1. Lecture 5: Electrochemistry Lecture 5 TopicChapter 20 1. Redox agents & half-equations Reducing & oxidizing agents 20.1 Solving redox by half-equation 20.2 Steps! 2. Voltaic cells are redox reactions 20.3 Separate “half-cells” 3. Batteries Electromotive force 20.4 Batteries & Calculating Ecell 20.7 Fuel Cells 4. Corrosion & Electrolysis Corrosion 20.8 Electrolysis 20.9

2. Corrosion Corrosion is unwanted spontaneous redox. Sacrificial anodes protect metals. Electrolysis reverses corrosion with input energy.

3. p. 858 - 9 Is an undesirable spontaneous redox reaction that corrodes, or eats away at, susceptible metals. Iron rusts easily in the presence of air and water. Acid, salts & contact with less easily oxidized metals increase that rate of rusting. Anode: Fe(s) -> Fe +2 (aq) + 2e-Ered = -0.44 V Cathode: O2(g) + 4H + (aq) + 4e- -> 2H2O(l) E red = +1.23 V 4Fe +2 (aq) + O2(g) + 4H2O + 2 xH2O(l) -> 2Fe2O3-xH2O(s) + 8H + anode cathode Fe +2 e- Fe is oxidized by O2 (notice that the Ered of Fe is lower than that of O2). Fe close to O2 exposure acts as cathode & is the site of O2 reduction Fe not exposed continues to supply e- as needed & therefore is anode. Iron is first oxidized from Fe to Fe +2, and then to Fe +3. Fe +3 combines with oxygen and water to form rust, Fe2O3-xH2O = RUST. Reduction of O2 requires H, often supplied by acid H. Why does rust form in a particular spot? O2 Rust forms where the metal is exposed to oxygen & H2O, and often acids. IRON WATER RUST Corrosion

4. p. 858 - 9 Why do iron nails left in a ditch quickly rust and begin to “disappear” while aluminum cans don’t? Look at the activity series in Chapter 4! Exposure of Al to O2 and water DOES oxidize the metal, forming a dull & hazy layer of hydrated aluminum oxide on the exposed surface. BUT this layer of oxide is IMPERMEABLE to water and oxygen and prevents, or LIMITS, further corrosion of the underlying aluminum. Aluminum is much higher on the series than iron, indicating that it oxidizes more easily, yet it doesn’t crumble when exposed to O2 and water, it just gets a hazy, dull surface. Al Al2O3-xH2O O2 & H2O Metals with ‘self-limiting’ corrosion Mg, stainless steel & silicon oxide (semiconductor chips) also undergo self-limiting oxidation.

5. p. 858 - 9 There are several ways to prevent corrosion; effectiveness varies 1. Cover the surface of metals with paint or another sealer Fe +2 + 2e- -> Fe Ered = -0.44 V Zn +2 + 2e- -> Zn E red = -0.76 V Ered values & activity series show that Zn is easier to oxidize then Fe, so Fe is protected until all Zn is “used up”. Fe cathode So what’s the iron doing here? How does galvanizing protect the iron? O2 Prevents O2 and water from contacting the metal & reacting with it 2. “Galvalnize” the metal e- Zn anode Zn +2 Zn -> Zn +2 + 2e- O2 + 4H + + 4e- -> 2H2O The iron is accepting e- from the oxidized Zn and passing them to O2 that is reduced. So the iron isn’t changed, but serves as the cathode in the sense that it is the surface upon which the O2 is reduced. Cathodic protection Sacrificial anode process of protecting a metal from oxidation by making it the cathode of an electrochemical reaction. the easily oxidized metal that is “sacrificed” to corrosion & protects the metal that is the cathode. See Fig 20.27, p.883 Preventing corrosion

6. There are several ways to prevent corrosion; effectiveness varies. CET application of sacrificial anode

7. p. 860 …that is used to precipitate or deposit atomic metals. Electricity is used to drive these non-spontaneous reactions forward. For example, electrolysis can be used to “decompose” NaCl: NaCl  2Na + Cl2(g) Electrolysis can be used to: 1. 2. Recover metals from aqueous solutions This is why you never want to sink “normal” batteries in salt water. The Cl2 gas is toxic &the salt water is a great electrolyte and will conduct the electrons right too you. Submarine movies of WWII vintage…. “Plate” cheap metals with nickel, silver, gold etc. V NiSO4 Ni anode Steel cathode e- Ni +2 How does electroplating work? The battery sends voltage (e-) to the steel cathode. The electrons can be used to reduce either water or the aqueous Ni +2 ions. 2H2O + 2e- -> H2 + 2OH - Ered = -0.83 V Ni -> Ni +2 + 2e-Ered = -0.28 V …so Ni is easier to reduce The battery “pulls” e- from the Ni anode, oxidizing it and creating Ni ions that are used in plating the steel. Overall reaction? Nickel is both oxidized and reduced so Ecell = 0V ! Electrolysis is ‘reversed’ corrosion