Fuel cells and Batteries

primary cell

• non re chargeable - due to the products slowly migrating away from the electrodes or are consumed by side reactions occurring in cell

• alkaline cells

• they contain no fluids

• simple and light

• cheap

• short lasting

• lower self discharge

• higher energy density

Properties of secondary cells

• Rechargeable- done by reversing reaction through attaching cell to an electrical charge which has a potential difference a little greater than that of the cell. Positive electrode of the charger to positive electrode of cell and negative electrode of the charger to negative electrode.

• electrical energy is converted to chemical in cell, in order for it to work products formed must stay in contact with electrodes

• These are made up of wet cells (flooded and liquid cells) and molten salt (liquid cells with different composition)

• complex and heavy

• expensive

• long lasting

• can withstand higher electrical currents

energy transformation in secondary cells

• when cell discharges it acts as a galvanic cell converting chemical energy into electrical

• when cell is recharged it acts as an electrolytic cell converting electrical to chemical energy

properties of fuel cells

• cells that are constructed with a continuous flow of reactants allowing for a constant production of chemical energy

• typically use hydrogen and oxygen gases as fuel and produce water

• transform chemical energy into electrical energy

• continuous electricity

• efficient

• a fuel cell using hydrogen produces electricity, water, heat and small amounts of nitrogen dioxide

anode of a leclanche cell

zinc case: Zn → Zn2+ + 2e

cathode of leclanche cell

carbon rod surrounded by paste: 2MnO2 + 2NH4+ + 2e → Mn2O3 + 2NH3 +H2O

lead-acid cell anode and cathode

anode: Made of lead and surrounded by sulfuric acid Pb + SO42- → PbSO4 +2e

Cathode: made of solid lead oxide: PbO2 + 4H+ + SO42- +2e → PbSO4 + 2H2O + 2e

wet corrosion

• at anodic site, electrons are released when iron atoms are oxidised to iron (II) ions

• at cathodic site the electrons reduce oxygen gas and water to produce hydroxide ions

• these combine to form insoluble iron (II) hyrdoxide

• iron (II) hydroxide is readily oxidised by more oxygen in the air to form iron (III) hydroxide

• Iron (III) hydroxide dehydrates to form rust, Fe2O3*H2O

types of methods to prevent corrosion

surface protection, galvanising, sacrificial anode and cathodic protection

Surface protection

• prevents iron from coming into contact with oxygen and water

• oil, grease, paint, plastic, other metals

• advantages: cheap and easy, can be flexible, effective for as long as coat remains in tact

• Disadvantages: coating must be reaplied when it deteriorates any scratch will leave metal susceptible

• For less reactive metal coating: any scratch will mean the more reactive iron will act as the anode and will corrode faster

Galvanising

• more reactive metal e.g. zinc is used as a coating

• more reactive metal will corrode quickly and form a thin layer of zinc oxide which protects the iron

• will protect the iron even if scratched since zinc is more reactive

Sacrificial anode

• force iron to become the cathode by electrically connecting it to a more reactive metal which will act as the anode

• Sacrificial anode will need to be replaced

Cathodic protection

• Use of sacrificial anode may be improved by use of a DC electrical current

• The (-) terminal of the DC power supply is electrically connected to iron making it the cathode of a galvanic cell

• the (+) terminal is connected to the sacrificial anode, which could be made of iron in this scenario but typically a more reactive metal

Electrolysis

Forcing a non-spontaneous to occur reaction by using electricity

Rechargeable batteries act as a voltaic cell when discharging and a electrolytic cell when recharging

under what conditions are standard reduction potentials measured

• solution concentration of 1.00 mol/L

• temp of 298 K (25C)

• pressure of 100 kPa

Why is the salt bridge needed?

it completes the circuit by allowing ions to move between the half cells

cell notation

anode on the left, cathode on the right

rules of standard reduction potentials

• reduction half equation must be above oxidation half on the table

• reduction is in the forward and oxidation in the reverse for reaction to occur

• voltage (E^0) must be positive for it to be spontaneous

Hydrogen half cell

How are the standard equation found 

by constructing a hydrogen half cell  with that element

electrolysis of water

• water can be split into hydrogen and oxygen gas

• small amount of sulphuric acid is added to increase the number of ions in the solution, since pure water is a poor conductor

• amount of Hydrogen gas produced is 2x the amount of Oxygen gas produced

• H2 is produced at the cathode and O2 is produced at the anode

what makes electrolysis of Aqueous solutions more complicated than molten

• water can be oxidised or reduced instead of the ions of the salt

electroplating

• electrolysis can be used to coat a thin layer of a metal over another metal

electrorefining of copper

electrolysis of aqueous solutions rules

• the oxidation reaction is above the reduction

• the oxidation reaction goes in reverse

• reduction reaction goes forward

types of voltaic cells

primary and secondary cells

dry cell (zinc-carbon cell or leclanche cell)

• primary cell

• the zinc case acts as the anode

• the cathode is a rod made from carbon surrounded by a paste of MnO2, ZnCl2, NH4Cl and powdered carbon

• initial overall voltage is 1.5 volts

Disadvantage of the dry cell

• since the zinc case takes part in the reaction, it will deteriorate and eventually leak

• also ammonium ions are weakly acidic which causes further deterioration of the zinc

• can’t be reversed because running the current through it in reverse will cause potentially dangerous side reactions which can produce gases (producing gases in a confined space will cause an explosion

alkaline battery

• primary cell

• Anode is a zinc rod: Zn + 2OH- → ZnO +H2O + 2e

• Cathode is the steel casing: 2MnO2 + H2O + 2e → Mn2O3 + 2OH

• similar voltage to dry cells but can sustain higher currents

• small and inexpensive

• cannot be recharged as products of the discharge reaction can move away from the electrodes and attempts to recharge can produce gases or cause the case to rupture

Lead - acid battery (car battery)

• secondary cell (can be recharged)

• made up of 6 cells, able to produce 2 volts each therefor the total volatge of the battery = 12 volts

• Anode is made up of lead surrounded by sulphuric acid: Pb + SO4²→ PbSO4 + 2 e

• Cathode is made of solid lead oxide: PbO2 + 4H+ + SO4² + 2 d → PbSO4 + 2 H2O

• because both reactions produce solid PbSO4 which remains on the electrodes, the battery can be effectively recharged by applying a voltage of at least 12v

• impractical due to their size and weight

• contains large amounts of lead and sulphuric acid, both of which are quite toxic, so they must be handled with care and disposed of properly

Lithium Ion Battery

• secondary cell (can be recharged)

• Anode is made from lithium ions attached to a thin layer of graphite. this is called intercalation 

• lithium ions (having a positive charge) are attached to the delocalised electrons in the graphite

• combination of lithium ions with atoms of carbon makes them behave like a lithium metal (a strong reducing agent) with a high tendency to donate electrons and become cations

• takes at least 6 carbon atoms per lithium ion for this to work perfectly 

• electrolyte is an organic solvent containing dissolved lithium ions

function: as battery discharges the lithium ions from the graphite anode move through the electrolyte to the cathode, releasing electrons through the external circuit and producing a voltage of about 3.7 v

Anode: LixC6 → X li+ + X e- + C6

Cathode: Li1-xCoO2 + XLi+ Xe- → LiCoO2

Recharge: exact opposite process occurs with electrons moving to graphite from the external circuit which attracts lithium ions back to graphite

Disadvantage: 

• expensive due to complexity and need for purity

solid oxide fuel cell

Anode: 2 H2 + 2 O2- → 2 H2O + 4 e

Cathode: O2 + 4e- → 2O²

disadvantage of fuel cells

• poor reliability over long periods of time

• expensive

• producing hydrogen gas requires a lot of electricity

• about 30% efficient

IPHE

Battery definition 

• two or more cells together is the traditional definition

dry corrosion

• example: the patina that forms on copper, copper reacts with the air to form copper oxide

• 2Cu(s) + O2(g) → 2CuO(s)