ELECTROCHEMISTRY
REDOX PROCESSES (reduction & oxidation reaction, both occur simultaneously)
redox reaction involves transfer of e^-s, 2 half reactions:- 1 oxidation, 1 reduction
oxidation: donate electrons or increase in oxid # (reducing agent/ reductant- oxidised)
reduction: gain electrons or decrease in oxid # (oxidising agent/ oxidant- reduced)
(oxid # : describes rel state of oxidation or reduction)

Rules for oxidation no
-oxid # of an atom / uncombined element is 0
-oxid # of simples ion = charge
-neutral molecules > algebraic sum of oxid # = 0
-ions > algebraic sum of oxid # = charge
-any subs: more electronegative atom has -ve oxid #
-some element fixed oxid # (group I: +1, group II: +2)
-hydrogen: +1 (except in metal hydrides => H: -1)
-oxygen: -2 (except in peroxides; -1, & OF₂; +2)

Electrode potential:-potential difference that exists bet metal & it aq ion
when metal placed in its aq sol of its ions > equilibrium setup > electric layers of metal ions and electrons create pd (electrode potential)
Mⁿ⁺(aq) +ne^- M(s)

electrode potential depends on: nature of metal, conc of aq ions of metal, temp, pressure (if gas)
Standard electrode potential, /V:- pd bet metal and sol containing 1 mol/dm³ of aq metal ions at 25°C (298K) & 1 atm (101325Pa) w/ reference to standard hydrogen electrode
standard electrode potential: conc of aq ion = 1mol/dm³, temp = 25°C (298K), pressure = 1atm (101325 Pa)

3 types of half-cell:
Metal/metal ions half-cell [eg Cu, CuSO₄(aq)]

metal immersed in sol containing its aq ions
metal: takes part in electrode reaction & acts as elec conductor

Redox pairs / Inert electrode half-cell [eg Fe²⁺(aq), Fe³⁺ (aq)]

inert metal (usually Pt) dipped in sol of a pair of aq ions in diff oxid # (redox pair)
inert metal: conducts electrons

Gas / Ions half-cell [eg Cl₂(g), Cl^-(aq)]

inert electrode dipped in sol w/ gas passing through it

Standard redox potential, : pd bet inert Pt electrode & sol w/ 2 aq ions (each of 1 mol/dm³ conc) of same element in diff oxid states at 25°C & 1 atm w/ ref to a standard H electrode


Standard hydrogen potential (used as a ref electode, arbitrary 0.00V)

Pt foil coated w/ finely divided Pt called platinum black immersed in sol of 1 mol/dm³ of H⁺(aq)/H(g)
H(g) at 298K & 1atm bubbled into sol
Pt black: elec cond & catalyst to facilitate equilibrium:
H⁺(aq) + e^- <(Pt black)> ½H₂(g), (reduction) = 0.00V

Measurement of standard electrode potential
Metals/ aq metal ions
Oxidation at metal plate

Zn²⁺(aq)/Zn(s) half-cell
oxidaton: Zn(s) - 2e- Zn²⁺(aq)
reducton: 2H⁺(aq) - 2e- H₂(g)
metal electrode undergoing oxidation (wrt H electrode) > -ve electrode > -ve
reduction (wrt H electrode) > +ve electrode > +ve
standard cell potential / standard emf, _cell: measured by voltmeter
_cell = (reduction) + (oxidation)
wrt H electrode, Zn electrode undergoes oxidation > forms -ve pole >
(Zn²⁺(aq)/Zn(s)) = -ve


salt bridge: contains saturated KCl(aq) as electrolyte
-connects 2 half-cell w/o direct mixing of sol (if mix > electrons moved within sol, not through wire > no current)
-allows flow of ions, K⁺ to +ve (electrons gather at +ve pole > attract K⁺) & Cl- to -ve (electrons leave -ve pole > net +ve charge > attracts Cl^-)

Reduction at metal plate
Cu²⁺(aq)/Cu(s) half-cell


Standard redox potential


IUPAC convention for standard electrode potential
(1) standard H electrode = ref electrode > assigned potential of 0V
(2) in form: oxidant + n e^- reductant
(3) if dirn of half-reaction is reversed > sign of reversed ( => - )
(4) independent of # of moles
(5) +ve : reduction (or forms +ve plate) wrt H electrode [-ve : oxidation (or forms -ve plate) wrt H electrode]
Voltaic cell (electrochemical cell formed by combining 2 half-cells)
-converts chemical energy in redox reaction > elec energy
-Daniel Cell = a voltaic cell; Zn electrode in 1mol/dm³ of ZnSO₄(aq) & Cu electrode in 1mol/dm³ of CuSO₄(aq) connected by salt bridge [contains conc KCl(aq)]

Cell diagram/notation
oxidation half-cell-anode (-ve) on left, reduction half-cell-cathode (+ve) on right
'|'- phase boundary, '||'- salt bridge
X(s)|Xⁿ⁺(aq) || M^N+(aq)|M(s)
oxidation: X(s) > Xⁿ⁺(aq)
reduction: M^N+(aq) > M(s)

if inert electrodes used, least oxidised species next to electrode
oxidation: Pt(s)|X₂(g)|X⁺(aq)
reduction: X₂(aq)|X^-(aq)|Pt(s)

Standard cell potential (standard emf) of voltaic cell,
voltaic cell consists of 2 half-cells (one- oxidation, one- reduction)
standard cell potential of voltaic cell, _cell = (reduction) + (oxidation)
_cell: always +ve

Prediction of feasibility of redox reaction under standard conditions
for spontaneously redox reaction _cell = +ve
_cell only tells if reaction possible, kinetics decide rate of reaction
reaction depends on _cell; large > readily, small > maybe reversible, 0 / -ve > doesn't occur at std conditions (may occur at diff conditions)

Variation of electrode potential w/ concentration of aq ions
Mⁿ⁺(aq) + ne^- M(s)
(reduction) ∝ conc
conc increased > (reduction) more +ve
-due to equilibrium shift to right, according to Le Chateliar's Principle > more electrons removed > more +ve (or less -ve)
(conc decrease > more electrons produced > more -ve/ less +ve)

Batteries (voltaic cells)(2 types)
-(1)primary cells (not rechargeable): Leclanché dry cell, Hg oxide/Zn cell, Zn/MnO₂ dry cell, Ag oxide/Zn cell, Li/I cell
-(2)secondary cells (rechargeable): Nickel-cadmium cell, lead storage battery (lead accumulator)

lead-acid battery
discharge
cathode: PbO₂(s) + 4H⁺(aq) + SO₄^2-(aq) + 2e^- > PbSO₄(s) + 2H₂O(l)
anode Pb(s) + SO₄^2-(aq) - 2e^- > PbSO₄(s)
overall: PbO₂(s) + Pb(s) + 4H⁺(aq) + 2SO₄^2-(aq) > 2PbSO₄(s) + 2H₂O(l)
charge- electrolyse: 2PbSO₄(s) + 2H₂O(l) > PbO₂(s) + Pb(s) + 4H⁺(aq) + 2SO₄^2-(aq)
fuel cells
like normal batteries but reactants & products not stored in cell, supplied into / removed from cell
alkaline hydrogen oxygen fuel cell
anode: 2H₂(g) + 4OH^-(aq) - 4e^- > 4H₂O(l)
cathode: O₂(g) + 2H₂O(l) + 4e^- > 4OH^-(aq)
overall: 2H₂(g) + O₂(g) > 2H₂O(l)
electodes: porous nickel
electrolyte: KOH aq
hydrogen absorbed by KOH(aq) > oxidised
oxygen reacts w/ water in KOH > reduced

ELECTROLYSIS
electrolysis- use of elec energy from voltaic cell to cause redox reaction [when current passed through electrolyte (in aq / molten state) via electrodes]
electrolyte- ionic cpd in aq / molten form (conducts elec current by ion movement, can be decomposed during reaction)
electrodes- conductors that allow current to flow in & out of electrolyte

anode:
-oxidation
-positive (for electrolysis)
-anions move to anode during electrolysis > lose electrons > oxidised
oxidation at anode (+): X^- - e^- > X

cathode:
-reduction
-negative (for electrolysis)
-cations move to cathode during electrolysis > gain electrons > reduced
reduction at cathode (-): M⁺ + e^- > X

Voltaic cell	Electrolytic cell
chem energy from reaction > elec energy	elec energy used to cause redox reaction
reduction at negative cathode	reduction at positive cathode
oxidation at positive anode	oxidation at negative anode

Faraday's law of electrolysis
(1) mass of subs produced directly proportional to quantity of electricity passed
quantity of elec(C) = current(A) × time(s) => Q = It
(2) quantity of electricity required to discharge 1 mole of an ion = simple multiple of a faraday (simple ion; no of faradays required to discharge 1 mole of ion = charge)

1 faraday of elec = electric charges carried by 1 mole of electrons = 96500 C/mol
1F = Le [L: Avogadro's constant= 6.022×10²³, e: electronic charge= 1.601×10^-19C]
Faraday constant, F = (6.022×10²³)(1.601×10^-19C) = 96500 C/mol
M^x+(aq) + xe^- > M(s) [x faradays / moles of electrons needed for oxid]
x = charge supplied for oxid of 1 mole of M^x+ ÷ 96500 coulomb

Selective discharge of ions during electrolysis
if more than 1 cation / anion in solution, reduction / oxid depends on
(1)position of ion in redox series (electrochemical series)
-ion higher in series, more likely to be oxidized (more negative std elecpot > oxid, more positve std elecpot > reduction)

Na₂SO₄(aq): OH^-(aq) & H⁺(aq) from water [H₂O(l) H⁺(aq) + OH^-(aq)] undergo oxid/reduction [OH^- higher in series than SO₄^2- > more likely oxidized; H⁺ lower in series than Na⁺ > more likely reduced] [vol H₂ twice of vol O₂]


(2)conc of ions
-higher conc > ion may be oxidized/reduced eventhough higher in series, due to conc

conc NaCl(aq): Cl^- oxidized due to conc, H⁺ reduced due to lower position in series
[sol > alkaline, NaOH formed, Cl₂ collected less than H₂ due to solubility in water through reaction: Cl₂(g) + H₂O(l) > HOCl(aq) + HCl(aq)]


(3)nature of electrodes used
-lower discharge voltage w/ respect to electrode > easier ion discharged

CuSO₄(aq)
Cu electrodes: anode oxidised > Cu²⁺(aq), Cu²⁺ reduced on cathode
-mass lost from anode gained by cathode, sol conc remains constant > blue colour remains

graphite electrodes: OH^-(aq) oxidized, Cu²⁺ reduced on cathode
-mass gained by cathode, sol con decreases > blue colour fades > colourless > dil H₂SO₄ formed


Applications of electrolysis in industry
-manufacture of chlorine by electrolysis of Brine in diaphragm cell
anode: titanium
cathode: steel
electrolyte: conc NaCl (brine)


-porous asbestos diaphragm separates cathode & anode
-purified saturated brine added to anode compartment (purified to remove Ca²⁺, Mg²⁺ which would formed insoluble hydroxide ppt > can block pores of diaphragm > useless)
-level of sol in anode part kept higher: facilitates ion movement
formation of ions:
    NaCl(aq) > Na⁺(aq) + Cl^-(aq)
    H₂O H⁺(aq) + OH^-(aq)
at anode: Cl^- & OH^- move here, Cl^- more conc > oxidised to chlorine gas: 2Cl^-(aq) - 2e^- > Cl₂(g)
at cathode: H⁺ & Na⁺ collect here, H⁺ reduced to H₂: 2H⁺(aq) + 2e^- > H₂(g)
products of electrolysis: H₂, Cl₂, NaOH (from resulting mixture left)
-extraction of Al from molten Al₂O₃

anode: graphite; 2O^2-(aq) - 2e^- > O₂(g)
cathode: graphite; Al³⁺(aq) + 3e^- > Al(s)
[Al v.dense > sink to bottom]
electrolyte: molten pure bauxite,Al₂O₃.2H₂O (Al ore), w/ cryolite, calcium fluoride + aluminium fluoride (cryolite- sodium hexafluoroaluminate, Na₃AlF₆, CaF₂ + AlF₃ : conductor & reduces mp > 850°C from 2050°C)
problems:
-need large amt of elec (Al readily recycled > recycle cheaper)
-oxide ion oxidised > oxygen > react w/ C > CO₂ > anode worn away
-large amount of waste
-fluoride ion (from cryolite) can be discharged > F cpds > toxic
-electroplating
anode: metal to be plated w/
cathode: obj to be plated
electrolyte: contains metal ions (of metal to be plated)
use: protect from corrosion (metals) & decorative purposes
-electrolytic purification of Cu
impure Cu anode > oxidised to Cu²⁺ which is reduced onto pure Cu cathode
impurities left in sol in aq or solid form [Zn: Zn²⁺(aq), Ag: Ag(s)]
electrolyte: aq H₂SO₄ (2mol/dm³) & aq CuSO₄(0.3 mol/dm³)
aq impurities in electrolyte > purify electrolyte, if not impurities reduced (due to conc)

extracting Cu: from copper ore, copper pyrites, CuFeS₂
2CuFeS₂(s) + 4O₂(g) > Cu₂S(s) + 2FeO(s) + 3SO₂(g)
Cu₂S(s) + 3O₂(g) > 2Cu₂O(s) + 2SO₂(g)
2Cu₂O(s)+Cu₂S(s) > 6Cu(s)+ SO₂(g)
molten Cu > moulds > on cooling SO₂, N₂ & O₂ released > blister Cu
blister Cu: 99% pure but elec conductivity affected & not pure enoguh for alloys
-anodising of Al
anode: Al
cathode: graphite [2H⁺(aq) +2e^- > H₂(g)]
oxygen produced at anode > oxygen reacts w/ Al > thicker oxide layer (can add dyes due to hydrated oxide > coloured Al)
electrolyte: hot sulphuric acid
[oxide layer of Al: non-porous > prevents reaction w/ water & acid, oxide layer of Fe: porous > more water > more oxidation, ie rust]


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