ELECTROLYSIS

Compare and contrast voltaic (galvanic) and electrolytic cells Explain the operation of an electrolytic cell at the visual, particulate and symbolic levels Include: molten and aqueous electrolytic cells Additional KEY Terms

Process of using electricity to force a nonspontaneous reaction Electrolysis: “electro” – electricity “lysis” – break Electrolytic cells Different: it is nonspontaneous so the opposite metals are going to be reduced and oxidized from a voltaic cell A direct current source (battery) is attached to inert electrodes – C or Pt Different: electrodes do not take part in reaction – “plate out” or “lose mass”

Inert electrodes must be charged to attract the appropriate ions to the appropriate side Anions migrate to (+) anode and are oxidized Cations migrate toward (-) cathode to be reduced Different: electrodes oppositely charged Same: e - still flow from anode to cathode Same: anode oxidized, cathode reduced

most metals occur naturally as compounds NaCl, Cu 2 O, TiO 2, CuF 2 Need very high temperatures to make salts molten - “melt” Unwanted elements of a compound are called “impurities” forcing the non-spontaneous redox is a common process for obtaining pure metals 1. Molten Electrolytic Cells

Na + (l) + Cl – (l) → Na (s) + Cl 2(g) E° R = V E° O = V 2 Cl - (l) → Cl 2(g) + 2e - 2 Na + (l) + 2e - → 2 Na (s) E o cell = V Notice: weaker element – Na is reduced Notice: stronger element – Cl is oxidized Down's Cell: A battery must supply 4.07 V of power to make this electrolytic cell work

Electroplating: object to be plated – cathode desired metal – anode immersed in a solution of same metal ions current used to plate layer of metal onto another surface by reducing its ions Only one element – the metal is oxidized and then reduces itself in second location

Corrosion: reaction of metals with oxygen - spontaneous metal – oxidized oxygen – reduced Most metals have reduction potentials below oxygen – except gold Sacrificial anode – second “weaker” element that sacrifices its electrons and oxidizes to save metal

Electrolysis of water: Water can be both reduced and oxidized Overall: 2 H 2 O (l) → 2 H 2(g) + O 2(g) Oxidation: Reduction: Net: 2 H 2 O (l) → 4 H + (aq) + O 2(g) + 4e – 4 H 2 O (l) + 4e – → 4 OH – (aq) + 2 H 2(g) 6 H 2 O (l) → 2 H 2(g) + O 2(g) + 4 H + (aq) + 4 OH – (aq) These ions recombine to form 4 H 2 O 2

2. Aqueous Electrolytic Cells: 3 possible reactants that will compete for electrons - cations, anions and water must predict which substances will be most likely oxidized and which reduced Because water is present – it can oxidize and reduce – you might not get what you expected in the reaction This is done by comparing the oxidation and reduction potentials of all possible reactants

What are the products formed at each electrode during the electrolysis of aqueous KI? possible reactants: K + I – H 2 O 2 I – (aq) → I 2(s) + 2e – E° o = – 0.54V H 2 O (l) → 2 H + (aq) + ½ O 2(g) + 2e – E° o = –1.23V I - has most positive oxidation potential o solid iodine is formed at the anode K + cannot lose more electrons - only I – and H 2 O can be oxidized Oxidization:

K + (aq) + 1e – → K (s) E° R = – 2.93V 2 H 2 O (l) + 2e – → 2 OH - (aq) + H 2(g) E° R = – 0.83V possible reactants: K + I – H 2 O I - cannot gain more electrons - only K + and H 2 O could be reduced Reduction: H 2 O has most positive reduction potential o hydrogen gas is formed at the cathode Overall: 2 H 2 O (l) + 2 I – (aq) → H 2(g) + I 2(s) Spectator ions created in the reaction would be K + and OH -

CAN YOU / HAVE YOU? Compare and contrast voltaic (galvanic) and electrolytic cells Explain the operation of an electrolytic cell at the visual, particulate and symbolic levels Include: molten and aqueous electrolytic cells Additional KEY Terms