1. Reduction and Oxidation of Metals
Chapter A2
Reduction and Oxidation of Metals

2. 2.1 - Compounds and Chemical Change
Chapter A2
2.1 - Compounds and Chemical Change

3. Metallurgy
Archeologist studies show that ancient civilizations were not only using metals, they learned to manipulate it
metallurgy – heating and hammering copper into shapes
by analyzing copper artifacts, archeologists can use chemical analysis to track trade routes, locations of ancient industry, etc.

4. Tarnishing
The colour change, and loss of luster, associated with metal over time
Silver tends to turn black over time
Copper tends to turn green

5. Silver tarnish
When a metal tarnishes, a chemical reaction is occurring between the metal, and molecules in the air.
Balance the following reaction between pure silver, and hydrogen sulfide and oxygen in the air.
__Ag(s) + __H2S(g) + __O2(g)  __Ag2S(s) + __ H2O(l)

6. Silver Tarnish
Draw Lewis dot diagrams for the entire reaction

7. Balanced Equations 4 Ag(s) + 2 H2S(g) + 1 O2(g)  2 Ag2S(s) + 2 H2O(l)
Notice in the last slide, the number of molecules you had to draw corresponded to the coefficients of the balanced equation.
4 Ag(s) + 2 H2S(g) + 1 O2(g)  2 Ag2S(s) + 2 H2O(l)
When an equation is balanced, you can use the coefficients to compare the ratios of one compound to another

8. Balancing Single Replacement Equations
**Recall from Sci 10, the rules for balancing an equation:
This means we must always have the same number of each type of atom on both sides of the equation
Determine the correct chemical formula for all reactants and products.
check for diatomic molecules.
check for polyatomic ions.
indicate correct state of compounds. (s, l, g, aq)
Balance metals
Balance nonmetals
Balance hydrogen
Balance oxygen
Recount all atoms
If every coefficient will reduce, rewrite the whole equation using the simplest ratio of coefficients.

9. Balance by Inspection
**Balance by adding a coefficient to the front of the chemical formula.
Coefficients must be whole numbers.
Do not change subscripts in chemical formula.
Do not place coefficients between atoms or ions in a formula.
Number of polyatomic ions must be the same on both sides of the equation

10. Examples: __KI(aq) + __Cl2(g) → __ KCl(aq) +__ I2(s)
__Al(s) + __H2SO4(aq) → __Al2(SO4)3(aq) + __ H2(g)
__Co(NO3)3(aq)+ __Zn(s)→ __Zn(NO3)2(aq)+__Co(s)

11. Mole ratios
mole ratio – the ratio of the coefficients in a balanced chemical equation
n required = coeffiecient required
n given coeffiecient given
A ratio does not tell you the exact amounts, but it tells you proportions
e.g. a mole ratio of 3:1 means three moles of the wanted for every mole of the given
if you know the number moles of the given, you can use the mole ratio to calculate the number of moles of the required

12. Steps to solving a mole ratios problem:
write the balanced chemical equation for the reaction
under the equation, write the number of moles of the given (nG) and “nR = ?” under the required substance
find the mole ratio by putting the coefficient for the required substance over the given substance’s coefficient
multiply the number of moles of the given by the mole ratio (R/G) to find the number of moles of the wanted

13. 4 Ag(s) + 2 H2S(g) + 1 O2(g)  2 Ag2S(s) + 2 H2O(l)
Practice problem #1:
Determine the amount of silver required to make mol of silver sulfide, based on the reaction given above.
write the balanced chemical equation for the reaction
4 Ag(s) + 2 H2S(g) + 1 O2(g)  2 Ag2S(s) + 2 H2O(l)

14. 4 Ag(s) + 2 H2S(g) + 1 O2(g)  2 Ag2S(s) + 2 H2O(l)
Practice problem #1:
Determine the amount of silver required to make mol of silver sulfide, based on the reaction given above.
under the equation, write the number of moles of the given (nG) and “nR = ?” under the required substance
R G
4 Ag(s) + 2 H2S(g) + 1 O2(g)  2 Ag2S(s) + 2 H2O(l)
nR = ? nG = mol

15. 4 Ag(s) + 2 H2S(g) + 1 O2(g)  2 Ag2S(s) + 2 H2O(l)
Practice problem #1:
Determine the amount of silver required to make mol of silver sulfide, based on the reaction given above.
find the mole ratio by putting the coefficient for the required substance over the given substance’s coefficient
4 Ag(s) + 2 H2S(g) + 1 O2(g)  2 Ag2S(s) + 2 H2O(l)
mole ratio = coefficientR = 4
coefficientG 2

16. Practice problem #1:
Determine the amount of silver required to make mol of silver sulfide, based on the reaction given above.
multiply the number of moles of the given by the mole ratio (R/G) to find the number of moles of the wanted
nR = nG (mole ratio) = mol (4/2) = 1.75 mol

17. Practice problem #2:
How many moles of oxygen are required to burn 5.0mol of ethane (C2H2)?

18. Practice problem #2 (solution):
How many moles of oxygen are required to burn 5.0mol of propane (C3H8(g))?
C3H8(g) + 5 O2(g)  3 CO2(g) + 4 H2O(g)
nG = 5.0mol nR = ?
nR = nG (mole ratio R/G) = 5.0mol (5/1) = 25mol
*remember that when there is no coefficient, it represents a “1” in a mole ratio

19. Homework
PRACTICE PROBLEMS
#4-6 page 63,
#4-7 page 68

20. A2.2 – The Gain and Loss of Electrons
Chapter A2
A2.2 – The Gain and Loss of Electrons

21. Making metal useful:
Metals are an important resource, but pure metals are rarely found in nature
the term ore refers to a rock containing enough useful metal to be mined

22. Formation of an ionic compound
a metal is made of cations surrounded by free floating electrons
some of these electrons will be lost to other atoms in the environment
the number of free-floating electrons is no longer equal to the number of positive charges
the metals that lose the electrons become positively-charged ions
Calcium atom
20 p+
20 e-
Net charge = 0
Calcium ion
20 p+
18 e-
Net charge = 2+

23. Formation of an ionic compound
the atoms that picked up the extra electrons become negatively-charged ions
the positively-charged ions are attracted to the negatively- charged ions and an ionic bond is formed
Phosphorus atom
15 p+
15 e-
Net charge = 0
Phosphorus ion
15 p+
18 e-
Net charge = 3 –

24. Oxidation: The Loss of Electrons
when an atom loses an electron it is oxidized
oxidation is a chemical process which a substance loses electrons
two types of oxidation reactions
metal atom (neutral)  metal ion + electrons
e.g. Fe(s)  Fe2+­(aq) + 2e-
non-metal ion (charged)  non-metal atom + electrons
e.g. 2Cl-(aq)  Cl2(g) + 2e-

25. Reduction: The Gain of Electrons
when an atom gains an electron it is reduced
two types of reduction reactions
metal ion (charged) + electrons  metal atom
e.g. Fe2+­(aq) + 2e-  Fe(s)
non-metal atom (neutral)  non-metal ion + electrons
e.g. Cl2(g) + 2 e-  2Cl-(aq)

26. Single Replacement Reactions
in order to get a valuable metal our of an ionic compound, a less valuable metal can be used to take its place in the compound
recall, in a single replacement reaction, an element reacts with a compound to produce a new element and a new compound
That means there are three elements involved altogether:
One element will be oxidized
One element will be reduced
One element will be the “spectator”

27. Example:
Most silver is found as the compound silver sulfate. The pure silver can be collected by reacting the silver sulfate with zinc metal.
Ag2SO4(s) + Zn(s)  ZnSO4(s) + 2Ag(s)

28. Example: Ag2SO4(s) + Zn(s)  ZnSO4(s) + 2Ag(s) 2Ag+(aq) + 2e-  2Ag(s)
the zinc loses electrons and is oxidized, because
Zn(s)  Zn2+(aq) + 2e-
the silver gains electrons and is reduced, because
2Ag+(aq) + 2e-  2Ag(s)
the sulfate is unaffected and is called a spectator ion, because
SO42-(aq)  SO42-(aq)
a reaction that includes oxidation and reduction is called a redox reaction

29. Homework PROBLEMS TO BE HANDED IN: Page 68 #5-7, Page 74 #18,

30. A2.3 – The Reactivity of Metals
Chapter A2
A2.3 – The Reactivity of Metals

31. Stability vs. Reactivity of Metals
Some metals are very stable – they can be found in pure form, and don’t often form chemicals (e.g. gold)
Other metals tend to corrode easily (tarnishing and rusting are both examples of corrosion)
corrosion is the oxidation of a metal

32. Simulation
Four different metals are immersed into four different ionic solutions, and the results recorded.
If nothing occurs, then “no reaction” will be noted.
If a change occurs, then “reaction” will be noted.
The results should be as follows:

33. Simulation - results Solutions Mg2+(aq) Cu2+(aq) Zn2+(aq) Ag+(aq)
Metals
magnesium
no reaction
reaction
copper
zinc
silver
rank the ions in order from most to least reactive
Ag+  Cu2+  Zn2+  Mg2+
rank the metal atoms in order from most to least reactive
Mg  Zn  Cu  Ag
what connection do you see between the reactivity of an ion and its stability as an atom?
The more stable a metal atom is, the more reactive it is as an ion

34. Activity Series
A list that organizes metal atoms from most stable to least stable
Simultaneously organizes metal ions from most reactive to least reactive
see page 80 in your text, or page 4 of your data booklet

35. Activity Series Cd2+(aq) + 2e-  Cd(s)
when read left to right you get a reduction half- reaction
Example:
Cd2+(aq) + 2e-  Cd(s)

36. Activity Series Example: Cd(s)  Cd2+(aq) + 2e-
when read right to left you get an oxidation half- reaction
Example:
Cd(s)  Cd2+(aq) + 2e-

37. Activity Series an activity series allows you to:
compare the relative reactivity of metal ions
using left side, whichever is listed closer to the top
compare the relative reactivity of metal atoms
using right side, whichever is listed closer to the bottom

38. Activity Series determine if a reaction will occur spontaneously
a spontaneous reaction is one that will occur by itself without any addition of energy
to determine if a reaction is spontaneous, locate the two half-reactions on the table
if the reduction reaction (LR) is located above the oxidation reaction (RL) it will occur spontaneously
a non-spontaneous reaction will have the reduction reaction located below the oxidation reaction
See Example Problem (page 82)

39. Oxidizing and reducing agents
A travel agent does not travel herself, but rather allows someone else to travel
In a similar way, the oxidizing agent is the one that causes the other substance to be oxidized, and the reducing agent causes the other substance to be reduced.

40. Oxidizing and reducing agents
in a net reaction, the reducing agent is the one that causes another substance to be reduced, and is itself oxidized
in a net reaction, the oxidizing agent is the one that causes another substance to be oxidized, and is itself reduced

41. Special properties of gold:
least chemically reactive metal
most likely to be found in pure form
least likely to form ionic compounds
excellent conductor that can be stretched into very fine wires
gold is favored for jewelry and coins
also favored for electronics and computer manufacturing

42. Homework
PRACTICE PROBLEMS 24-32

43. Chapter A2
A2.4 – Voltaic Cells

44. Voltaic cells
The focus of our studies on metals has so far been focused on what’s happening on the surface of the metal (e.g. tarnishing)
The rest of this chapter focuses on another area of the reaction – the movement of electrons between metals
just like the word “cell” in biology, the term cell in chemistry refers to a distinct structure that interacts with the environment
an example of a cell is a AA battery

45. Voltaic cells
though commonly we refer to a voltaic cell as a “battery”, technically cells are only referred to as a battery when several are together
when an electronic device is operating, voltaic cells provide a continuous flow (current) of electrons, which is converted into current to power the device

46. Voltaic cells the voltaic cell that you know looks like this:
this is simple and compact version of the voltaic cell we will make in the lab.

47. Voltaic cells
an electrode is a solid piece of metal that is suspended in a solution and connected to an external circuit.

48. Voltaic cells
the electrode zinc, is immersed into an electrolyte solution, where the zinc electrode acquires an excess of electrons, becoming negatively charged

49. Voltaic cells
the other electrode is usually composed of a different material (usually copper) and will become positively charged

50. Voltaic cells
once a current is closed between the two electrodes, the electrons will repel from the zinc electrode, pass through the circuit and be used by whatever device is connected, then flow through the positive electrode

51. Voltaic cells
the reaction will continue until one substance can no longer be sustained.

52. Voltaic cells
while cells can often be built with both electrodes in the same container, the cells that we will study in Science 20 are a little more complex
instead of both electrodes being suspended in the same electrolyte solution, each is in its own container
each electrode is suspended in a solution containing the same metal’s ions - e.g. Zn is suspended in Zn2+(aq)

53. Voltaic cells
a salt bridge is a glass U-shaped tube that is filled with an ionic solution
this is to allow for free flow of electrons from one solution to the other

54. How the cell works:
because it is the more reactive of the two metals, the zinc electrode will become oxidized, giving away electrons
these electrons will travel from the electrode, through a metal wire, and then into an electronic device
in the lab, the device we will be using will be a voltmeter, allowing us to measure the amount of electricity passing through

55. How the cell works:
the electrons will pass through the device, back through another wire, into the copper electrode
these electrons will be attracted to the Cu2+(aq) ions in the solution
over time, the zinc electrode will shrink in size (as Zn  Zn2+) and the copper electrode will grow (Cu2+  Cu)
if the two solutions were not connected, the zinc would run out of electrons and the cell would stop working
the solution in the salt bridge allows a continuous flow of electrons back into the zinc solution

56. How the cell works:

57. Analyzing a voltaic cell: Step #1: identify the electrode where oxidation occurs
locate the two metals on the activity series (right side)
the metal closer to the bottom will be oxidized = reducing agent
the electrode that is oxidized is called the anode
the other electrode is reduced, and is called the cathode

58. Analyzing a voltaic cell: Step #2: describe the oxidation process in the anode
write the oxidation half-reaction
the anode will decrease in size over time because the metal is turning into metal ions
electrons leave the anode and travel to the external circuit running the electronic device
in the lab, a voltmeter is used instead of an electronic device – this allows you to measure the amount of electricity being produced
because the anode is the electrode where the electrons originate, it is considered the negative electrode

59. Analyzing a voltaic cell: Step #3: describe the reduction process in the cathode
the electrons will travel through the electronic device and back into the cathode
the electrons will be attracted to the positively-charge metal ions in the cathode solution
the cathode ions will be deposited on the metal electrode

60. Analyzing a voltaic cell: Step #4: describe how the salt bridge completes the circuit
eventually, the anode would run out of electrons and the voltaic cell would stop working
the salt bridge connects the cathode back to the anode to allow the replenishment of electrons on the anode side
the salt bridge contains a third solution
the positive ions from the solution will be attracted to the cathode, while the negative ions from the solution will migrate toward the ion

61. Example: in this voltaic cell: zinc is the anode – it is oxidized
copper is the cathode – it is reduced
the solution in the salt bridge is KCl(aq)
chloride ions are a spectator ion – their job is to replenish the electron supply at the anode

62. Practice Problem: 1. Draw a voltaic cell using the following supplies:
two beakers
U-tube & cotton balls
wire & voltmeter
tin and magnesium strips
solutions of SnSO4(aq), MgSO4(aq), and NaNO3(aq)
2. Label the direction of e- flow, the anode, cathode, OA, RA, - and + electrodes, voltmeter and salt bridge

63. Practice Problem (Solution):

64. Cell notation
voltaic cells can also be represented using short hand cell notation
Zn(s) | Zn2+(aq) || Cu2+(aq) | Cu(s)
anode salt bridge cathode
 oxidation reduction
the anode is listed on the left, the cathode on the right
the vertical line | represents a boundary between a metal and its solution
the double line || represents the salt bridge

65. Homework:
VOLTAIC CELLS WORKSHEET

66. Chapter A2
A2.5 – Electrolytic Cells

67. Electrolytic vs. voltaic
an electrolytic cell is a system where a non- spontaneous redox reaction occurs
recall that a reaction that is non-spontaneous will only occur if energy is added
in an electrolytic cell, energy is added in the form of electricity
Voltaic
Electrolytic
spontaneous?
yes no
requires energy?
produces voltage?
use
energy source
electroplating
change in energy
exothermic
endothermic

68. Electroplating
the metals, like gold, that are the most stable and corrosion-resistant are also the most precious (and expensive)
if you want to manufacture a metal object that is resistant to corrosion it would not be cost-effective to make the whole thing out of gold
instead, a thin coating of gold could be applied to the surface of a more affordable metal

69. Electroplating
the object to be coated will be submerged in a solution of metal ions (e.g. silver ions)
an external energy source supplies energy to the electrons to force them to flow from the electrode into the solution
this turns the plating metal ions into metal atoms, which will accumulate on the surface of the object to be plated.

70. Electrolytic cells Step #1:
electrons from the plating metal (e.g. gold) are attracted to the + electrode of the power source
this leaves the Au(s) atoms short on electrons, which causes Au+(aq) ions to be added into the solution

71. Electrolytic cells Step #2: Step #3:
electrons flow through the power source
Step #3:
electrons accumulate on the surface of the object to be plated

72. Electrolytic cells Step #4:
Au+(aq) ions from the solution are attracted to the electrons on the object to be plated
when they gain these electrons, they turn into solid gold again and form a gold coating

73. Electroplating
electroplating is a particularly good way to protect metals that are easily oxidized, like iron
metals that work as good electroplaters are chromium, platinum, silver and gold

74. Gold jewelry solid gold gold plated
if you have a piece of gold plated jewelry, care must be taken to avoid any deep scratches
deep scratches will expose the oxidizable metal underneath
solid gold
karats - pure gold is 24K
gold is a soft metal, so it is often combined with other metals like brass (copper and zinc) and nickel to make it more durable
the number of karats in the gold refers to how many 1/24th of gold it contains
by law, every piece of gold jewelry must be stamped with the karat mark

75. Other uses for electrolytic cells
refining metals
a sample of impure metal (anode), pure metal (cathode)
ions of the pure metal will travel from the anode to the cathode to build up the atoms of pure metal
electrolysis
decomposition of a compound by means of an electric current
e.g. electrolysis of water makes it decompose into O2 and H2

76. Other uses for electrolytic cells
producing non-metals
non-metals, especially the halogens, are difficult to obtain in pure form because they are so reactive
non-metal atoms will accumulate around the anode of an electrolytic cell
recharging voltaic cells
when you use a battery recharger, you are using an electrolytic cell to reverse the process that occurs normally in the voltaic cell
you are literally re-charging the voltaic cell with a new supply of electrons

77. Homework:
ELECTROLYTIC CELLS WORKSHEET
Chapter A2 Quiz:
DATE:

78. Reduction and Oxidation of Metals
Chapter A2
Reduction and Oxidation of Metals

79. 2.1 - Compounds and Chemical Change
Chapter A2
2.1 - Compounds and Chemical Change

80. Metallurgy
Archeologist studies show that ancient civilizations were not only using metals, they learned to manipulate it
metallurgy – heating and hammering copper into shapes
by analyzing copper artifacts, archeologists can use chemical analysis to track trade routes, locations of ancient industry, etc.

81. Tarnishing
the colour change, and loss of luster, associated with metal over time
silver tends to turn black over time
copper tends to turn green

82. Silver tarnish
when a metal tarnishes, a chemical reaction is occurring between the metal, and molecules in the air
balance the following reaction between pure silver, and hydrogen sulfide and oxygen in the air
__Ag(s) + __H2S(g) + __O2(g)  __Ag2S(s) + __ H2O(l)