Chapter 23 - Redox + Electrode Potentials

oxidising agent

species that is reduced in reaction causing other species to be oxidised - takes electrons

reducing agent

species that is oxidised in reaction causing other species to be reduced - adds electrons

writing redox equations from HALF EQUATIONS

-multiply half equations if needed so same number of electrons in both

-then cancel out electrons and combine equations to form redox equation

reduction half + oxidation half = redox equation

if OH^- and H₂O in half equation

subtract them from each other

e.g. 10OH^- - 6OH^- = 4OH^- , 5H₂O - 3H₂O = 2H₂O

writing redox equations from OXIDATION NUMBERS

1) write equation using species mentioned in question

2) then balance any atoms that are changing oxidation number first

3) then write out reduction equation and oxidation equation and multiply if needed so same number of electrons

4) then write out overall equation and to balance H and O, add H⁺, OH^-, H₂O

5) check if charge on left = charge on right

half cell

comprises of an element in two different oxidation states

types of half cells 1) METAL/METAL ION (draw)

metal element placed in an aqueous solution of its ions

e.g. Zn²⁺_(aq) + 2e^- ⇌ Zn_(s)

forward = reduction, backward = oxidation

types of half cells 2) METAL ION/METAL ION (draw)

ions of the same element in different oxidation states

-consists of solution with same concentration of each ion, inert platinum electrode

types of half cells 3) NON-METAL/NON-METAL ION (draw) = hydrogen half cell

standard cell made up of:

-1mol dm^-3 HCl acid

-H₂ gas at 100kPa pressure

-temperature of 298K

-inert platinum electrode

electrochemical cells structure (DRAW)

two half cells connected together

-electrodes connected by a wire from voltage source = voltmeter allows flow of electrons carrying charge

-has salt bridge between = connects solutions together so charge of ions can be transferred between half cells

standard electrode potential E^o

the emf (electromotive force) of a half cell compared with a standard hydrogen cell (0V) measured at 298K with solution concentration of 1mol dm^-3 and gas pressure of 100kPa

electrode potential value

if more positive, the greater tendency for species within to gain electrons so reduction reaction favoured so is a stronger oxidising agent

VICE VERSA for more negative

calculating E^o of a cell

E^o (more positive) - E^o (more negative)

if E^o is positive then reaction is feasible

reasons why redox reactions may not occur even if feasible from E^o value

-activation energy is too high so rate of reaction is too slow

-reaction does not occur under standard conditions

-reaction does not occur under aqueous conditions

-standard cell potential is too small → E^o < 0.4V

types of electrochemical cells 1) non-rechargeable cells

provide electrical energy until all chemicals have reacted = only one use

-used in low current, long storage devices

types of electrochemical cells 2) rechargeable cells

provide electrical energy until all chemicals have reacted but can be regenerated when recharged = flow of electrons reversed so reactions will be reversed

common types: Ni + Cd batteries, Li ion batteries → portable/light density but are unstable at high temps

types of electrochemical cells 3) fuel cells

converts chemical energy to electrical energy from reaction of a fuel with oxygen to create a voltage

-however fuel + oxygen must be continuously supplied

EXAMPLE of fuel cell = hydrogen fuel cell

overall equation = ½O₂ + H₂ → H₂O = +1.23V

alkaline: ox = H₂ + 2OH^- → 2H₂O +2e^- , re = ½O₂ + H₂O + 2e^- → 2OH^-

acidic: ox = H₂ → 2H⁺ + 2e^- , re = ½ O₂ + 2H⁺ + 2e^- → H₂O

advantages of using methanol fuel cell over hydrogen

-less volatile

-easier to store + transport

-can be sourced from renewable sources e,g, biomass