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Recall: the Galvanic Cell
Cu2+
Cu(s)
Zn(s)
Conductor
Zn2+
Zn2+
Cu(s)
Cu2+
Current →
Load
Oxidation happens here Reduction happens here
AnodeCathode
Electrolyte
Zn2+
Salt Bridge
Questions:
• Which way did the electrons go?
• Why?
• What happens when we use a different pair of metals?
Voltage
• Voltage is also known as electromotive force and potential difference
• It is a measure of how much energy electrons have to get them moving
• It is related to the distance between two metals are on the activity series
AnalogyLithium
Potassium
Barium
Calcium
Sodium
Magnesium
Aluminium
Zinc
Iron
Nickel
Lead
(Hydrogen)
Copper
Silver
Gold
e-
e-
Potential Difference
This activity series is the inverse of the reduction table
Similarities and Differences
Galvanic Cells• Oxidation occurs at the
anode• Reduction occurs at the
cathode• Anode is negative• Cathode is positive• Does not require external
voltage source• Changes chemical
reactions into electrical energy
Electrolytic Cells• Oxidation occurs at the
anode• Reduction occurs at the
cathode• Anode is positive• Cathode is negative• Requires external voltage
source• Changes electrical
energy into chemical reactions
Etymology
• Electrolysis comes from two Greek words: electron and lysis (meaning to break.) Therefore, the word means “breaking apart using electrons
Water Reaction
Oxidation at the Anode: 2 H2O(l) → O2(g) + 4 H+(aq) + 4 e-
Reduction at the Cathode: 4 H2O(l) + 4 e- → 2 H2(g) + 4 OH-(aq)
6 H2O(l) → 2 H2(g) + O2(g) + 4 H+ + 4 OH-(aq)
Practice
• Where would the hydrolysis reaction be useful?
• Draw an electrolytic cell for AgBr(l)?
• Draw a galvanic cell for the reaction
Au+3 (aq) + Ag(s) Ag+1 (aq) + Au (s)
Review of Batteries
• Define the following terms:
• Primary Cell
• Secondary Cell
• Power Density
• Memory Effect
Nickel-Cadmium
Advantages• Can be recharged 1000 times
or more• Low cost/cycle• Tough, stands up to abuseDisadvantages• Low energy density• Memory effect• Contains toxic metalsPopular Uses• Two-way radios, power tools,
medical equipment
Lithium-Ion
Advantages• High energy density• No memory effectDisadvantages• More expensive than Ni-Cd• Not fully mature, technology is
still evolvingPopular Uses• Cell-phones, iPods, laptop
computers
Lead-Acid
Advantages• Mature technology, well
understood• Cheap and easy to manufacture• No memory effectDisadvantages• Very low energy density; most
applications require huge batteries
• Limited number of full discharge cycles
• Environmental concernsPopular Uses• Electric cars, golf carts,
scooters
Reusable Alkaline
Advantages• Cheap to manufacture• More economical than primary
alkaline cellsDisadvantages• Limited current, cannot be
made on large scale• Limited cycle life (about 10
cycles); fully discharging shortens life
Popular Uses• Personal CD players, radios,
flashlights
The Limit of Batteries
• A battery is a fancy type of galvanic cell. It changes chemicals into electricity.
• Eventually, all the chemical are reacted and the battery goes dead.
• If the battery is a secondary cell, you can recharge it, but this takes time and energy. Also, there is a limited number of times you can do this.
• Wouldn’t it be nice to be able to just open up a battery, and pour in some more chemicals, like refueling a car?
Fuel Cells: The Ultimate Battery
• A fuel cell is a type of galvanic cell that allows you to add fresh chemicals continuously. It will continue to run as long as you keep adding fuel.
Fuel Cell Comparisons
Battery Fuel Cell Engine
Energy Conversion
Chemical>
Electrical
Chemical>
Electrical
Chemical>
Mechanical>
Electrical
Fuel Zinc or other metal
H2, Methane, Propane, Methanol, etc.
Propane, Methane, Gasoline
Powered By Electrochemical Reaction
Electrochemical reaction
Combustion Reaction
Power Output Low Variable High
Advantages of Fuel Cells
• Extremely versatile – can power everything from cell phones to buses
• Can run on a variety of fuels• More environmentally friendly than combustion
Disadvantages of Fuel Cells
• Technology is still somewhat unreliable
• Some types still produce greenhouse gas emissions
• EXPENSIVE
The Promise of Hydrogen
• Many Fuel Cells are emission-free because they run on hydrogen
H2(g)
Anode CathodeElectrolyte
2 H+
O2(g)
2 e-
The Promise of Hydrogen
• Many Fuel Cells are emission-free because they run on hydrogen
H2(g)
Anode CathodeElectrolyte
H2O(l)
Types of Fuel Cells
• Proton Exchange Membrane (PEM)
• Solid Oxide Fuel Cells (SOFC)
• Alkaline Fuel Cells (AFC)
• Direct Methanol Fuel Cells (DMFC)
• There are many more, but we won’t get into them here
Your Task
• You will work in groups of about 5• Half of each group will argue “for” a
particular type of fuel cell the other half will argue “against”
• Each half-group will prepare an extremely short (60 s) presentation to convince the audience of their stance
• The class will vote on who was most convincing
What if there was a chemical reaction that:
• Turned vehicles and buildings into dust
• Caused billions of dollars worth of damage per year
• Was virtually unstoppable
• Had the potential to destroy an entire planet’s atmosphere
Why does rust happen?
• Iron, like most metals, is a strong reducing agent
• Earth’s atmosphere is 21% O2, which is a powerful oxidizing agent
• Galvanic cells are easy to set up, and can be as simple as a drop of water
The Rust Galvanic Cell
Fe(s)
2 H2O(l)Particle
O2(g)
Oxidation: Fe(s) Fe2+ + 2 e-
Reduction: O2(g) + 2H2O + 4e- 4 OH-
The Rust Galvanic Cell
Fe(s)
2 H2O(l)Particle
O2(g)
Oxidation: Fe(s) Fe2+ + 2 e-
Reduction: O2(g) + 2H2O + 4e- 4 OH-
2 e-Fe2+
The Rust Galvanic Cell
Fe(s)
Particle
Oxidation: Fe(s) Fe2+ + 2 e-
Reduction: O2(g) + 2H2O + 4e- 4 OH-
Fe2+4 OH-
2 Fe(s) + O2(g) + 2H2O 2 Fe2+ + 4 OH-
The Rust Galvanic Cell
Fe(s)
Particle4 Fe(OH)24 e-
O2(g)
2 H2O(l)
4 Fe(OH)2 + O2(g) + 2 H2O(l) 4 Fe(OH)3
The Rust Galvanic Cell
Fe(s)
4 Fe2O3 · 3 H2O
4 Fe(OH)2 + O2(g) + 2 H2O(l) 4 Fe(OH)3
Fe(OH)3 Fe2O3·3 H2O
Questions
• How did the water become an electrolyte?• What was the anode?• What was the cathode?• Would this happen for other metals? Which
ones? How would it be different?• Corrosion costs billions of dollars a year in
damage as boats, cars, trains, building, etc. all gradually turn to dust. What can we do to prevent rusting from causing so much damage?
Rust Prevention: Protective layer
• Adding a protective layer of paint, plastic, or glass prevents the iron from coming in contact with the electrolyte
• What happens if the protective layer develops a scratch?
Rust Prevention: Galvanizing
• If you coat iron in a thin layer of zinc, it is called galvanization. The layer both protects the iron and will act as the anode if a scratch develops
• What happens when all the zinc is oxidized?
Fe2O3 + 3 Zn 3 ZnO + 2 Fe
Rust Prevention: Sacrificial Anode
• Some ships and gas pipelines are protected by putting a block of zinc, aluminum, or magnesium on them. The more reactive metal is oxidized and the iron stays intact.
• Who pays to replace the sacrificial anode every year?
Rust on Mars
• The surface of mars is completely covered in rust. Scientists think that Mars might once have had an atmosphere like earth’s, but all of that oxygen is now tied up in Fe2O3.
• Question: why hasn’t this happened on Earth?
Practice Questions
• In 2000, Transport Canada recalled thousands of cars with corroded engine mounts in Nova Scotia, New Brunswick, and PEI. Why was corrosion such a problem in these provinces?
• A small scratch in a car door can quickly develop into a major rust spot. Why does this happen?
• Does acid rain promote or prevent corrosion? Explain?