1. Oxidation Reduction Chemistry
Redox Chemistry

2. Oxidation and Reduction reactions always take place simultaneously.
Loss of electrons – oxidation
(Increase in Oxidation Number)
Ex:Na > Na+1 + e-1
Gain of electrons – reduction
( Decrease in Oxidation Number)
Cl2 + 2 e > 2 Cl-1

3. LEO the lion says GER:
Lose Electrons, Oxidation; Gain Electrons, Reduction

4. Oxidation occurs when a molecule or atom loses electrons
Reduction occurs when a molecule or atom gains electrons.

6. Zinc is being oxidized while the copper is being reduced. Why?
Zn(s) + CuCl2 (aq)  ZnCl(aq) + Cu(s)
Net ionic equation
Cu2+ (aq) + Zn(s)  Cu(s) + Zn2+aq)
Zinc is being oxidized while the copper is being reduced. Why?

7. Redox reactions involve electron transfer: Lose e - =Oxidation
Cu(s) + 2 Ag1+(aq) Cu2+(aq) + 2 Ag(s)
Gain e - =Reduction

8. Rules for Assigning Oxidation States
1.The oxidation number of an atom in an uncombined element is always The oxidation number of any ion equals its ionic charge
Rules for Assigning Oxidation States

9. 3. In compounds, the oxidation number of many elements corresponds to the elements position in the periodic table
a. alkali metals are always +1
b. alkaline earth metals are always +2
c. Aluminum is +3
d. Halogens are almost always –1 unless with an oxianion
e. H has an oxidation # of +1 when combined with nonmetals
f. Oxygen is always -2

10. Rules for Assigning Oxidation States
4. The oxidation numbers of elements in compounds are written per atom 5. The sum of the oxidation numbers of all atoms in a compound must equal The sum of the oxidation numbers of all atoms in a polyatomic ion must equal the charge of the ion.

11. Identifying oxidized species, reduced species or spectator species1+ 1- 2+ 1-
Ca + 2 H Cl  Ca Cl 2 + H2
Oxidation = loss of electrons; becomes more positive in charge. Ca0  Ca+2, so Ca0 is the species that is oxidized.
Reduction = gain of electrons; becomes more negative in charge. H+1  H0, so the H+1 is the species that is reduced.
Spectator Ion does not gain or lose any electrons; no change in charge. Cl-1  Cl-1, so the Cl-1 is the spectator ion.

12. SO2 CO32– Na2SO4 (NH4)2S Assigning Oxidation Numbers to Atoms +1 +6 –2
What is the oxidation number of each kind of atom in the following ions and compounds?
a. SO2 c. Na2SO4
b. CO32– d. (NH4)2S
Na2SO4 –2
SO2
+4 –2
(NH4)2S
– –2
CO32–
+4 –2

13. Oxidizing Agents - Reducing Agents:
An oxidizing agent gets reduced itself by “taking” electrons from another substance.
That substance is then a reducing agent and gets oxidized itself, because it “gave” electrons.
So a strong oxidizing agent should have a great tendency to accept e-
A strong reducing agent should be willing to lose e- easily.
What are strong oxidizing agents- metals or non metals? Why?

14. Oxidizing and Reducing Agents
Ca0 + 2 H+1Cl-1  Ca+2Cl-12 + H20
Since Ca0 is being oxidized and H+1 is being reduced, the electrons must be going from the Ca0 to the H+1.
Since Ca0 would not lose electrons (be oxidized) if H+1 weren’t there to gain them, H+1 is the cause, or agent, of Ca0’s oxidation. H+1 is the oxidizing agent.
Since H+1 would not gain electrons (be reduced) if Ca0 weren’t there to lose them, Ca0 is the cause, or agent, of H+1’s reduction. Ca0 is the reducing agent.

15. Identifying Oxidized and Reduced Reactants
Silver nitrate reacts with copper to form copper nitrate and silver. From the equation below, determine what is oxidized and what is reduced. Identify the oxidizing agent and the reducing agent.
2AgNO3(aq) + Cu(s) → Cu(NO3)2(aq) + 2Ag(s)
Oxidation: Cu → Cu2+ + 2e– (loss of electrons)
Reduction: 2Ag+ + 2e– → 2Ag (gain of electrons)
Cu : reducing agent. Ag+ : oxidizing agent.

16. The element that donates electrons in a redox reaction is called the reducing agent. Which of the following is always true of the reducing agent?
A. It is oxidized.
B. It is reduced.
C. It is ionic.
D. It is covalent.

17. Identifying Oxidized and Reduced Atoms
Use changes in oxidation number to identify which atoms are oxidized and which are reduced in the following reactions. Also identify the oxidizing agent and the reducing agent.
a. Cl2(g) + 2HBr(aq) → 2HCl(aq) + Br2(l)
b. C(s) + O2(g) → CO2(g)

18. Cl2(g) + 2HBr(aq) → 2HCl(aq) + Br2(l)
– –
Cl2(g) + 2HBr(aq) → 2HCl(aq) + Br2(l)
Chlorine is reduced, so Cl2 is the oxidizing agent.
The bromide ion from HBr(aq) is oxidized, so Br– is the reducing agent.

19. 2HNO3(aq) + 3H2S(g) → 2NO(g) + 4H2O(l) + 3S(s)
Use changes in oxidation number to identify which atoms are oxidized and which are reduced in the following reaction.
+1 +5 – – – –
2HNO3(aq) + 3H2S(g) → 2NO(g) + 4H2O(l) + 3S(s)
Sulfur is oxidized because its oxidation number increases (–2 to 0). Nitrogen is reduced because its oxidation number decreases
(+5 to +2).

20. Identifying Redox Equations
In general, all chemical reactions can be assigned to one of two classes:
oxidation-reduction: electrons are transferred
single-replacement, combination, decomposition, and combustion
no electron transfer:
double-replacement
acid-base reactions

21. 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?
δ+N  Oδ−
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.

22. Question
Answer
Is the reaction redox?
If any atoms change their oxidation number, the reaction is redox.
Which element is oxidized?
The element that increases its oxidation number is oxidized
Which element is reduced?
The element that decreases its oxidation number is reduced.
What’s the reducing agent?
The reactant that contains the element that is oxidized is the reducing agent.
What’s the oxidizing agent?
The reactant that contains the element that is reduced is the oxidizing agent.

23. Identifying Redox Reactions
Use the change in oxidation number to identify whether each reaction is a redox reaction or a reaction of some other type. If a reaction is a redox reaction, identify the element reduced, the element oxidized, the reducing agent, and the oxidizing agent.
a. Cl2(g) + 2NaBr(aq) → 2NaCl(aq) + Br2(aq)
b. 2NaOH(aq) +H2SO4(aq) → Na2SO4(aq)+ 2H2O(l)

24. Cl2(g) + 2NaBr(aq) → 2NaCl(aq) + Br2(aq)
– –
Cl2(g) + 2NaBr(aq) → 2NaCl(aq) + Br2(aq)
This is a redox reaction.
The chlorine is reduced.
The bromide ion is oxidized.
Chlorine is the oxidizing agent; the bromide ion is the reducing agent.

25. 2NaOH(aq) + H2SO4(aq) → Na2SO4(aq) + 2H2O(l)
+1 – – – –2
2NaOH(aq) + H2SO4(aq) → Na2SO4(aq) + 2H2O(l)
None of the elements change in oxidation number.
This is not a redox reaction.

26. Balancing Redox Equations (in acid solution)
The fundamental principle involves electron transfer. The transfer of electrons MUST be the same in the reduction process and the oxidation process.
First write half reactions
1. assign oxidation numbers for all atoms.
2. Identify the element that is oxidized and the element reduced. (LEO says GER). Split the reaction into half-reactions (reduction, oxidation)
Balance elements in half reactions
3. Balance all elements except hydrogen and oxygen.
4. Balance oxygen atoms by adding H2O. Then balance the hydrogen atoms on the other side by adding H+.

27. Balancing Redox Equations (in acid solution)
Balance charge in each half reaction
5. Add sufficient electrons (e-)to one side of each half reaction to balance the charges. Electrons appear on the left in the reduction half reaction and appear on the right in the oxidation half reaction
Balance electrons from oxidation to reduction
6. Balance the electrons by finding a common factor for the electrons. Then multiply each half reactions by some integer so the # of electrons are the same in each half-reaction.
Add half reactions
7. Add the two half-reactions, canceling and reducing any elements or compounds. The electrons should cancel on each side.
8. Check your work!

28. Fe2+(aq) + MnO4-(aq) Fe3+(aq) + Mn2+(aq) (acidic solution)
1. Write the half reaction:
Fe2+(aq)  Fe3+(aq) (oxidation)
MnO4-(aq)  Mn2+(aq) (reduction)
2. Balance each half reaction:
Fe2+(aq)  Fe3+(aq) + e-
MnO4-(aq) + 8H+(aq) + 5e-  Mn2+(aq) + 4H2O
3. Combine the half reactions to give the overall reaction:
5Fe2+(aq)  5Fe3+(aq) + 5e-
MnO4-(aq) + 8H+(aq) + 5e-  Mn2+(aq) + 4H2O
5Fe2+(aq) +MnO4-(aq) + 8H+(aq) 5Fe3+(aq) + Mn2+(aq) + 4H2O

29. Practice Half-Reactions
Don’t forget to determine the charge of each species first! 4 Li + O2  2 Li2O Oxidation Half-Reaction: Reduction Half-Reaction: Zn + Na2SO4  ZnSO4 + 2 Na

30. Balancing Redox Reactions: Basic solution
Balance following “acidic”steps.
Add an equal number of OH- ions to both sides of the equation to “neutralize” H+ ions.
Cross out waters on each side.
4 H+ + MnO4- --> MnO2 + 2 H2O
4 OH- + 4 H+ + MnO4- --> MnO2 + 2 H2O + 4OH-
4 H2O + 4OH- + 4 H+ + MnO4- --> MnO2 + 2 H2O + 4 OH-
2 4 H2O + MnO4- --> MnO2 + 2 H2O + 4 OH-

31. Standard Cell Potential
Just as the water tends to flow from a higher level to a lower level, electrons also move from a higher “potential” to a lower potential. This potential difference is called the electromotive force (EMF) of cell and is written as Ecell.
The standard for measuring the cell potentials is called a SHE (Standard Hydrogen Electrode).
Description of SHE (Standard Hydrogen Electrode)
Reaction 2H+(aq, 1M)+ 2e - H2(g, 101kPa) E0= 0.00 V

32. Standard Reduction Potentials
Many different half cells can be paired with the SHE and the standard reduction potentials for each half cell is obtained. Check the table for values of reduction potential for various substances:
Would substances with high reduction potential be strong oxidizing agents or strong reducing agents?
Why?

33. Higher up the list: More active metals – lose electrons more easily
Higher up the list: More active metals – lose electrons more easily. More easily oxidized. A metal high on the list can transfer electrons to a metal lower on the list.

34. Zn is above Pb in the activity series of metals
Zn is above Pb in the activity series of metals. Which of the following statements is correct?
A. Zn will react with Pb2+.
B. Pb2+ will react with Zn2+.
C. Zn2+ will react with Pb.
D. Pb will react with Zn2+.

35. Metal Activity
Metallic elements start out with a charge of ZERO, so they can only be oxidized to form (+) ions. The higher of two metals MUST undergo oxidation in the reaction, or no reaction will happen. Refer to the activity series, which reaction will happen? Which won’t? 3 K + FeCl3  3 KCl + Fe Fe + 3 KCl  FeCl3 + 3 K

36. Some applications of Redox Reactions: 1
Some applications of Redox Reactions: 1. Preventing Corrosion: Corrosion is the oxidation of some metal caused by some substance in the environment. (Rust) 2. Bleaching: A way to remove unwanted color from fabric or other materials. Color is caused by movement of electrons between different energy levels and bleaching removes these electrons. Bleaches are oxidizing agents. Two common bleaches are hypochlorite, ClO-, and hydrogen peroxide, H2O2.

37. Some applications of Redox Reactions: 3
Some applications of Redox Reactions: 3. Fuels and Explosives: Fuels, like gasoline, are oxidized by oxygen and form CO2 and H2O. Explosives, like nitroglycerin (C3H5(NO3)3 contain both oxidizing and reducing agents. 4. Photography: Film processing is based on the oxidation of Bromide ions and the reduction of silver ions.

38. Reduction at the Cathode
Voltaic Cells (Galvanic Cells)
A voltaic cell converts chemical energy from a spontaneous redox reaction into electrical energy.
A reaction is spontaneous if the metal with higher reduction potential (lower on activity series) is made the cathode.
Higher on activity series = (-) anode
(lower reduction potential)
Lower on activity series = (+) cathode
(higher reduction potential)
Reduction at the Cathode

39. Voltaic Cells
cathode
anode
Zn(s)
Cu(s)
Zn(NO3)2
Cu(NO3)2

40. A voltaic cell is formed from a piece of iron in a solution of Fe(NO3)2 and silver in a solution of AgNO3. Which is the cathode, and which is the anode? Why?
The iron electrode is the anode because it is the most easily oxidized. The silver electrode is the cathode because silver is below iron in the activity series and is therefore reduced in the spontaneous redox reaction.

41. Electrolytic Cells
Use electricity to force a nonspontaneous redox reaction to take place.
Uses for Electrolytic Cells:
decomposition of alkali metal compounds
electrolysis of water
electroplating

42. Electrolytic cells – NOT spontaneous Need energy to proceed Anode is +; Cathode is – No salt bridge

43. Decomposing Alkali Metal Compounds
2 NaCl  2 Na + Cl2
The Na+1 is reduced at the (-) cathode, picking up an e- from the battery
The Cl-1 is oxidized at the (+) anode, the e- being pulled off by the battery (DC)

44. Decomposing Water 2 H2O  2 H2 + O2
The H+ is reduced at the (-) cathode, yielding H2 (g), which is trapped in the tube.
The O-2 is oxidized at the (+) anode, yielding O2 (g), which is trapped in the tube.

45. Electroplating
The Ag0 is oxidized to Ag+1 when the (+) end of the battery strips its electrons off.
The Ag+1 migrates through the solution towards the (-) charged cathode (ring), where it picks up an electron from the battery and forms Ag0, which coats on to the ring.

46. Voltaic (galvanic)
Electrolytic
Spontaneous
Yes No
Anode
negative
positive
Cathode
# of cells 2 1
Salt bridge