox process can be written as the sum of two half reactions: one for the oxidation and one for the reduction

25. Six habits of the redox equation balancer

26. STEP 1: the unbalanced equation
Dichromate ion reacts with chloride ion to produce chlorine and chromium (III)

27. STEP 2: identify the oxidized and reduced and write the half reactions
Oxidation half-reaction
Reduction half-reaction

28. STEP 3: Balance the half reactions
Oxidation
Reduction

29. STEP 4: Material balance
As with the oxidation number method, add H2O to the side lacking O and add H+ to the other side (for reactions in acid solution)
Oxidation reaction – unchanged
Reduction reaction

30. STEP 5: Balance half-reactions for charge by addition of electrons
No explicit calculation of oxidation numbers is required; we balance the charges on both sides of each half-reaction

31. STEP 5 cont: Multiply by factors to balance total electrons
Overall change in electrons must be zero
Multiply the oxidation half reaction by 3

32. STEP 6: Add half reactions and eliminate common items+ =
Atoms and charges balance

33. Redox Titrations
Acid-base titration is used to determine an unknown concentration (either acid or base)
The endpoint is manifested in a color change (of an indicator) or by measuring pH
In redox titrations, the concentration of one of the reagents can be measured, provided there is a sharp distinction between the oxidized and reduced states

34. Using the roadmap Oxalic acid is oxidized by MnO4-
A known quantity of oxalic acid is used to determine the concentration of the MnO4-
MnO4- has an intense purple colour, whereas Mn2+ is almost colourless

35. Strategy A known amount of H2C2O4 is used
A solution of KMnO4 is titrated till the first purple colour – the endpoint. All the H2C2O4 is oxidized.
The equation gives the number of moles of MnO4-
The volume of solution yields the concentration

36. Naming Salts

37. We learned last time that acids form salts when they react
We learned last time that acids form salts when they react. But how do we name these salts? There are two parts to the name of a salt – one from the acid and one from the base/metal.

38. The first part of the salt name
The first part comes from the base/metal:
e.g.
magnesium carbonate  magnesium _____________
calcium hydroxide  calcium ________________
sodium oxide  sodium ______________
aluminium  aluminium _____________
ammonium carbonate  __________________

39. The second part of the salt name
The second part comes from the acid:
e.g.
hydrochloric acid  _____________ chloride
nitric acid  ________________ nitrate
sulphuric acid  ______________ sulphate
phosphoric acid  _____________ phosphate
ethanoic acid  _____________ ethanoate
carbonic acid  _____________ carbonate
etc.

40. Putting that together + + + + sodium hydroxide nitric acid sodium 
water
nitrate
this is the salt
magnesium oxide +
sulphuric
acid
magnesium  +
water
sulphate

41. Putting that together + + + + + hydrochloric acid iron iron  hydrogen
chloride
ammoniumcarbonate
ammonium carbonate
phosphoric
acid
ammonium +  +
water
phosphate +
carbon
dioxide

42. Now try the naming salts worksheet.