1. REVIEW BOOK TOPIC 9
RED OX

2. Redox Reactions
Chapter 18
+ O2 

3. LEARNING OBJECTIVES
Recognize an oxidation-reduction reaction, in which electrons are transferred from one material (the substance that is oxidized) to another material (the substance that is reduced).
Learn the role of oxidation/reduction in metabolism.
c. Know the difference between the chemical reactions in batteries and those used in electrolysis.

4. DEFINITIONS Oxidation-Reduction Reactions Defined
Oxidation-reduction (“redox”) reactions involve the transfer of electrons from one substance to another.
Oxidized substances lose electrons and reduced substances gain electrons.

5. OXIDATION and REDUCTION
An oxidation-reduction reaction
provides us with energy from food.
provides electrical energy in batteries.
occurs when iron rusts.
4Fe(s) + 3O2(g)  2Fe2O3(s)

6. ELECTRON LOSS OR GAIN An oxidation-reduction reaction
transfers electrons from one reactant to another.
loses electrons in oxidation (LEO) or (OIL)
Zn(s)  Zn2+(aq) + 2e (loss of electrons)
gains electrons in reduction (GER) or (RIG)
Cu2+(aq) + 2e-  Cu(s) (gain of electrons)

7. Zn and Cu +2 Zn(s)  Zn2+(aq) + 2e- oxidation Silvery metal
Cu2+(aq) + 2e-  Cu(s) reduction
Blue orange

8. LEARNING CHECK Identify each of the following as
1) oxidation or 2) reduction.
__A. Sn(s) Sn4+(aq) e−
__B. Fe3+(aq) e−  Fe2+(aq)
__C. Cl2(g) e−  2Cl-(aq)

9. LEARNING CHECK ANSWERS
Identify each of the following as
1) oxidation or 2) reduction.
1 A. Sn(s) Sn4+(aq) e−
2 B. Fe3+(aq) e−  Fe2+(aq)
2 C. Cl2(g) e−  2Cl-(aq)

10. Oxidation-Reduction (Redox) Reactions
“redox” reactions: rxns in which electrons are transferred from one species to another
oxidation & reduction always occur simultaneously
we use OXIDATION NUMBERS to keep track of electron transfers

11. Rules for Assigning Oxidation Numbers:
1) the ox. state of any free (uncombined) element is zero.
Ex: Na, S, O2, H2, Cl2, O3

12. Rules for Assigning Oxidation Numbers:
2) The ox. state of an element in a simple ion is the charge of the ion.
Mg2+  oxidation of Mg is +2

13. Rules for Assigning Oxidation Numbers:
3) the ox. # for hydrogen is +1
(unless combined with a metal, then it has an ox. # of –1)
Ex: NaOH (H bonded to O) v. NaH (H bonded to Na)
H = +1
H = -1

14. Rules for Assigning Oxidation Numbers:
4) the ox. # of fluorine is always –1.

15. Rules for Assigning Oxidation Numbers:
5) the ox. # of oxygen is usually –2.
Why USUALLY? Not -2 when it’s in a peroxide, such as hydrogen peroxide:
H2O2

16. Rules for Assigning Oxidation Numbers:
6) in any neutral compound, the sum of the oxidation #’s = zero.

17. Rules for Assigning Oxidation Numbers:
7) in a polyatomic ion, the sum of the oxidation #’s = the overall charge of the ion.

18. Rules for Assigning Oxidation Numbers:
**use these rules to assign oxidation #’s; assign known #’s first, then fill in the #’s for the remaining elements:

19. NaNO3 oxidation number of a single atom of Na _____ N ______ O ______
LETS PRACTICE
NaNO3 oxidation number of a single atom of Na _____ N ______ O ______

20. Examples: Assign oxidation #’s to each element:
a) NaNO3
Na = +1 O = -2
Therefore, N = +5

21. LETS PRACTICE
SO3 2- oxidation number of a single atom of S _____ O _____

22. Examples: Assign oxidation #’s to each element:
b) SO32- +4 -2
O = -2, therefore S must = +4 to balance the charges and have an overall charge of 2-

23. LETS PRACTICE
HCO3 - oxidation number of a single atom of H _____ C _____ O _____

24. Examples: Assign oxidation #’s to each element:
c) HCO3- +1 +4 -2

25. Examples: Assign oxidation #’s to each element: Do on your own:
H3PO4
Cr2O72-
K2Sn(OH)6

26. Definitions
Oxidation: the process of losing electrons (ox # increases)
Reduction: the process of gaining electrons (ox # decreases)
Oxidizing agents: species that cause oxidation (they are reduced)
Reducing agents: species that cause reduction (they are oxidized)

27. To help you remember…
OIL RIG
Oxidation Is Loss
Reduction Is Gain

28. OXIDATION- REDUCTION Reactions
The chemistry of photography
For photographic film that contains AgBr, light causes loss of an electron by bromide (Br-) and gain of that electron by silver ion (Ag+).
Grains of reduced metallic silver (Ag) form the photographic image.

29. Are all rxns REDOX rxns? You must determine this…
a reaction is “redox” if a change in oxidation # happens; if no change in oxidation # occurs, the reaction is nonredox.

30. Examples
MgCO3 is an ionic compound, so what is Mg’s charge in an ionic compound?
MgCO3  MgO + CO2 +2 -2 +4 -2 +4 +2 -2
The carbonate ion CO32- is the other ion, let’s figure out C because we already know O.
Is this a redox or nonredox reaction?
NONREDOX (no change in oxidation numbers)

31. Examples Zn + CuSO4  ZnSO4 + Cu
Which oxidation numbers do we already know?
Zn + CuSO4  ZnSO4 + Cu +6 +2 -2 +2 +6 -2
0, free element
Break down this ionic compound into its ions
Is this a redox or nonredox reaction?
Cu and SO42-
So, Cu must be Cu2+
REDOX reaction
O = -2 in SO42-, so S must be…?

32. Example… Redox or not?
Zn + CuSO4  ZnSO4 + Cu

33. Example 2 … Redox or not?
CO2 + H2O  C6H12O6 + O2

34. Examples CO2 + H2O  C6H12O6 + O2 +1 -2 +1 +4 -2 -2-2
What happened to Carbon? It went from +4 oxidation # to 0.
Was Carbon oxidized or reduced?
REDUCED
OIL RIG (oxidation is losing electrons so oxidation number increases, where as reduction is gaining electrons so oxidation number decreases)

35. WRITING OXIDATION and REDUCTION Reactions
Write the separate oxidation and reduction reactions for the following equation.
2Cs(s) F2(g)  2CsF(s)
A cesium atom loses an electron to form cesium ion.
Cs(s)  Cs+(s) e− oxidation
Fluorine atoms gain electrons to form fluoride ions.
F2(s) + 2e-  2F−(s) reduction

36. LEARNING CHECK
In light-sensitive sunglasses, UV light initiates an oxidation-reduction reaction.
UV light
Ag+ + Cl−  Ag Cl
A. Which reactant is oxidized?
B. Which reactant is reduced?

37. LEARNING CHECK ANSWERS
In light-sensitive sunglasses, UV light initiates
an oxidation-reduction reaction.
UV light
Ag+ + Cl−  Ag Cl
A. Which reactant is oxidized? Cl−  Cl e−
B. Which reactant is reduced? Ag+ + 1e−  Ag

38. LEARNING CHECK
Identify the substances that are oxidized and reduced in each of the following reactions: A. Mg(s) + 2H+(aq)  Mg2+(aq) + H2(g) B. 2Al(s) + 3Br2(g)  2AlBr3(s)

39. LEARNING CHECK ANSWERS
Mg is oxidized Mg(s)  Mg2+(aq) + 2e−
H+ is reduced 2H+ + 2e−  H2
B. Al is oxidized Al  Al3+ + 3e−
Br is reduced Br + e−  B-

40. Examples NaCl + AgNO3  AgCl + NaNO3
Double displacement reactions are _______ _______!!
Redox or nonredox?

41. Common uses of the terms oxidation and reduction
MEANING
Oxidation
To lose electrons
To increase in oxidation number
Reduction
To gain electrons
To decrease in oxidation number

42. OXIDATION- REDUCTION reactions
Role in Metabolism
Oxidation can also be considered to be gain of oxygen or loss of hydrogen in a molecule.
Reduction can also be considered to be loss of oxygen or gain of hydrogen in a molecule.
Respiration is a redox process whereby living organisms oxidize food to release energy.

43. OXIDATION- REDUCTION reactions
The Chemistry of Batteries
Combining a readily oxidized substance with an easily reduced substance can create a battery.
The oxidized material is the anode and the reduced material is the cathode of the battery

44. Electrochemistry and Redox
Oxidation-reduction: “Redox”
Electrochemistry:
study of the interchange between chemical change and electrical work
Electrochemical cells:
systems utilizing a redox reaction to produce or use electrical energy

45. (Balancing redox reactions)
REDOX REVIEW
Redox reactions: electron transfer processes
Oxidation: loss of 1 or more e-
Reduction: gain of 1 or more e-
Oxidation numbers: imaginary charges
(Balancing redox reactions)

46. Assigning oxidation Oxidation Numbers review … the rules
1. Pure element O.N. is zero
2. Monatomic ion O.N. is charge
3. Neutral compound: sum of O.N. is zero
Polyatomic ion: sum of O.N. is ion’s charge
*Negative O.N. generally assigned to more electronegative element

47. Assigning oxidation Oxidation Numbers review … the rules p.2
4. Hydrogen
assigned (metal hydrides, -1)
5. Oxygen
assigned (peroxides, -1; OF2, +2)
6. Fluorine
always -1

48. REDOX review cont. Oxidation is loss of e- causes reduction
“reducing agent”
Reduction is gain of e causes oxidation “oxidizing agent”

49. REDOX HALF-REACTION review
1. Write separate equations (half-reactions)
for oxidation and reduction
2. For each half-reaction
a. Balance elements involved in e- transfer
b. Balance number e- lost and gained
3. To balance e-
multiply each half-reaction by whole numbers

50. TYPES of CELLS
Voltaic (galvanic) cells: a spontaneous reaction generates electrical energy Electrolytic cells: absorb free energy from an electrical source to drive a nonspontaneous reaction

51. Common Components Electrodes:
conduct electricity between cell and surroundings
Electrolyte:
mixture of ions involved in reaction or carrying charge
Salt bridge:
completes circuit (provides charge balance)

52. ELECTRODES Anode: Oxidation occurs at the anode Cathode:
Reduction occurs at the cathode
Active electrodes: participate in redox

53. Voltaic (Galvanic) Cells
A device in which chemical energy is changed to electrical energy.
Uses a spontaneous reaction.

57. Zn2+(aq) + Cu(s)  Cu2+(aq) +Zn(s)
Zn gives up electrons to Cu
“pushes harder” on e-
greater potential energy
greater “electrical potential”
Spontaneous reaction due to
relative difference in metals’ abilities to give e-
ability of e- to flow

58. Cell Potential Cell Potential / Electromotive Force (EMF):
The “pull” or driving force on electrons
Measured voltage (potential difference)

60. Batteries
A battery is a galvanic cell or, more commonly, a group of galvanic cells connected in series.

62. Corrosion
Some metals, such as copper, gold, silver and platinum, are relatively difficult to oxidize.
These are often called noble metals

64. ELECTROCHEMICAL CELLS
Voltaic cell
Chemical energy is convert to electrical energy
* Spontaneous
chemical reaction
Electrolytic cell
Electrical energy is convert to chemical energy
* Nonspontaneous
chemical reaction

67. Electrolysis
Forcing a current through a cell to produce a chemical change for which the cell potential is negative.