c23 - redox and electrode potentials

when are manganate redox titrations used

to estimate amount of metal ions in a solution or % of metal in a solid

how are redox titrations different to neutralisation titrations

redox titrations involve transfer of electrons rather than neutralisation, titrations with potassium manganate are self indicating so don’t requitre an indicator unlike neutralisation

problems of using potassium manganate solution

kmno4 oxidises cl- and br- so solution has to be acidified with h2so4 not hcl, difficult to read at meniscus because too dark, not 1ry standard (too strongly coloured to see when all crystals dissolved, over time oxidises o2 to form manganese oxide so conc not guaranteed)

how to make fe in +2 oxidation state (required state for analysis)

if element fe(0) then react with sulfuric acid to oxidise to fe2+, if fe3+ then react with zn to reduce to fe2+ and then remove excess zn

equation for manganate redox titration with fe2+

MnO4-(aq purple) + 8H+(aq) + 5Fe2+(aq) → Mn2+(aq almost colourless) + 5Fe3+(aq) + 4H2O(l)

equation for manganate redox titration with ethanedioic acid (COOH)2

5C2O42- + 2MnO4- + 16H+→ 10CO2 + 2Mn2+ + 8H2O (requires warming at start of reacrtion as -ve charged ions repel)

how does the Mn2+ produced in manganate redox titration with ethanedioic acid act as a catalyst

autocatalysis - 1: 4Mn2+ + MnO4- + 8H+ → 5Mn3+ + 4H2O then 2: 2Mn3+ + C2O42- → 2CO2 + 2Mn2+

overall redox equation for iodine thiosulfate redox titration

2S2O32-(aq) + I2(aq brown) → 2I-(aq yellow then black w starch) + S4O62-(aq)

indicator required for ioding thiosulfate reactions

starch, only add when end point approaching and iodine is pale yellow straw colour, at end point solution should go from blue black to colourless

use of iodine thiosulfate titrations to determine cu content in copper alloys

1-convert Cu to Cu2+ with nitric acid, 2-Cu2+ reacts with excess I- to form solution of I2 2Cu2(aq) + 4I-(aq) → 2CuI(s) + I2(aq)

use of iodine thiosulfate titrations to analyse ClO- content in household bleach

1- ClO-(aq) + 2I-(aq) + 2H+(aq) → Cl-(aq) + I2(aq) + H2O(l) 2- I2 reacts with S2O32- ions in titration

metal/metal ion half cell

metal rod dipped in solution of its own aqueous metal ions, metal rod enables electrons to move to or from half cell, metal in 2 different oxidation states, at equilibrium forward reaction always shown as reduction

ion/ion half cell

aqueous ions of same element in 2 different oxidation states, no metal to transport electrons to or from half cell so inert platinum electrode used instead, at equilibrium forward reaction always shown as reduction

standard electrode potential

the e.m.f. of a half cell measured against the standard hydrogen electrode undr standard conditions of 298K, 100kPa, solution conc of 1.0moldm-3

what do electrode potentials tell us

+ve e- indicates substance more likely to be reduced and gain electrons than h+, -ve e- indicates substance less likely to be reduced and gain electrons than h+,

cell potential/ e cell

most +ve half cell - least +ve half cell

what happens at each electrode in half cells

electrons flow from least +ve half cell (anode) to most +ve half cell (cathode), reduction takes place at cathode (+ve) and oxidation takes place at anode (-ve)

role of salt bridge in half cells

to allow ions to flow

predictions from electrode potentials

reduction reaction with highest standard electrode potential will occur, redox system with more +ve electrode potential will react from left to right and gain electrons, redox system with less +ve electrode potential will react from right to left and lose electrons,

non-standard conditions - limitation of using predictions from standard electrode potentials

e cell values refer to standard conditions, if cell operates under non-standard conditions then electrode potential is not same as standard, as cell operates, conc of solutions change so e cell values change

just because it can happen, does not mean it will - limitation of using predictions from standard electrode potentials

e- values greater than or equal to 0 mean the reaction is feasible, if roro very slow it may appear as if nothing is happening, if ea very high then it may prevent feasible reaction from occurring

primary storage cells

used for low current, long storage devices e.g. clocks, electrical energy generated by irreversible redox reactions at electrodes, non-rechargeable, chemicals eventually used up and battery must be discarded/recycled, improper disposal risks fire + environmental contamination

secondary storage cells

rechargeable so can be reused as cell reaction that generates electrical energy is reversible, regenerating the chemicals e.g. lead-acid batteries in cars, li+ batteries in phones

fuel cells

generate voltage from reaction of fuel with o2, electrolyte remains in cell whereas reactants and products flow in and out, can operate continuously provided fuel and o2 are supplied, h2 fuel cells do not produce co2 only h2o but harder to store h2

acid h2 fuel cell

oxidation (anode): h2(g) → 2h+(aq) + 2e-, reduction (cathode): 0.5o2(g) + 2h+(aq) + 2e- →h20(l), overall: h2(g)+0.5o2(g) →h2o(l)

alkali h2 fuel cell

oxidation (anode): h2(g) + 2oh-(aq) →2h2o(l) + 2e-, reduction (cathode): 0.5o2(g) + h2o(l) + 2e- → 2oh-(aq), overall: h2(g)+0.5o2(g) →h2o(l)