electrolysis

• what are electrolytic cells?

  opposite to galvanic cells

  • electrical energy → chemical energy

  • involves passage of electrical energy from direct current DC power supply (battery) thru conducting liquid

  • non spontaneous redox reaction (electricity forces to occur)
• molten


  ionic compound heated above m.pt → liquid + no water
• 4 key features of electrolytic cells

  • inert electrodes typically used to conduct electricity

  • molten ionic compound (NaCl) → electrolyte (conducting liquid)

  • No water present - only Na+ and Cl

  • Electrical energy comes from power supply (acts as ‘electron pump’) → takes e- from anode + pushes onto cathode
• example of molten electrolysis

  electrolysis of molten sodium chloride

  • neg electrode connect to neg terminal

  • pos electrode connect to pos terminal
• substances that are preferentially reduced and oxidised

  NEGATIVE electrode = CATHODE (reduction)

  • Weak oxidant preferentially reduced → LHS bottom (higher E value = stronger reduction)

  POSITIVE electrode = ANODE

  • Weak reductant preferentially oxidised → RHS top (lower E value = stronger oxidation)
• strongest oxidant and reductant

  RHS down = strongest oxidant (R)

  LHS up = strongest reductant (O)

  
• overall reaction


  • Non-spontaneous reaction → weak oxidants don’t react w weak reductants

  • Reverse reaction is spontaneous

  • Occur in one container → products need to be separated or else reform reactants

  • Anions to anode, cations to cathode

  • OIL RIG CAT true but polarities diff

  • RHOL is OPPOSITE - reduction lower, oxidation higher
• competition at electrodes


  • several chemicals present at electrodes that might be able to react

  • water potential reactant in aqueous solutions

  • electrode may participate in reaction (anode)

  USE ECS TO PREDICT
• 3 principles of electrolysis

  • Oxidation at anode; reduction at cathode

  • Strongest oxidant (LHS top): reduction at cathode

  • Strongest reductant (RHS bottom): oxidation at anode
• diagram of electrolytic cells

  • e- flow
  • power supply charge
  • cathode + anode
  • electrode material
  • charges of electrodes (pos or neg)
  • ion flow (cation to cathode, anion to anode)
  • states

  
• aqueous


  contains water
• example of aqueous electrolysis

  • have to consider water ( +1.23, -0.83 ) → water reacts at both electrodes

  
• differences b/w molten, aqueous and concentrated NaCl


  • concentrated NaCL = brine = > 1.0M → Cl- reacts at anode instead of water

  
• comparing electrolytic and galvanic cells

  similarities:

  • oxidation at anode + reduction at cathode
  • cations to cathode, anions to anode
  • contain electrolyte
  • electrodes (cathode + anode)

  differences:

  • galvanic cathode is positive + anode is negative ; electrolytic cathode is negative + anode is positive
  • galvanic is spontaneous ; electrolytic is non-spontaneous
  • galvanic is chem to electrical energy ; electrolytic is electrical to chem energy
  • galvanic produces energy ; electrolytic consumes energy
• real life example of galvanic + electrolytic cells


  galvanic example = mobile phone battery (discharging)

  electrolytic example = electric cars, charging mobile phone, electroplating
• what is a commercial example of a molten electrolyte?

  DOWNS CELL → produce sodium + chloride

  cathode: Na+ (l) + e- → Na (s)

  anode: 2Cl- (l) → Cl2 (g) + 2e-

  electrodes are inert:

  • anode: conductive material (graphite, platinum)

  • cathode: iron (effectively inert bc e- pumped onto electrode, preventing iron from reacting → cannot form a neg iron bc its a metal)

  
• electrodes in an electrolytic cell (excludes electroplating)

  electrodes are inert:

  • anode: conductive material (graphite, platinum →inert )

  • cathode: iron (effectively inert bc e- pumped onto electrode, preventing iron from reacting → cannot form a neg iron bc its a metal)
• electrodes in electroplating cell

  • anode (+) : metal used for plating

  • cathode (-) : object being plated
• advantages + disadvantages of molten electrolytes

  A) more energy

  B) maintain

  C) no water present

  D) molten state

  E) dangerous

  F) burns

  G) high

  H) increase

  I) wasted energy

  
• electrolyte in a Downs cell

  electrolyte contains NaCl and CaCl2

  • CaCl2 can’t participate in reaction + lowers mpt of electrolyte (800 °C to 600 °C) → saves on energy costs
• purpose of the iron mesh screen in the Downs cell

  iron mesh screen:

  • separates products from each other

  • products of electrolysis are strong oxidants/ reductants + will readily react to reform reactants → construction of Downs cell minimises contact
• commercial example of an aqueous electrolyte

  membrane cells used for aqueous electrolytes → 2 compartments separated by semipermeable membrane

  • used to produce: NaOH, Cl2 and H2 from concentrated NaCl solution (aka brine)

  cathode (-): 2H2O (l) + 2e- → H2 (g) + 2OH- (aq)

  
  

  anode (+): 2Cl- (aq) → Cl2 (g) + 2e-

  • water is stronger reductant than Cl-, but using conc. NaCl, Cl- is oxidised instead

  
  

  inert electrodes typically used

  
  

  • if cathode made from metal it’ll be inert (can’t gain e- or e neg charged)

  
  

  • production of Al uses inert + reactive electrodes

  
• purpose of semipermeable membrane

  semi-permeable membrane helps prevent contact b/w reactive products

  • made from polymer: only allows cations to pass

  • prevents mixing of products formed at electrodes

  • only Na+ moves from one chamber to other → neg ions cant pass thru

  • results in pure NaOH w little Cl- contamination
• advantages of using the membrane cell

  adv:

  • NaOH produced not contaminated w NaCl

  • no need to heat electrolyte + cost of production is reduced bc med temps used (80 - 90 °C)
• how is aluminium produced through electrolysis?

  made from aluminium oxide (Al2O3 -alumina) by electrolysis

  • alumina melts at v high temps (2050 celsius)

  • dissolved in molten cryolite → performed at much lower temp (950-1000 celsius) → avoids v high energy costs

  • Hall-Heroult cell used

  
  

  cathode (-): Al3+ (in cryolite) + 3e- → Al (l)

  • molten aluminihv sinks to bottom + is siphoned

  
  

  anode (+): 2O2- (in cryolite) → O2 (g) + 4e-

  • O2- reacts w carbon anode to make CO2

  
  

  • carbon anodes hv taken part in this reaction → replace regularly ($)

  
• what are the common design features and general principles of commercial cells?

  • separation of products: reactive products = spontaneous reaction to reform
  • inert vs reactive electrodes → depends on electrode cost (Pt expensive to replace) + mpt of electrode (if molten need high mpt)
  • molten ($) vs aqueous electrolyte → presence of water + reaction
  • chemical additives → reduce mpt of molten electrolyte or solvent for compound being electrolysed

  
• electroplating


  thin surface coating of metal placed over another surface

  • important commercial application of electrolysis
• electroplating cells


  • electroplating performed in electrolytic cells

  • object to be plated connected by wire to neg terminal of power supply → becomes neg electrode (cathode) of cell

  • object immersed in electrode

  
• what are the key features in an electroplating cell?

  • electrolyte solution must contain ions to be plated

  • cations (Sn2+) move twrds cathode (object to be plated)

  • anions (NO3-) move twrd anode

  • movement of electrolyte ions allows current to pass thru cell → forms internal circuit

  • power supply: forces e- into cathode + removes from anode

  
• cathode + anode reactions during electroplating

  cathode (-): cations attracted to cathode

  • coating of metal forms on object as ions reduced

  anode (+): anions attracted to anode

  • electrode of selected metal acts as anode

  • power supply withdraws electrons, forcing oxidation, metal electrode slowly dissolves + cations formed

  overall [electrolyte] remains constant→ same reversible reaction at both electrodes
• 3 key factors that determine amount of products that will form in an electrolytic cell

  • charge of ion involved in electrode reaction

  • current flowing thru cell

  • length of time that current flows
• Faraday’s First Law of Electrolysis


  ‘mass of metal produced at cathode directly proportional to electrical charge passed through cell’ (m ∝ Q)

  Q = I x t

  • Q - coulomb (C)

  • I - amps (A)

  • t - seconds (s)
• Faraday’s Second Law of Electrolysis

  Q = n(e-) x F

  • F - 96500C

  • n(e-) - moles of electrons

  • Q - coulombs (C)
• copper electrorefining


  impure copper purified by electrolysis

  • sheets of impure copper placed in a large tank of sulfuric acid w sheets of positioned b/w them

  • impure copper acts as anode; pure copper acts as cathode

  • copper in the impure copper oxidised into copper ion + migrate to cathode where they’re reduced

  • impurities less reactive than copper (silver and gold) fall from anode + collect at bottom of tank

  • pure copper reduced onto cathode
• purpose of the power supply

  • acts as 'electron pump'
  • withdraws electrons from anode + forces into cathode
  • allows for non-spontaneous redox reaction to occur
  • allows for relatively weak oxidant to react with a relatively weak reductant
• purpose of the variable resistor

  to maintain a constant current by restricting the flow of current in a cell
• why might mass of metal plated be too high or too low?

  too high:

  • soluble ions remain on the cathode = increase actual mass →dip electrode in distilled water to remove any soluble ions
  • water remains on cathode after dipping = increase actual mass → place in oven at 110CELSIUS until constant mass achieved

  too low:

  • patting dry = low mass → place in oven at 110CELSIUS until constant mass achieved
  • metal formed may not have attached to electrode → use scourer to clean electrodes OR use filter paper to filter the electrolyte + account for the metal formed