1. Chapter 15. Oxidation-Reduction Reactions

2. Oxidation-Reduction Reactions
Oxidation was originally understood as reaction of a substance with oxygen.
Combustion:
CH4(g) + 2 O2(g)  CO2(g) + 2 H2O(g)
+ Energy!
Metabolism:
C6H12O6(s) + 6 O2(g)  6 CO2(g) + 6 H2O(g)

3. Oxidation-Reduction Reactions
Corrosion:
2 Mg(s) + O2(g)  2 MgO(g)
4 Fe(s) + 3 O2(g)  2 Fe2O3(s)
Rust!

4. Oxidation-Reduction Reactions
Reduction was originally understood as the loss of mass of metal ores in smelting.
2 Fe2O3(s) + 3 C(s)  4 Fe(s) + 3 CO2(g)

5. Oxidation is a process in which a substance loses electrons
Oxidation is a process in which a substance loses electrons. Reduction is a process in which a substance gains electrons.
4 Fe(s) + 3 O2(g)  2 Fe2O3(s)
4 (Fe0  Fe3+ + 3e1) Iron is oxidized
4 Fe0  4 Fe e1 forms cation
3 (O2 + 4 e1  2 O2) Oxygen is reduced
3 O e1  6 O2 forms anions

6. Oxidation-Reduction Reactions
A half-reaction is a chemical equation that shows either the oxidation or reduction part of an oxidation-reduction reaction.
Electrons appear as products in oxidations.
Zn(s)  Zn2+(aq) e1
Electrons appear as reactants in reductions.
I2(aq) e1  2 I1 (aq)

7. L E O the lion says G E R! Lose Electrons Oxidation
Gain Electrons Reduction

8. Oxidation-Reduction Reactions
Examples:
Which of these half-reactions show oxidations? Which ones show reductions?
Zn(s)  Zn2+(aq) e1
Cu2+(aq) e1  Cu(s)
Ag(s)  Ag1+(aq) + e1
Cr2O72 + 14H e1  2 Cr H2O
2 F1(aq)  F2(g) e1

9. Oxidation-Reduction Reactions
In ionic compounds, we can look at changes in charge to figure out what's reduced and what's oxidized.
What happens with molecular compounds?
CH4(g) + 2 O2(g)  CO2(g) + 2 H2O(g)
Use Oxidation Numbers

10. Determining Oxidation Numbers
An oxidation number (or oxidation state) is the charge that an atom appears to have when the electrons in each bond in which it is participating are assigned to the more electronegative of the two atoms participating in the bond.

11. Determining Oxidation Numbers
Draw Lewis Structure of molecule or ion.
Use Electronegativities (p 281, Fig 7.12) to determine which atom "owns" electrons.
Lone pairs belong to the atom.
If two atoms of the same element are bonded together, 1 electron to each.
ONAtom = VEAtom - TEAtom

12. Determining Oxidation Numbers
The oxidation number of a main-group element can be anything from Column # (loses all e1) to (Column #  8) (fills valence shell).
The sum of all oxidation numbers in a species must equal its charge.
The oxidation number of an atom in its elemental form is zero.
If a Lewis structure is complex, just work with the atom of interest and the atoms bonded to it.

13. Determining Oxidation Numbers
Examples:
Determine the oxidation number of:
N in N2 H in H2 S in S8
C and H in CH4 N and O in N2O
S in H2SO4 Cl in ClO41
C in Urea, CH4N2O

14. Determining Oxidation Numbers
Polyatomic ions in which the central atom is a transition metal:
Oxygen's ON is 2. Hydrogen's ON is +1.
The sum of all oxidation numbers in a species must equal its charge.
What are the oxidation states of the metals?
MnO41 CrO42 Cr2O72

15. Vocabulary A substance is oxidized if: it loses electrons
it loses hydrogen atoms
it gains oxygen atoms
its charge or oxidation number increases (becomes more positive).

16. Vocabulary A substance is reduced if: it gains electrons
it gains hydrogen atoms
it loses oxygen atoms
its charge or oxidation number decreases (becomes less positive).

17. Vocabulary An agent causes something to happen.
An oxidizing agent causes oxidation.
It does this by being reduced (gaining electrons).
A reducing agent causes reduction.
It does this by being oxidized (losing electrons).

18. Vocabulary Typical Oxidizing Agents: O2 Oxygen MnO41 Permanganate
ClO41 Perchlorate CrO42 Chromate
ClO31 Chlorate Cr2O72 Dichromate
NO31 Nitrate Cl2, other halogens
O3 Ozone H2O2 Peroxides

19. Vocabulary Typical Reducing Agents: Hydrogen sources: H2 Hydrogen
NaBH4 Sodium borohydride
LiAlH4 Lithium aluminum hydride
Active metals (1A and 2A metals, zinc)
Reduced Carbon (C, CO)

20. Examples:
Identify the substance being oxidized, the sub-stance being reduced, the oxidizing agent, and the reducing agent:
2 H2(g) + HCCH(g)  C2H6(g)
2 HCl(aq) + Zn(s)  ZnCl2(aq) + H2(g)
4 C3H6O(l) + NaBH4(s) H2O(l)  C3H8O(l) + NaB(OH)4(aq)

21. Examples:
Identify the substance being oxidized, the sub-stance being reduced, the oxidizing agent, and the reducing agent:
2 Na(s) + Cl2(g)  2 NaCl(s)
C2H4O(l) + H2(g)  C2H6O(l)
2 KMnO H2O H2SO4  MnSO4 + K2SO O2(g) H2O

22. Predicting Reactions Who does what to whom?
Cu2+(aq) Ag(s)  Cu(s) Ag1+(aq)
Cu(s) Ag1+(aq)  Cu2+(aq) Ag(s)
Cu2+(aq) + Zn(s)  Cu(s) + Zn2+(aq)
Cu(s) + Zn2+(aq)  Cu2+(aq) + Zn(s)
2 Ag1+(aq) + Zn(s)  Zn2+(aq) Ag(s)
2 Ag(s) + Zn2+(aq)  2 Ag1+(aq) + Zn(s)

23. A Daniell Cell

24. Daniell Cells
Voltages from Daniell Cells provide information about energy changes:
Cu(s) Ag1+(aq)  Cu2+(aq) Ag(s)
0.34 V
Cu2+(aq) + Zn(s)  Cu(s) + Zn2+(aq)
1.10 V
2 Ag1+(aq) + Zn(s)  Zn2+(aq) Ag(s)
1.56 V

25. An Energy Scale for Redox RXN’s
Standard Reduction Potentials are potentials measured against a standard hydrogen elec-trode, with all solutions at 1.0 M and gas pressures at 1.0 atm.
A Standard Hydrogen Electrode is an elec-trode in which the half-reaction is:
2 H e1  H2(g)
Materials are at standard conditions, and the half-cell potential is 0.00 V.

26. An Energy Scale for Redox RXN’s
From chemguide.co.uk

27. An Energy Scale for Redox RXN’s
Some Standard Reduction Potentials:
Reaction E0, V
Li1+(aq) + e1  Li(s) 3.04
Al3+(aq) e1  Al(s) 1.66
Zn2+(aq) e1  Zn(s) 0.76
2 H 1+(aq) e1  H2(g)
Cu2+(aq) e1  Cu(s)
Ag1+(aq) + e1  Ag(s)
Cr2O7 + 14H e1  2 Cr H2O
F2(g) e1  2 F1(aq)

28. An Energy Scale for Redox RXN’s
Strongest oxidizing agents have highest standard reduction potentials (E0's, in volts). Strong oxidizing agents have low-energy vacant orbitals.
Strongest reducing agents have lowest (often negative) standard reduction potentials. Strong reducing agents have electrons in high-energy orbitals.

29. Predicting Reactions Balancing equations for Redox Reactions
Use Table of Standard Reduction Potentials:
Find half-reaction for reduction, write it as given.
Find half-reaction for oxidation, write it as reverse of what is given (electrons are products).

30. Predicting Reactions
Balance half-reactions so number of electrons is equal for each.
Add half-reactions, cancelling out electrons and any other species that appear on both sides of equation.

31. Predicting Reactions Will the Reaction be Spontaneous?
Calculate E0 for RXN: E0 reduction
- E0 oxidation
E0 RXN
If E0 RXN is positive, the reaction is exo- thermic and will proceed spontaneously.
If E0 RXN is negative, the reaction is endo-thermic and will not proceed spontaneously.

32. Predicting Reactions Balance the reactions.
Which of them are spontaneous ?
Cu(s) + Zn2+(aq)  Cu2+(aq) + Zn(s)
Zn(s) + Ag1+(aq)  Zn2+(aq) + Ag(s)
Cu(s) + H1+(aq)  Cu2+(aq) + H2(g)
Fe(s) + Al3+(aq)  Fe2+(aq) + Al(s)
Cu(s) + Cr2O72–(aq)  Cu2+(aq) + Cr3+(aq)

33. Batteries
Batteries are galvanic (a.k.a. voltaic) electro-chemical cells in which a spontaneous reac-tion is used to convert chemical energy to electrical energy.
The earliest batteries common galvanic cells were Daniell Cells, from 1836.

34. Batteries
Batteries usually contain an electrolyte, which is a solution that contains ions. Electric cur-rent can flow through the electrolyte.
The anode in a battery is the half-cell with the lower (more negative) E0. Oxidation occurs at the anode. Anions flow toward the anode.
The cathode in a battery is the half-cell with higher (more positive) E0. Reduction occurs at the cathode. Cations flow toward the cathode.

35. A Daniell Cell

36. Batteries
Dry cell (ca. 1900, developed from Leclanché cell, Alkaline cell, ca. 1950)

37. Batteries Dry cell, acid form; Cell potential ~1.5 V
Anode half-reaction:
Zn(s)  Zn2+(aq) e1–
Cathode half-reaction:
2 MnO2(s) NH4Cl(aq) e1–  Mn2O3(s) NH3(aq) + H2O(l) Cl1–(aq)

38. Batteries Dry cell, alkaline form; Cell potential ~1.5 V
Anode half-reaction:
Zn(s) OH1–(aq)  ZnO(s) + H2O(l) e1–
Cathode half-reaction:
2 MnO2(s) H2O(l) e1–  Mn(OH)2(s) OH1–(aq)

39. Batteries
“Button” battery, ca. 1940, HgO or Ag2O

40. Batteries Button battery; Cell potential ~1.5 V Anode half-reaction:
Zn(s) OH1–(aq)  ZnO(s) + H2O(l) e1–
Cathode half-reaction:
Ag2O(s) + 2 H2O(l) + 2 e1–  2 Ag(s) + 2 OH1–(aq)
HgO(s) + H2O(l) + 2 e1–  Hg22+(aq) + 2 OH1–(aq)

41. Batteries
Lead-Acid Battery (ca. 1859)

42. Batteries
Detail of Lead-Acid Battery: Six cells in series produce about 12 volts.

43. Batteries Lead-acid battery; Cell potential ~2.1 V
Anode half-reaction:
Pb(s) + SO42–(aq)  PbSO4(s) e1– –0.36 V
Cathode half-reaction:
PbO2 (s) + 4 H1+(aq) + SO42–(aq) + 2 e1–  PbSO4(s) + 2 H2O(l)
+1.69 V

44. Batteries Nickel-Metal Hydride (Ni-MH) Battery; cell potential ~1.4V
Anode half-reaction:
MH(s) + OH1–(aq)  M(s) + H2O(l) + e1–
Cathode half-reaction:
NiO(OH)(s) + H2O(l) + e1–  Ni(OH)2(s) + OH1–(aq)
MH(s) is metal alloy, e.g. LaNi5, into which hydrogen is absorbed.

45. Batteries Lithium ion battery; cell potential ~3.7V
Chemistry is complex because water can-not be used in the electrolyte.
Lithium embedded in graphite is the anode.

46. Electrolytic Cells
An electrolytic cell is an electrochemical cell in which a non-spontaneous chemical change is caused to occur by application of electrical energy.
Many important chemical processes are carried out in electrochemical cells.

47. Electrolytic Cells
A chlor-alkali cell

48. Electrolytic Cells
A Hall-Heroult Cell for Aluminum

49. Electrolytic Cells
An electrorefining cell for purifying copper

50. Electrolytic Cells
An electroplating cell