Galvanic Cells From Chemistry to Electricity
Electrolytic Cells From Chemistry to Electricity . . . And back again!
Zn2+ Zn2+ Zn2+ 2e- Recall: the Galvanic Cell Cu2+ Cu(s) Zn(s) Cu2+ Cu2+
Load Salt Bridge Electrolyte Cu2+ Zn2+ Zn2+ Zn2+ Cu(s) Recall: the Galvanic Cell Current → Conductor Cu(s) Zn(s) Anode Cathode Cu2+ Oxidation happens here Reduction happens here
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
e- e- Analogy Lithium Potassium Barium Calcium Sodium Magnesium Aluminium Zinc Iron Nickel Lead (Hydrogen) Copper Silver Gold This activity series is the inverse of the reduction table Potential Difference
Question: • Can we push the electrons back up again?
- + Na+ Na+ Cl- Cl- The Electrolytic Cell External Voltage Anode Cathode Electrolyte: eg. NaCl
- + Na(s) Na(s) Cl2(g) The Electrolytic Cell External Voltage Anode Cathode Electrolyte: eg. NaCl
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 Similarities and Differences
Etymology • Electrolysis comes from two Greek words: electron and lysis (meaning to break.) Therefore, the word means “breaking apart using electrons
Water Reaction 2 H2O(l)→ O2(g) + 4 H+(aq) + 4 e- 2 H2O(l) + 2 e-→ H2(g) + 2 OH-(aq)
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)
- + O2(g) H2(g) The Electrolytic Cell External Voltage Anode Cathode Electrolyte
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 abuse Disadvantages • Low energy density • Memory effect • Contains toxic metals Popular Uses • Two-way radios, power tools, medical equipment
Lithium-Ion Advantages • High energy density • No memory effect Disadvantages • More expensive than Ni-Cd • Not fully mature, technology is still evolving Popular Uses • Cell-phones, iPods, laptop computers
Lead-Acid Advantages • Mature technology, well understood • Cheap and easy to manufacture • No memory effect Disadvantages • Very low energy density; most applications require huge batteries • Limited number of full discharge cycles • Environmental concerns Popular Uses • Electric cars, golf carts, scooters
Reusable Alkaline Advantages • Cheap to manufacture • More economical than primary alkaline cells Disadvantages • 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
Fuel Cells Behind the hype
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.
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 O2(g) H2(g) 2 H+ 2 e- Anode Cathode Electrolyte
The Promise of Hydrogen • Many Fuel Cells are emission-free because they run on hydrogen H2(g) H2O(l) Anode Cathode Electrolyte
Overall Reaction: H2 + O2→ H2O Question: Where does hydrogen come from?
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
Rust The Silent Killer
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
Particle The Rust Galvanic Cell Oxidation: Fe(s) Fe2+ + 2 e- Reduction: O2(g) + 2H2O + 4e- 4 OH- O2(g) 2 H2O(l) Fe(s)
Fe2+ 2 e- The Rust Galvanic Cell Oxidation: Fe(s) Fe2+ + 2 e- Reduction: O2(g) + 2H2O + 4e- 4 OH- O2(g) Particle 2 H2O(l) Fe(s)
The Rust Galvanic Cell Oxidation: Fe(s) Fe2+ + 2 e- Reduction: O2(g) + 2H2O + 4e- 4 OH- 2 Fe(s) +O2(g) + 2H2O 2 Fe2+ + 4 OH- Particle 4 OH- Fe2+ Fe(s)
The Rust Galvanic Cell Particle 4 OH- 2 Fe2+ Fe(s)
The Rust Galvanic Cell 2 Fe2+ + 4 OH- 2Fe(OH)2 Particle 2 Fe(OH)2 Fe(s)
4 e- The Rust Galvanic Cell 4 Fe(OH)2 + O2(g) + 2 H2O(l) 4 Fe(OH)3 O2(g) Particle 4 Fe(OH)2 2 H2O(l) Fe(s)
The Rust Galvanic Cell 4 Fe(OH)2 + O2(g) + 2 H2O(l) 4 Fe(OH)3 4 OH- Particle 4 Fe(OH)2 Fe(s)
The Rust Galvanic Cell 4 Fe(OH)2 + O2(g) + 2 H2O(l) 4 Fe(OH)3 Particle 4 Fe(OH)3 Fe(s)
The Rust Galvanic Cell 4 Fe(OH)2 + O2(g) + 2 H2O(l) 4 Fe(OH)3 Fe(OH)3 Fe2O3·3 H2O 4 Fe2O3· 3 H2O Fe(s)
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?