Electrode Potentials and Fuel Cells

In a redox reaction, both reduction and oxidation occur

Redox reactions occur when electrons are transferred

Oxidation is a loss of electrons or an increase in oxidation number

Reduction is a gain of electrons or a decrease in oxidation number

The substance that is reduced is an oxidising agent

The substance that is oxidised is a reducing agent

In the oxidation half-equation, the electrons are lost so will be on the right

In the reduction half-equation, the electrons are gained so will be on the left

You can identify the oxidation and reduction processes using the oxidation numbers

Changes in the oxidation number of each atom in the equation show what is reduced and what is oxidised

When you have combined two half-equations and are short of 'H' on one side you can add H+, OH- or H2O:

• If the reaction has been carried out in acid conditions, add the number of H+ ions that are needed
• If the reaction has been carried out in alkaline conditions, add OH- ions
• If 'H' and 'O' are needed, add H2O

An electrochemical cell controls the electron ransfer in redox reactions to produce electrical energy

A half cell is made up of an element in two oxidation states. The simplest half cell has a ***** of metal placed in a solution of its aqueous ions. For example, a copper half cell is made up of a ***** of copper (oxidation state 0) placed in a solution of Cu2+ ions (oxidation state 2+).

An equilibrium exists at the surface of the copper between the oxidation states:

The forward reaction involves gain of electrons  and so is a reduction

The reverse reaction involves loss of electrons and so is an oxidation

The electrons are usually on the left-hand side. The electrode potential of the half cell indicates its tendency to lose or gain electrons 

Contains ions of the same element in different oxidation states.

Fe3+(aq) and Fe2+(aq)

A standard Fe2+(aq)/Fe3+(aq) half cell is made up of:

• A solution containing Fe2+(aq) and Fe3+(aq) ions with the same concentrations (equimolar)
• An inert platinum electrode to allow the transfer of electrons into and out of the half cell via a connecting wire

An electrochemical cell can be made by connecting two half cells with different electrode potentials

One half cell releases electrons, the other half-cell gains electrons

The two half cells are joined using a wire and a salt bridge

• The wire connects the metals. This allows electron transfer to take place. A voltmeter is connected to the wire. It has a high resistance and is used to minimise the current that flows.
• The salt bridge connects the two solutions and allows ions to be transferred between them. The salt bridge can be a ***** of filter paper soaked in an aqueous solution of an ionic compound that will not react with either solution (usually KNO3 or NH4NO3)

Hydrogen Gas in contact with H+(aq) ions

No electrode to connect the wire to.
Platinum electrode placed in the solution - in contact with H2(g) and H+(aq) ions

Platinum is inert. Allows transfer of electrons into and out of the half cell via a wire. The surface of the platinum is coated with platinum black - a spongy coating that allows the transfer of electrons between the non-metal and its ions.

A standard hydrogen half cell is used as the reference for measurement of voltages in electrochemical cells. A standard hydrogen half cell is made up of:

• Hydrochloric acid, HCl, conc 1 mol dm^-3, which is the source of H+ ions
• Hydrogen gas, H2(g), at 100kPa pressure
• Inert platinum electrode to allow electrons to pass into or out of the half-cell via a wire

The standard electrode potential of a half cell is the e.m.f of a half cell compared with a standard hydrogen half cell, measured ak 298K with solution concentrations of 1 mol dm^-3 and a gas pressure of 100 kPa

The electromotive force, e.m.f, is the voltage produced by a cell when no current flows

The standard electrode potential of a hydrogen half cell is 0V

A standard electrode potential is measured by connecting a standard half cell to a standard hydrogen half cell

The standard cell potential of cell is the e.m.f between the two half cells making up the cell under standard conditions - the difference between the standard electrode potentials of each half cell

The cell reaction is the overall chemical reaction occuring in the cell - the sum of the reduction and oxidation half reactions taking place in each half cell

Examples:

The Electrochemical Series

Electrode potentials and concentrations

The half-equation for the copper half-cell is:

From le Chatelier's principle, on increasing the concentration of Cu2+(aq):

• the equilibrium opposes the change by shifting to the right 
• electrons are removed from the equilibrium
• the electrode potential becomes less negative or more positive

A change in electrod potential resulting from concentration changes means predictions based on a standard value may not be valid

Sources of electrical energy
Two redox equilibria with different electrode potentials
Zinc and Copper: 

1. Non-rechargeable cells - provide electrical energy until the chemicals have reacted to the extent that the voltage falls. The cell is then 'flat' and is discarded.

2. Rechargeable cells - the chemicals in the cell react, providing electrical energy. The cell reaction is reversed during charging -  the chemicals in the cell are regenerated and the cell can be used again.

• nickel and cadium (NI-Cad) batteries, used in rechargeable batteries
• lithium-ion and lithium polymer batteries, used in laptops

3. Fuel Cells - the cell reaction uses external supplies of a fuel and an oxidant, which are consumed and need to be provided continuously. The cell will continue to provide electrical energy as long as there is a supply of fuel and oxidant

The hydrogen-oxygen fuel cell

A fuel cell uses energy from the reaction of a fuel with oxygen to create a voltage

• Reactants flow in, products flow out, electrolyte remains in the cell
• Can operate virtually continuously as long as the fuel and oxygen continue to flow into the cell
• Do not need to be recharged

Hydrogen-oxygen fuel cell with an alkaline electrolyte:

Fuel Cell Vehicles (FCVs) use hydrogen gas or hydrogen-rich fuels (methanol, natural gas, petrol)

Fuel mixed with water and converted to hydrogen gas by an onboard 'reformer' which operates at 250-300 degrees celsius

For example:

Fuel cells developed that use methanol as the fuel instead of hydrogen

• liquid fuel is easier to store
• methanol can be generated from biomass

Only generate a small amount of power. CO2 is a waste emission

Less pollution & less CO2

• Combustion of hydrocarbon fuels produces CO2 which contributes to the greenhouse effect. Incomplete combustion produces toxic carbon monoxide, which must be used by catalytic converters
• Hydrogen-rich fuels produce only small amounts of CO2 and air pollutants

Greater efficiency

• Petrol engines are less than 20% efficient in converting chemical energy by combustion of petrol. Most of the energy is wasted as heat.
• Hydrogen fuel cell vehicles are 40-60% efficient in converting the fuel's energy. Fuel consumption drops by more than half

As a liquid under pressure

• Very low temperature required
• Stored in a giant 'thermos flask' to prevent boiling

Adsorbed onto the surface of a solid material

• Similar to gases adsorbed to a catalyst
• H2 molecules attach to the surface of a material

Absorbed within some solids

• H2 molecules dissociate into H atoms which are incorporated as hydrides within a solid lattice

• Large-scale storage and transportation
• Feasibility of storing a pressurised liquid
• 'Adsorbers' and 'absorbers' have a limited lifetime
• Fuel cells have a limited lifetime - require regular replacement and disposal, high production costs
• Toxic chemicals used in the production of fuel cells

• Use of hydrogen must be accepted by politicians and by the public
• Logistical problems in the handling and maintenance of hydrogen systems 
• Hydrogen is an 'energy carrier', not an energy source. It must be manufactured by electrolysis of water or by reacting methane (finite fuel) with steam. More energy may be used making he hydrogen than is saved by using it. One solution could be to use renewable sources to generate the hydrogen.