1. Oxidation-Reduction Reactions
Oxidation Numbers
Oxidation Numbers and Nomenclature
Identifying Oxidation-Reduction Reactions
Writing Equations for Oxidation-Reduction Reactions
Oxidation-Reduction Titrations
Oxidation by Oxygen

2. Oxidation Numbers Definition:
An oxidation number is a number that reflects the electrons gained, lost, or shared when an element reacts.
Remember that when you lose electrons you become more positive and when you gain electrons you become more negative.

3. Rules for Assigning Oxidation Numbers
The oxidation # of any uncombined element is zero.
The oxidation # of a monoatomic ion equals its’ charge.
The more electrongative element in a binary compound is assigned the number equal to the charge it would have if it were an ion.
The oxidation number of fluorine is always -1.
Oxygen has an oxidation number of -2.
Hydrogen has an oxidation number of +1.
In compounds, Group 1 and 2 elements and aluminum have oxidation numbers of +1, +2, and +3 respectively.
The sum of the oxidation numbers of atoms in a compound is zero.
The sum of the oxidation # in a polyattomic ion is equal to the charge of the ion.

4. Oxidation number or state
When dealing with simple ions, this is easy to determine. It is simply the charge on the ion.
Examples
Group IA (1) +1
Group IIA (2) +2
Group VIIA (17) -1
Oxygen -2 usually
Hydrogen +1 if bonded to nonmetal
Hydrogen -1 if bonded to metal

5. Oxidation number
For elements in their elemental state, the oxidation number is also pretty straightforward.
Since all of the atoms are the same, the electrons are shared equally so the oxidation number is zero.
Examples
The atoms in N2, Na, P4, H2 and O2 all have oxidation numbers of zero.

6. Oxidation numbers
With polar covalent bonds electrons are shared but not equally.
For electrons that are shared in these compounds, we assign the shared electrons to the most electronegative element.
We are just acting as though the electronegativity difference was large enough for the transfer of electrons to occur.

7. Example Assign the oxidation states for all elements in water.
The electronegativities are:
H = 2.2, O = 3.5
The electrons from both hydrogen are assigned to the oxygen.
Oxidation numbers: O = -2
H = +1

8. Oxidation numbers
Many elements have more than one possible oxidation number.
Often, it is possible to determine the oxidation number of those elements in a compound simply by looking at what you do know.
Follow the previous rules and then assign an oxidation number that insures that the overall compound has no net charge.

9. HNO3 Example Find the oxidation state for all elements in:
Hydrogen we know - it must be +1
Oxygen should be -2 in this case.
What about nitrogen?

10. Example See what you do know and find the difference. We know that H
HNO3
We know that H
is assigned a
value of +1
Oxygen is -2
and we have 3
of them = -6
OK, so what's left over??
(+1) + (-6) + charge on nitrogen = 0
Nitrogen must have a value of +5

11. Determining Oxidation Numbers
SO2 NO3-
PCl5 SO42-
H2O NH4+
H2SO3 MnO4-

12. Common oxidation numbersLi +1 Na Cs Rb K Fr Ba +2 Be Mg Sr Ca Ra B +3 C +4 -2 -4 N
+5 +4
+3 +2
+1 -3 O -1 -2 F -1 Ne H +1 He Al +3 Si +4 -4 P +5 +3 -3 S
+6 +4 +2 -2 Cl
+7 +5
+3 +1 -1 Ar Sc 3+ Ti +4 +3 +2 V
+5 +4 +3 +2 Cr +6 +3 +2 Mn
+7 +6
+4 +3 +2 Fe +3 +2 Co +3 +2 Ni +2 Cu +2 +1 Zn +2 Ga +3 Ge +4 -4 As 5+ 3+ 3- Se 6+ 4+ 2- Br +5 +1 -1 Kr +4 +2 Y +3 Zr +4 Nb +5 +4 Mo +6 +4 +3 Tc +7 +6 +4 Ru
+8 +6 +4 +3 Rh +4 +3 +2 Pd +4 +2 Ag +1 Cd +2 In +3 Sn +4 +2 Sb +5 +3 -3 Te +6 +4 -2 I
+7 +5 +1 -1 Xe +6 +4 +2 Lu +3 Hf +4 Ta +5 W +6 +4 Re +7 +6 +4 Os +8 +6 Ir +4 +3 Pt +4 +2 Au +3 +1 Hg +2 +1 Tl +3 +1 Pb +4 +2 Bi +5 +3 Po +2 At -1 Rn Lr +3
Common oxidation numbers

13. Oxidation numbers and the periodic table
Some observed trends in compounds.
Metals have positive oxidation numbers.
Transition metals typically have more than one oxidation number.
Nonmetals and semimetals have both positive and negative oxidation numbers.
No element exists in a compound with an oxidation number greater than +8.
The most negative oxidation numbers equals 8 - the group number

14. Oxidation number and nomenclature
Stock system
For metals with several possible oxidation numbers, use Roman numeral in the name.
FeSO4 iron(II) sulfate
Fe2(SO4)3 iron (III) sulfate
Cu2O copper(I) oxide
CuO copper(II) oxide
PbCl2 lead(II) chloride
PbCl4 lead(IV) chloride

15. Identifying oxidation-reduction reactions.
Oxidation-Reduction - REDOX
A chemical reaction where there is a net change in the oxidation number of one or more species.
Both an oxidation and a reduction must occur during the reaction.
Mg (s) + Cl2 (g) MgCl2 (s)
Here the oxidation number of Mg has changed from
zero to +2. Cl has changed from zero to -1.

16. REDOX reactions Oxidation An increase in oxidation number. Reduction
A decrease in oxidation number.
If the oxidation number of any element changes in the course of a reaction, the reaction is oxidation-reduction.
Example.
2 Fe(NO3)3 (aq) + Zn(s) Fe(NO3)2 (aq) + Zn(NO3)2 (aq)

17. Redox Reaction Fe(s) + O2(g)  Fe2O3(s) Iron is oxidized (0 to +3)
and O2 is reduced (0 to -2)

18. Half Reactions Fe  Fe3+ + 3e- (oxidation)
O e-  2 O2- (reduction)

19. Half reactions Example.
Half-reactions can be of the ‘net ionic’ form. Balance the follow
Fe3+ + Zn (s) Fe2+ + Zn2+
2 ( Fe3+ + e Fe2+) (reduction)
Zn(s) Zn2+ + 2e- (oxidation)
2Fe3+ + Zn (s) Fe2+ + Zn2+

20. Example Fe3+ is reduced to Fe2+ Zn is oxidized to Zn2++3 +2 +2
2Fe(NO3)3 (aq) + Zn(s) Fe(NO3)2 (aq) Zn(NO3)2 (aq)
Fe3+ is reduced to Fe2+
Zn is oxidized to Zn2+
NO3- is a spectator ion.

21. Zinc - Copper Voltaic Cell

23. A Zinc-Copper Voltaic Cell
Zn electrode
Anode (-)
Oxidation
Half Reaction:
Zn  Zn e-
Cu electrode
Cathode (+)
Reduction
Half Reaction:
Cu e-  Cu

24. Voltaic Cell cont…
In a voltaic cell the redox reaction is spontaneous and produces an electric current. This current can be measured using a voltmeter.

26. Standard Electrode Potentials
A standard electrode potential, Eo, is based on the tendency for reduction to occur at the electrode.
The cell voltage, called the standard cell potential (Eocell), is the difference between the standard potential of the cathode and that of the anode.
Eocell = Eo (cathode) – Eo (anode)

27. Standard Electrode Potentials

28. Balancing REDOX equations
Many REDOX equations can be balanced by inspection.
H2S (g) + H2O2 (aq) S (s) + 2 H2O (l)
However, others are more difficult.
2KMnO4 (aq) + H2O2 (l) + 3H2SO4 (aq)
2MnSO4 (aq) + K2SO4 (aq) + 3O2 (g) + 4H2O (l)

29. Balancing REDOX equations
Half-Reaction method.
With this approach, the reaction is broken into two parts.
Oxidation half-reaction. The portion of the reaction where electrons are lost.
A An+ + ne-
Reduction half-reaction. The portion of the reaction where electrons are gained.
me- + B Bm-

30. Balancing REDOX equations
The goal is then to make sure that the same number of electrons are being produced and consumed.
(m) ( A An+ + ne- )
(n) (me- + B Bm )
nB + mA mAn+ + nBm+
When properly balanced, the electrons will cancel out.

31. Half reactions Another Example
Determine the balanced equation for the reaction of Fe2+ with Cr2O72- in an acidic solution.
Fe2+ + Cr2O Fe3+ + Cr3+
The two half-reactions would be:
Fe2+ Fe3+
Cr2O72- Cr3+ H+

32. Half reactions
First, balance each half-reaction for all elements except hydrogen and oxygen.
Fe2+ Fe3+
Cr2O Cr3+
Next, balance each half-reaction with respect to oxygen by adding an appropriate number of H2O.
Fe2+ Fe3+
Cr2O Cr3+ + 7H2O

33. Half reactions
Remember that this reaction occurs in an acid solution so we can add H+ as needed.
Fe2+ Fe3+
14H+ + Cr2O Cr3+ + 7H2O
Now we need to know how many electrons are produced or consumed and place them in our half-reactions.
For iron, one e- is produced.
For dichromate, six e- are consumed.

34. Half-reactions Fe2+ Fe3+ + e- 6e- + 14 H+ + Cr2O72- 2Cr3+ + 7H2O
We need the same number of electrons produced and consumed so:
6Fe Fe3+ + 6e-
6e- + 14H+ + Cr2O Cr3+ + 7H2O
As our final step, we need to combine the half-reactions and cancel out the electrons.

35. Half-reactions 6Fe2+ + 14H+ + Cr2O72- 6Fe3+ + 2Cr3+ + 7H2O
In this reaction, Fe2+ is oxidized and the dichromate ion is reduced.
This reaction is used for the determination of iron by titration.

36. Disproportionation reactions
In some reactions, the same species is both oxidized and reduced.
Examples
2H2O2 (l) H2O (l) + O2 (g)
3Br2 (aq) +6OH- (aq) BrO3-(aq) +5Br-(aq) +3H2O(l)
For this to occur, the species must be in an intermediate oxidation state. Both a higher and lower oxidation state must exit.

37. Oxidation-reduction titrations
REDOX reactions can also serve as the basis for titrations.
For example, we can determine the amount of iron in an ore by titration.
Initially, we must dissolve the sample. This results in both iron(II) and iron(III) being produced in solution.
The first step is to get all of the iron into one oxidation state.

38. Sample preparation
One option is to use a reductor. You slowly wash your sample through the column with water.
Jones Reductor
Zn(Hg) Zn2+ + Hg(l) + 2e
An amalgam is used to prevent
Zn + 2H+ Zn2+ + H2(g)
For iron, we get
Fe2+ + Fe Fe2+
reductor

39. Titrants
Now that all of our iron is in a single oxidation state, we’re ready to do a titration.
We need an oxidizing agent to convert all the iron from Fe2+ to Fe3+.
Primary standard
A material that is available in pure form.
Ideally, it should be something that we can directly weigh out, dissolve and then titrate with.

40. Common titrants Oxidizing titrants Dichromate - Cr2O72-
The potassium salt is a primary standard material.
Very stable solutions. If air is kept out, it can last for years.
It is a very strong oxidizing agent.
Need an indicator such as diphenylamine sulfonic acid.

41. Common titrants Oxidizing titrants Permanganate - MnO4-
The potassium salt is the most commonly used. It is not a primary standard.
Solutions must be standardized - typically use Na2C2O4 ( a primary standard material.)
Reagent slowly degrades and MnO2 must be removed
No indicator is needed - excess reagent produces a pink solution.

42. Common titrants
The standardization of MnO42- with oxalate involves the following reaction.
2MnO42-(aq) + 5C2O42-(aq) + 8H+(aq)
2Mn2+(aq) + 10CO2 (g) + 8H2O (l)
Since MnO42- in an intense magenta color and Mn2+ is a faint green, detecting the endpoint for the titration is easy.

43. Oxidation by oxygen
Oxygen is not the strongest of oxidizing agents but it is about 19% of our atmosphere.
It is able to react with all other elements except: noble gases
halogens
noble metals like gold
One very common reaction for oxygen is:
Combustion - rapid oxidation accompanied by heat and usually light.

44. Combustion Examples CH4(g) + 2O2(g) CO2(g) + 2H2O(g)
S(s) + O2(g) SO2(g)
N2(g) + O2(g) NO(g)
2NO(g) + O2 (g) 2 NO2(g)
Note. All of these oxides contribute to air pollution. CO2 contributes to the “greenhouse effect.”