ELECTRODE POTENTIALS AND CELLS

ELECTROCHEMICAL CELLS can be made from

2 different metals dipped in salt solutions of their own ions and connected by a wire (AKA the external circuit)

there are 2 reactions within an electrochemical cell

oxidation and reduction- this process is a REDOX

an electrochemical cell made of copper and zinc

zinc loses electrons more easily than copper, so at the ZINC electrode there is OXIDATION to form Zn²⁺

electrons from the zinc electrode are released into the external circuit into the copper electrode

reducing the Cu²⁺ ions into Copper atoms

electrons flow though the wire FROM the most reactive metal TO the least

the voltage between two half-cells can be calculated using a voltmeter to find the CELL POTENTIAL/ EMF

cathode charge

negative

anode charge

positive

reaction at cathode

reduction

reaction at anode

oxidation

you can have half cells involving solutions of 2 aqueous ions of the same element

for example using Fe²⁺ and Fe³⁺

platinum is used because

its inert and conducts electricity

the conversion of Fe²⁺ to Fe³⁺

occurs on the surface of the platinum electrode

the direction of the conversion of 2+ to 3+ depends on the other half-cell in the circuit

• if it contains a metal LESS reactive than iron, Fe2+ will be oxidised to Fe3+

• if the other cell contains MORE reactive metal fe3+ will be reduced to fe2+ at the electrode

the direction of reaction depends on the metal able to lose electrons more easily (HOW EASILY OXIDISED)

How easily a metal oxidises is measured using electrode potentials

easily oxidised metals have VERY NEGATIVE potentials

easily reduced metals

have less negative/ positive electrode potentials

MORE NEGATIVE FIRST MORE POSITIVE SECOND

cell potential is ALWAYS positive

more positive - more negative

factors affecting the electrode potential

• temperature 25℃

• pressure 1atm

• concentration 1 mol dm^-3

changing the equilibrium position

changes the cell potential

standard conditions used to measure electrode potentials

so you get the same value for the electrode potential and can compare values for different cells

THE STANDARD HYDROGEN ELECTRODE

you measure the electrode potential of a half cell against the standard hydrogen electrode

in the standard hydrogen electrode, hydrogen gas bubbled through a solution of aqueous H+ ions

a platinum electrode is used as a platform for the oxidation/reduction reactions

standard conditions when measuring electrode potentials using SHE

• all solution ceoncentrations 1.00 mol md-3

• 298K

• 100kPa

Standard electrode potential of a half cell is the voltage measured under standard conditions when

the half cell is connected to s SHE

SHE is always on the left

no matter of which half cell is more positive

SHE can be used to calculate standard electrode potentials because

the electrode potential of a SHE is 0.00V therefore the voltage will be equal to the electrode potential on the RIGHT

electrochemical series

long list of electrode potentials for different electrochemical half cells in order

half cells are always written as reduction reactions (GAIN ELECTRONS)

but reactions are reversible and can go the opposite way

when two half equations are put together in an electrochemical cell

• the one with the more negative electrode potential goes in the direction of oxidation BACKWARDS

• the more positive electrode potential goes in the direction of reduction FORWARDS

you can use electrode potentials to predict whether a redox reaction will occur and show which direction it goes in

• find the 2 half equations for redox reactions and write them BOTH as reduction reactions (GAIN ELECTRONS)

• Use the electrochemical series to work out which half equation has more negative electrode potential

• write out the half equation with the more negative potential going in the backwards direction OXIDATION and the more positive potential going forwards REDUCTION

• combine the 2 half equations and write out a full redox equation

when the overall electrode potential is positive

the reaction is feasible in that direction

batteries types

• rechargeable

• non rechargeable

non rechargeable cells

use irreversible reactions

common type of non rechargeable cells

dry cell alkaline battery

short use gadgets that dont require alot of power use dry cell alkaline batteries:

• TV remotes

• smoke alarms

• torch

Zinc-Carbon dry cells batteries

negative electrode: zinc

Positive electrode: Manganese dioxide and carbon mixture

Electrolyte: Ammonium chloride paste

zinc electrode forms the casting of the battery which becomes thinner as the zinc is oxidised

which is why overtime batteries leak

devices with rechargeable batteries:

• mobile phones

• laptops

• cars

common type of rechargeable battery

lithium cells

lithium cell batteries

Positive electrode: Lithium cobalt oxide (LiCoO₂)

Negative electrode: graphite, where lithium ion reacts

electrolyte: lithium salt in organic solvent

half equations of lithium battery

negative electrode: Li → Li⁺ + e^-

positive electrode: Li⁺ + CoO2 +e^- → Li⁺[CoO2]^-

there are 2 other types of rechargeable battery

• NiCad (nickel-cadmium)

• Lead-acid

for rechargeable batteries

a current is supplied to force electrons to flow in the opposite direction in the circuit which reverses the reaction

rechargeable reactions only work because

none of the substances in a rechargeable battery escape or get used up

the reactions that occur in non rechargeable batteries

are difficult or impossible to reverse in this way

for most cells

chemicals that generate the electricity are contained in the electrodes and electrolyte that form the cell

for fuel cells

chemicals are stored separately outside the cell and are fed in when electricity is required

alkaline-hydrogen fuel cell powers electric vehicles, contain platinum electrodes

• hydrogen and oxygen gases are fed into 2 seperate platinum-containing electrodes

• these electrodes are made up by coating a porous ceramic material with a thin layer of platinum

• this is cheaper than solid platinum rods and providesa larger surface area for a faster reaction

alkaline-hydrogen fuel cell, NEGATIVE ELECTRODE

• the electrodes are separated by an anion-exchange membrane that allows anion (OH-) and water to pass through it

• however it stopes hydrogen and oxygen gas from passing through it

• the electrolyte is an aqueous alkaline KOH solution

• hydrogen is fed to a negative electrode: H₂ +2OH^- → 2H₂O + 2e^-

alkaline-hydrogen fuel cell

• electrons flow from the negative electrode through an external circuit to the positve electrode

• OH- pass through the anion-exchange membrane towards the negative electrode.

• oxygen is fed to the positive electrode: O₂ + 2H₂O + 4e^- → 4OH^-

alkaline-hydrogen fuel cell overall reaction

2H₂ + O₂ → 2H₂O

advantages of fuel cells in cars

• fuel cells are more efficient, they convert more of their available energy into kinetic energy to get the car moving

• internal combustion engines waste a lot of their energy producing heat

• the only waste product of fuel cell is water, no toxic chemicals or CO2

• fuel cells don’t need to be recharged like batteries

• as long as hydrogen and oxygen are supplied the cell will continue to produce electricity

disadvantages of fuel cells

• you need energy to produce a supply of hydrogen, produced from electrolysis of water but this requires electrocity thats normally generated from the burning of fossil fuels- this whole process is NOT carbon neutral

• hydrogen is also highly flammable so needs to be handled carefully when stored and transported

• infrastructure to prodive hydrogen fuel for cars doesnt exist on a large scale so refuling stations are very rare