Ch.23 redox and electrode potentials

takes electrons from species being oxidised

adds electrons to the species being reduced

MnO4- ions are reduced to Mn2+ KMnO4 in burette (deep purple so read from top of meniscus) pipette measured amount of analysed solution into conical flask with an excess of HCl to provide H+ to act as reducing agent no indicator reaction is self indicating end point when first permanent pink colour indicating an excess of MnO4- ions

Fe2+ –> Fe3+

acidified potassium dichromate Cr2O72- (orange) —> Cr3+ (green)

IO3 - + 5I - + 6H + —> 3I2 + 3H2O I2 + 2S2O3 2- —> 2I - + S4O6 2- 2 S2O3 2- is oxidised to S4O6 2- I2 is reduced to 2I- add standard solution of Na2S2O3 to burette pipette analysing solution into conical flask add excess of potassium iodide oxidising agent reacts with iodine ions forming iodine making the solution yellow brown as titration occurs iodine is reduced back to I- colour fades when straw colour add starch indicator solution turns deep blue black colour disappears at end point

chlorate(1) ClO- ir reduced to Cl- copper(11) Cu2+ is reduced to CuI

copper salts can be disolved in water to produce Cu2+ Cu can be reacted with an acid to form Cu2+ copper alloys can be dissolved in conc nitric acid then neutralised to form Cu2+ Cu2+ react with I- forming I2 and white CuI precipitate mixture appears brown

converts chemical energy into electrical energy

contains chemical species present in redox half equation

metal rod dipped into a solution if its aqueous metal ions

ions of the same state but with different oxidation states platinum electrode is used

when both half cells are metal/metal ion the more reactive metal releases electrons more readily and is oxidised electrode with more reactive metal is oxidised - - electrode electrode with less reactive metal is oxidised - + electrode

E* measure of the tendency to be reduced and gain electrons e.m.f of a half cell connected to a standard hydrogen half cell under standard conditions E* of a standard hydrogen half cell is 0V

298K 25*C 100 kPa

half cell connected to standard H electrode connected by a wire to allow a controlled flow of e- 2 solutions connected by a salt bridge allowing ions to flow composed of an unreactive electrolyte solution filter paper soaked in KNO3 forward reaction on data tables is always reduction most - E* value at top

greater tendency to lose electrons less tendency to gain more reactive a metal is less reactive a non metal is stronger the reducing agent it is

E* cell = E(+) - E(-)

no indication of the rate of the reaction

non rechargeable

rechargeable

the stronger an oxidising agent it is greater tendency to gain electrons less tendency to lose electrons less reactive a metal is more reactive a non metal is

E* values with most - at top and most + at bottom e- always in the reactants top right will be strongest reducing agent bottom left will be strongest oxidising agent

when the E* cell is -