U2- Electrochemistry Flashcards
displacement
an element displaces another element from a compound
addition
one molecule attaches to another across a double bond
decomposition
a substance breaks down into one or more smaller particles
condensation
produces a small molecule byproduct, usually water
oxidation
loss of electrons
increases in oxidation number
gain O2/ loss of H
reduction
gain of electrons
decreases in oxidation number
loss of O2/ gain of H
what type of reaction is a redox reaction
displacement
oxidant/ oxidising agent
reactant that cause another substance to be oxidised (what is reduced)
reductant/ reducing agent
reactant that causes another substance to be reduced (what is oxidised)
oxidation number
measure of electron density around an atom, compared to its elemental form (sign b4 no.)
Ox rule: 1
sum of oxidation numbers = total charge
Ox rule 2 and 3
2: F= -1
3: G1= +1, G2= +2
Ox rule 4
H bonded to non metals = +1
H bonded to metal = -1
Ox rule 5
Oxygen = -2 unless in peroxides the O= -1
Ox rule 6
If 2 or more different G7 elements present, more electronegative one is -1
KOHES
Key element
balance oxygens
balance hydrogens
balance charges by adding electrons
STATES
(separate into 2 half equations)
why do oxidation and reduction occur simultaneously during any redox
bc any loss of e by one substance must be accompanied by a gain in e by something else
conjugate oxidant/reductant
substance produced when a reductant loses/ oxidant gains e, containing the element that has increased/decreased in oxidation number
standard redox conditions
all solutions = 1M
all gases @ 1atm
all solids = pure
electrode= made of reductant if it is solid and conductive otherwise inert electrode eg graphite or platinum
salt bridge
structure containing free-moving unreactive ions that provides a connection bw two half cells. (often filter paper soaked w KNO3)
electrochemical series
higher on left= oxidants at cathode that are preferentially reduced
lower on right= reductants at anode that are preferentially oxidised
predicting spontaneous reaction: electrochemical series
if oxidant is higher in the series than reductant, reaction will occur
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anode
site of oxidation
galvanic cell= negative (bc a loser)
electrolytic cell = positive (bc PANIC)
cathode
site of reduction
galvanic cell= positive
electrolytic cell = negative (bc PANIC)
galvanic cells vs direct contact
direct contact: nag released in the form of heat
galvanic cell: nag released in the form of electrical energy
cell potential
difference bw standard electrode potential value of the cathode and that of the anode
= E (cat or oxidant) - E (anode or reductant)
(NEVER MULTIPLY BY MOLES)
electrical potentual
pressure with which electrons are forced around the circuit
standard cell notation
anode I conj oxidant II oxidant I cathode
or (if inert electrodes)
anode I reductant, conj oxidant II oxidant, conj reductant I cathode
galvanic vs electrolytic
G. E
produce e. Consume e
ions spontaneously react ions don’t spontaneously react
convert chem nrg -> e nrg convert e nrg -> chem nrg
anode = neg; cathode = pos anode = pos; cathode = neg
conjugate redox pair
two substances (one oxidant and one reductant) that share an atom (or atoms) that have gained or lost electrons
applications of electrolysis
electrolysis of water
production of reactive metals
production of the substances
electroplating
electrorefining
electrolysis of water
diaphragm/ separator may be required bc products are able to spontaneously react
extraction of AL
treat bauxite (Al2O3) with NaOH to remove impurities
Dissolve Al2O3 in molten cryolite (Na3AlF6)- cryolite has a lower mp vs alumina which reduces operating costs
electrolyse using graphite electrodes
side reaction: C(s) +O2 (g) –> CO2 (g)
why are graphite electrodes used and why do they have to be replaced regularly
inert + cheap + high MP
react w oxygen produced in side reaction to form CO2 gas
why does Al sink to the bottom during extraction
denser than cryolite
corrosion + rusting
corrosion= disintegration of a metal as a result of a redox chemical reaction
rusting= specific eg involving iron
role of water droplet in rusting
enables ion transfer between the anodic site and cathodic site
prevention of the corrosion of iron
barrier: remove O2 or water from system eg paint
chemical: coat iron with more reactive metal that preferentially oxidises vs iron
commercial batteries
galvanic, storage cells the favour spontaneous redox reactions producing e current.
reactant chemicals are not replenished and eventually run out
electroplating
metal substance to be electroplated @ cathode
electrorefining
purifying metals
eg impure copper anode + pure copper cathode
anodic sludge formed
