Chapter 8: Energy from Electron Transfer

A battery is a system for the direct conversion of chemical energy to electrical energy.

Batteries are found everywhere in today’s society because they are convenient, transportable sources of stored energy.

The “batteries” shown here are more correctly called galvanic cells.

A series of galvanic cells that are wired together constitutes a true battery – like the one in your car.

8.1

A galvanic cell is a device that converts the energy released in a spontaneous chemical reaction into electrical energy.

This is accomplished by the transfer of electrons from one substance to another.

Consider a nickel-cadmium (NiCad) battery:

Electrons are transferred from cadmium to nickel.

8.1

Chemical to electrical energy

The electron transfer process involves two changes: the cadmium is oxidized, and the nickel is reduced.

Each process is expressed as a half-reaction:

Oxidation half-reaction: Cd Cd2+ + 2 e–

Reduction half-reaction: 2 Ni3+ + 2 e– 2 Ni2+

Overall cell reaction: 2 Ni3+ + Cd 2 Ni2+ + Cd2+

The cadmium releases two electrons, resulting in a 2+ ion.

Two Ni3+ ions accept the two electrons and their respective charges go from 3+ to 2+.

8.1

The overall cell reaction does not have any electrons written in it – they must cancel.

Oxidation is loss of electrons: Cd Cd2+ + 2 e–

Cd loses electrons

Reduction is gain of electrons: 2 Ni3+ + 2 e– 2 Ni2+

Ni3+ gains electrons

Oil Rig is a useful mnemonic device.

The transfer of electrons through an external circuit produces electricity, the flow of electrons from one region to another that is driven by a difference in potential energy.

8.1

To enable this transfer, electrodes (electrical conductors) are placed in the cell as sites for chemical reactions.

Oxidation occurs at the anode

Reduction occurs at the cathodeThe cathode receives the electrons.

The difference in electrochemical potential between the two electrodes is the voltage (units are in volts).

8.1

A Laboratory Galvanic Cell

8.1

Oxidation (at anode):Zn(S) Zn2+(aq)

Reduction (at cathode):Cu2+(S) Zn2+(aq)

An Alkaline Cell

8.2

8.2

– Mercury batteries had the advantage of being very small – Used in watches, cameras, hearing aids – Toxicity and disposal concerns led to alternatives

8.2

Lead–Acid Storage BatteriesThis is a true battery as it consists of a series of six cells.

Anode = PbCathode = PbO2

Pb(s) + PbO2(s) + 2 H2SO4(aq) 2 PbSO4(s) + 2 H2O(l)discharging

recharging

Rechargeable:

Hybrid Electric Vehicles (HEVs)

8.4

Combining the use of gasoline and battery technology

Many hybrids, unlike conventionalgasoline-powered cars, deliverbetter mileage in city driving thanat highway speeds.

8.5

Honda FCX Powered by Fuel Cells

8.5

A fuel cell is a galvanic cell that produces electricity by converting the chemical energy of a fuel directly into electricity without burning the fuel.

Both fuel and oxidizer must constantly flow into the cell to continue the chemical reaction.

Chemical Changes in a Fuel CellH2 + ½ O2 H2O

Anode reaction: (Oxidation half-reaction)H2(g) 2 H+(aq) + 2 e–

Cathode reaction: (Reduction half-reaction) ½ O2(g) + 2 H+(aq) + 2 e– H2O(l)

H2(g) + ½ O2(g) + 2 H+(aq) + 2 e– 2 H+(aq) + 2 e– + H2O(l)

Overall (sum of the half-reactions)

H2(g) + ½ O2(g) H2O(l)

Net equation:

8.5

8.5

Comparing Combustion with Fuel Cell Technology

8.5

Methanol or Methane Gas Could be Converted to Hydrogen Gas Using a Reformer

8.6

High Pressure Hydrogen Storage

Hydrogen Storage in Metal Hydrides

Refueling a Honda FCX Clarity with hydrogen gas

Hydrogen is absorbed onto metal hydrides and released with increased temperature or decreased

pressure.

8.6

If the “Hydrogen Economy” becomes reality, where will we get the H2?

One potential source is fossil fuels:

165 kJ + CH4(g) + 2 H2O(g) → 4 H2(g) + CO2(g)

247 kJ + CO2(g) + CH4(g) → 2 H2(g) + CO(g)

But these reactions produce carbon dioxide and carbon monoxide.

Electrolysis of water could also be used:

249 kJ + H2O(g) → H2(g) + ½ O2(g)

8.6

Hydrogen (and oxygen) gas produced by the electrolysis of water.

Electrolysis of Water:

But this process requires energy.

8.6

• Using fossil fuels or electrolysis uses a lot of energy.

• Instead, use energy from the sun for this process.

Use Energy from the Sun to Split Water:

Photovoltaic Cells

8.7

Energy of the future? An awesome source of energy.

Our society has been using photovoltaic cells (solar cells) at a minimum level, but will we all begin picking up on this natural form of energy as a way to conserve our natural coal and petroleum supply?

8.7

Other countries making use of solar energy

Solar Park Gut Erlasee in Bavaria. At peak capacity, it can generate 12 MW.

Harnessing the energy of the sun for pumping water

It’s time to make important decisions and advances in alternative energy technologyand new sources of renewable energy.

8.7

How does a photovoltaic cell generate electricity?

• A Semiconductor can be used (usually made of crystalline silicon). Semiconductors have a limited ability to conduct electricity.

• Two layers of semiconducting materials are placed in direct contact to induce a voltage in the PV cell.

• n-type semiconductor – has an abundance of electrons• p-type semiconductor – deficient in electrons

Arsenic-doped n-type silicon semiconductor

Gallium-doped p-type silicon semiconductor

8.7

Schematic diagram of a solar cell:

8.8

Other Renewable Energy Sources: Solar Thermal

Aerial view of the Solar Millennium Andasol project in

Spain

Close-up of a portion of the mirrored array

Other renewable energy sources include wind, water, and geothermal energy.