U4 S2 L3 Electrolytic cells Textbook Readings MHR page 776: Electrolytic Cells pages : Electrolysis of Molten Salts pages : Electrolysis of Water pages : Electrolysis of Aqueous Solutions Textbook Practice Items MHR pages : items 9, 10, 11 and 12 page 783: items 15 and 16

Upon completion of this lesson, you should be able to: explain the historical development of electrolysis; for example, Sir Humphrey Davy's use of electrolysis to separate table salt (NaCl) into sodium and chlorine define electrolytic cells as requiring electrical energy to cause non- spontaneous oxidation-reduction reactions to occur draw and label an electrolytic cell including the anode, cathode, salt bridge, power supply, and direction of flow of electrons predict and write balanced half reactions for reactions at the cathode and anode of electrolytic cells write equations for electrolytic cells from half-reactions

Electrolytic cells: –essentially the same as a galvanic cell but;… we add an electric current to reverse the flow of electrons. This energy reverses the chemical reaction. (non- spontaneous – the electrons need to be pushed!) The overall reaction of an electrolytic cell is non- spontaneous therefore we have a negative cell potential (- E o cell ) In electrolytic cells the polarity of the anode and cathode are reversed! –Electrons are now being pushed on to the cathode to cause reduction. Thus, there will be a build up of negative charge (the e- are being gathered up waiting for their chance to jump off and oxidize something – lemings! You may or may not need a porous barrier or salt bridge. As with a Galvanic cell – reduction occurs at the cathode and oxidation occurs at the anode (Leo is an ox) –Polarity: Galvanic cell – anode is negative and cathode is positive Electrolytic cell – anode is positive and cathode is negative.

First electrolytic cell created by Davy. –Electrolysis of molten (800 o C) NaCl (l) to produce molten Na (l) and Cl 2 (g) (p 776)

In an electrolytic cell: –The external electrical force pushes electrons onto one electrode. This electrode is now negative. The positive ions of the electrolyte flow toward this negative electrode, pick up an electron and get reduced (cathode). –The negative ions then move to the positive electrode, lose an electron and become oxidized (anode) As with a Galvanic cell – reduction occurs at the cathode and oxidation occurs at the anode (Leo is an ox) Polarity: Galvanic cell – anode is negative and cathode is positive Electrolytic cell – anode is positive and cathode is negative. Explain why the negative terminal of a battery must be connected to the cathode of an electrolytic cell?

Draw an electrolytic cell of the electrolysis of molten calcium chloride.

Galvanic cell Vs: electrolytic cell p 780 Galvanic CellElectrolytic cell Reaction Purpose Anode Cathode Oxidation Reduction Reduction equation Cell reaction E 0 cell Zn|Zn 2+ ||Cu 2+ |Cu Cu|Cu 2+ ||Zn 2+ |Zn

Electrolysis of water:

Electrolysis of water p. 778 When electrolyzing aqueous solutions water is also present in the electrolyte. –That is, the water may be electrolyzed instead of the dissolved electrolyte. –Half reactions for the electrolysis of water. Mountain range analogy: electrons will take the lowest energy route! When more than one reaction is possible, the redox reaction that gives a negative cell potential closest to zero will be the reaction that occurs. One way to determine this is to write all possible reactions and determine The Ecell for each.

Predict the products of the electrolysis of 1.0 mol/L LiBr (aq). P 781

Short Cut!! List all species as OA or RA The SOA and the SRA will be the two half reactions! Predict the products of the electrolysis of 1.0 mol/L NiCl 2 (aq). RA Ni 2+ Cl - H 2 O (found on both sides) OA

Predict the products of the electrolysis of 1.0 mol/L NaCl (aq) #13 p 83.