 17.1 Explain how a non-spontaneous redox reaction can be driven forward during electrolysis  17.1 Relate the movement of charge through an electrolytic cell to the chemical reactions that occur  17.1 Apply the principle of electrolysis to its applications such as chemical synthesis, refining, plating, and cleaning.

 17.2 Relate the construction of a galvanic cell to how it functions to produce a voltage and an electrical current  17.2 Trace the movement of electrons in a galvanic cell  17.2 Relate chemistry in a redox reaction to separate reactions occurring at electrodes in a galvanic cell

 Oxidation reduction reactions involve a transfer of electrons.  OIL- RIG  Oxidation Involves Gain  Reduction Involves Loss  LEO-GER  Lose Electrons Oxidation  Gain Electrons Reduction

 Galvanic Cells (Voltaic) = A battery which uses spontaneous chemical processes to produce electricity ◦ The amount of electricity depends on how bad the atoms (molecules) want the electrons or want to give them up.

 Corrosion: The oxidation of metals over time from being oxidized by surrounding oxidizing agents (such as oxygen). ◦ Generally very slow, but some are more quickly oxidized (depending on activity of metal as a solid)

 How to stop corrosion: Sacrificial Anodes. Since some metals corrode easier than others, we have the metal we want safe (steel, iron) in contact with a metal that is more easily oxidized (like zinc). ◦ The zinc gets oxidized first, and loses electrons and takes the hit instead of the iron or steel.

 Running a reaction backwards ◦ Forcing electrons onto the atoms/molecules  Separating Atoms ◦ Give everyone an octet without each other  Used to separate metals from their salts

 Using electrolysis to place aqueous metals onto a surface. ◦ This is how jewelry is plated in gold and silver, how silver ware is coated in silver, but not completely out of silver ware.

 Moving electrons is electric current.  8H + +MnO Fe +2 +5e -  Mn Fe +3 +4H 2 O  Helps to break the reactions into half reactions.  8H + +MnO e -  Mn +2 +4H 2 O  5(Fe +2  Fe +3 + e - )  In the same mixture it happens without doing useful work, but if separate

H + MnO 4 - Fe +2  Connected this way the reaction starts  Stops immediately because charge builds up.

Reducin g Agent Oxidizin g Agent e-e- e-e- e-e- e-e- e-e- e-e- AnodeCathode

 Oxidizing agent pushes the electron.  Reducing agent pulls the electron.  Unit is the volt(V)

Zn +2 SO M HCl Anode M ZnSO 4 H + Cl - H 2 in Cathod e

1 M HCl H + Cl - H 2 in  This is the reference all other oxidations are compared to  E º = 0  º indicates standard states of 25ºC, 1 atm, 1 M solutions.

 Zn(s) + Cu +2 (aq)  Zn +2 (aq) + Cu(s)  The total cell potential is the sum of the potential at each electrode.  E º cell = E º Zn  Zn +2 + E º Cu +2  Cu  We can look up reduction potentials in a table.  One of the reactions must be reversed, so change the sign.

 Determine the cell potential for a galvanic cell based on the redox reaction.  Cu(s) + Fe +3 (aq)  Cu +2 (aq) + Fe +2 (aq)  Fe +3 (aq) + e -  Fe +2 (aq) E º = 0.77 V  Cu +2 (aq)+2e -  Cu(s) E º = 0.34 V  Cu(s)  Cu +2 (aq)+2e - E º = V  2Fe +3 (aq) + 2e -  2Fe +2 (aq) E º = 0.77 V

 solid  Aqueous  Aqueous  solid  Anode on the left  Cathode on the right  Single line different phases.  Double line porous disk or salt bridge.  For the last reaction  Cu(s)  Cu +2 (aq)  Fe +2 (aq),Fe +3 (aq)

 Rusting - spontaneous oxidation.  Most structural metals have reduction potentials that are less positive than O 2. ◦ If you are more positive on the chart, you can oxidize anything below you (the product)

Water Rust Iron Dissolves- Fe  Fe +2 e-e- Salt speeds up process by increasing conductivity

 Coating to keep out air and water.  Galvanizing - Putting on a zinc coat  Has a lower reduction potential, so it is more. easily oxidized.  Alloying with metals that form oxide coats.  Cathodic Protection - Attaching large pieces of an active metal like magnesium that get oxidized instead.

1.0 M Zn +2 e-e- e-e- Anode Cathode 1.10 Zn Cu 1.0 M Cu +2

Cathode (Reduction) Half-Reaction Standard Potential E ° (volts) Li + (aq) + e -  Li(s) K + (aq) + e -  K(s) Fe 2+ (aq) + 2e -  Fe(s) Pb 2+ (aq) + 2e -  Pb(s) Cu+(aq) + e-  Cu(s) 0.52 I 2 (s) + 2e-  2I-(aq) 0.54 Ag + (aq) + 1e-  Ag (s) 0.80 Pt +2 (aq) + 2e-  Pt (s) 1.23 Cl 2 (g) + 2e-  2Cl- (aq) Write the chemical shorthand for a Lead and Lithium battery. 2.What is the cell potential for a Iron and Platinum battery? 3.Which element (and charge) is the best oxidizing agent? 4.Which element (and charge) is the best reducing agent?