1. Volumetric Analysis Unit 3
Redox reactions
Volumetric Analysis Unit 3

2. Types of chemical reactions
Fall into several categories.
1) Precipitation reactions
The general form of the equation for such a reaction is: AB + CD  AD + CB
2) Acid-Base reactions. An acid substance react with a substance called a base. Such reactions involve the transfer of a proton (H+) between reactants.
3) Oxidation-reduction reactions. These involve the transfer of electrons between reactants.

3. Oxidation: Reduction:
Gain of oxygen Loss of oxygen
Loss of electrons Gain of electrons
Increase in
oxidation
number
Decrease in
oxidation
number

4. You can’t have one without the other!
Reduction (gaining electrons) can’t happen without an oxidation to provide the electrons.
You can’t have 2 oxidations or 2 reductions in the same equation. Reduction has to occur at the cost of oxidation
Oxidation and reduction always occur simultaneously

5. 2Mg(s) + O2(g)  2MgO(s) Mg  Mg2+ +2e- O2  O2- +2e-
Oxidised – loss of e-
O  O2-
Reduced – gain of e-
+2e-

6. Review of Oxidation numbers
The charge the atom would have in a molecule (or an ionic compound) if electrons were completely transferred.
Free elements (uncombined state) have an oxidation number of zero.
Na, Be, K, Pb, H2, O2, P4 = 0
In monatomic ions, the oxidation number is equal to the charge on the ion.
Li+, Li = +1; Fe3+, Fe = +3; O2-, O = -2
Oxidation numbers are hypothetical charges on atoms.

7. Assigning Oxidation Numbers
For ionic compounds, oxidation numbers can be assigned
using the expected charges from the periodic table.
The sum of the oxidation numbers of the atoms in the compound must equal 0.

8. Assigning Oxidation Numbers
For ionic compounds, oxidation numbers can be assigned
using the expected charges from the periodic table.
Alkali Metals = +1

9. Assigning Oxidation Numbers
For ionic compounds, oxidation numbers can be assigned
using the expected charges from the periodic table.
Alkaline Earth Metals = +2

10. Assigning Oxidation Numbers
For ionic compounds, oxidation numbers can be assigned
using the expected charges from the periodic table.
Group 13 Boron Group = +3

11. Assigning Oxidation Numbers
For ionic compounds, oxidation numbers can be assigned
using the expected charges from the periodic table.
Group 15 Nonmetals = -3

12. Assigning Oxidation Numbers
For ionic compounds, oxidation numbers can be assigned
using the expected charges from the periodic table.
Group 16 Nonmetals = -2

13. Assigning Oxidation Numbers
For ionic compounds, oxidation numbers can be assigned
using the expected charges from the periodic table.
Halogens = -1

14. Assigning Oxidation Numbers
For ionic compounds, oxidation numbers can be assigned
using the expected charges from the periodic table.
Transition Metals depend on anion

15. Ionic Compounds
Example: +2 -1
MgCl2

16. Ionic Compounds
Example: +1 -2
Na2O

17. Fe2O3 Ionic Compounds ?? -2 Example:
We don’t know iron’s oxidation number from the periodic table since it is a transition metal. ?? -2
Fe2O3

18. Fe2O3 Ionic Compounds +3 -2 Example:
But since we know the compound is neutral, the oxidation numbers must add up to zero. Therefore, Fe has a +3 oxidation number in this compound.
Example: +3 -2
Fe2O3

19. Cr2O72- Ionic Compounds ?? -2 Example:
We don’t know chromium’s oxidation number from the periodic table since it is a transition metal. ?? -2
Cr2O72-

20. Cr2O72- Ionic Compounds +6 -2 Example:
But since this ion has a charge of -2, the oxidation numbers must add up to negative two. Therefore the oxidation number of Cr is +6
Example: +6 -2
Cr2O72-

21. Covalent Compounds
Covalent compounds are made of two nonmetals, which from the periodic table are always expected to be negative

22. Covalent Compounds
But since covalent compounds are neutral species, it is not possible for every element to retain its negative oxidation number

23. Covalent Compounds
ONLY THE MORE ELECTRONEGATIVE ELEMENT keeps its negative oxidation number. Other nonmetals must adapt to keep the compound neutral

24. Electronegativity Trend
Increases

25. SO2 Covalent Compounds -2
Example: -2
SO2
Since oxygen is the more electronegative element, it will have its normal oxidation number.

26. SO2 Covalent Compounds +4 -2
Example: +4 -2
SO2
The compound is neutral, so the oxidation number of sulfur will be sufficient to balance out the two oxygen atoms.

27. OF2 Covalent Compounds -1
Example: -1
OF2
Since fluorine is the more electronegative element, it will have its normal oxidation number.

28. OF2 Covalent Compounds +2 -1
Example: +2 -1
OF2
The compound is neutral, so the oxidation number of oxygen will be sufficient to balance out the two fluorine atoms.

29. PO43- Covalent Compounds -2
Example: -2
PO43-
Since oxygen is the more electronegative element, it will have its normal oxidation number.

30. PO43- Covalent Compounds +5 -2 Example:
The ion has a charge of negative three, so the oxidation numbers must add up to the total charge of the ion.
The sum of the oxidation numbers in the formula of a polyatomic ion is equal to its ionic charge

31. Ionic Compounds with Polyatomic
Example: +2
CaSO4
This is an ionic compound, so the charge of the metal cation is its oxidation number

32. Ionic Compounds with Polyatomics
Example: +2
CaSO4
The anion is a polyatomic ion, sulfate, and the charge of sulfate is negative two. So the oxidation numbers of sulfur and oxygen must add to -2

33. Ionic Compounds with Polyatomics
Example: +2 -2
CaSO4
Oxygen is the more electronegative of the two, so it keeps its normal oxidation number.

34. Ionic Compounds with Polyatomics
Example: +2 +6 -2
CaSO4
Sulfur and the four oxygen atoms must add to negative two (the charge of the sulfate anion).

35. Ionic Compounds with Polyatomics
Example:
Pb(OH)4
This is an ionic compound, so the charge of the metal cation is its oxidation number. But this is a transition metal, so we cannot know it from its position on the periodic table.

36. Ionic Compounds with Polyatomics
Example: +4
Pb(OH)4
But the anion, the hydroxide ion, carries a charge of negative one. All four hydroxides are negative one, but since the compound is neutral, the oxidation number of lead must balance it out.

37. Ionic Compounds with Polyatomics
Example: +4 -2
Pb(OH)4
Within the anion, oxygen is the more electronegative of the two elements, and keeps its normal oxidation number.

38. Ionic Compounds with Polyatomics
Example: +4 -2 +1
Pb(OH)4
Within the hydroxide ion, the oxygen and hydrogen must add to the charge of the ion, -1

39. Assign oxidation numbers to all of the elements:
Li2O
Li = +1
O = -2
PF3
P = +3
F = -1
HNO3
H = +1
N = +5
O = -2
MnO4-
Mn = +7
O = -2
Cr2O72-
Cr = +6
O = -2

40. What is the oxidation state of the highlighted element?
P2O5 +5
diphosphorous pentoxide
NaH -1
sodium hydride
SnBr4 +4
tin (IV) bromide
BaO2 -1
barium peroxide

41. Identifying Redox Equations
In general, all chemical reactions can be assigned to one of two classes:
oxidation-reduction, in which electrons are transferred:
Single-replacement, combination, decomposition, and combustion
this second class has no electron transfer, and includes all others:
Double-replacement and acid-base reactions

42. Identifying Redox Equations
In an electrical storm, nitrogen and oxygen react to form nitrogen monoxide:
N2(g) + O2(g) → 2NO(g)
Is this a redox reaction?
If the oxidation number of an element in a reacting species changes, then that element has undergone either oxidation or reduction; therefore, the reaction as a whole must be a redox.

43. they are then balanced separately, and finally combined
A half-reaction is an equation showing just the oxidation or just the reduction that takes place
they are then balanced separately, and finally combined
Step 1: write unbalanced equation
Step 2: write separate half-reaction equations for oxidation and reduction
Step 3: balance all non H or O atoms in the half-reactions (e.g. Cr, Cu, Au, N)

44. Step 4: add H2O to balance any O atoms
Step 5: add H+ to balance the H atoms
Step 6: add electrons to the most positive side of each half-reaction to balance the charges (remember one will be reduction, one will be oxidation)
Step 7: multiply each half-reaction by a number to make the electrons equal in both
Step 8: add the balanced half-reactions to show an overall equation

45. Oxidizing and Reducing Agents
Now the confusing part…
CuO + H2  Cu + H2O
Cu goes from +2 to 0
Cu is reduced, therefore it is called an oxidizing agent because it causes some other substance to be oxidized
H goes from 0 to +1
H is oxidized, therefore it is called a reducing agent because it causes some other substance to be reduced.

46. Some common oxidizing agents
Oxygen! Is an oxidant because it CAUSES oxidation, it itself is reduced though
Oxidized coal in electric power
Gas in automobiles
Wood in campfires
Food we eat
Antiseptics
Hydrogen Peroxide
Benzoyl peroxide
Disinfectants
Chlorine

47. Some common reducing agents
Metals are REDUCTANTS! They cause REDUCTION, but themselves are oxidized (loss electrons)
Antioxidants
Ascorbic acid is used to prevent the browning of fruits by inhibiting air oxidation
Many antioxidants are believed to retard various oxidation reactions that are potentially damaging to vital components of living cells