1. Lesson 2

2. Galvanic Cells In the reaction between Zn and CuSO 4, the zinc is oxidized by copper (II) ions. Zn 0 (s) + Cu 2+ (aq) + SO 4 2-  Cu 0 (s) + Zn 2+ + SO 4 2- (aq) Zn becomes Zn 2+, a loss of 2e - (oxidation) Cu 2+ becomes Cu, a gain of 2e - (reduction)

3. Separating the zinc metal from the solution containing copper ions, and placing a metal conductor between the metals creates electricity. Electrons that are lost by zinc are forced to travel through the metal conductor to reach the copper ions. The movement of electrons is known as electric current.

4. Galvanic cell This apparatus is called a galvanic cell. A device that converts chemical energy from redox reactions into electrical energy.

5. Galvanic cell A Galvanic cell is a spontaneous reaction A reaction that proceeds on its own without outside assistance (energy). The oxidation of zinc and reduction of copper occur in separate beakers called half cells. One of the two compartments in a galvanic cell Composed of an electrode and a electrolytic solution.

6. Galvanic cell The metal in each beaker is called a electrode. A solid electrical conductor where the electron transfer occurs.

7. Galvanic cell Each electrode has a special name Anode – The electrode where oxidation occurs (think of anion, a negative ion) Cathode – The electrode where reduction occurs (think cation, a positive ion) REDCAT (Reduction Cathode) Metals and non-reactive conductors such as graphite are often used as conductor electrodes.

8. Galvanic cell Each electrode is immersed in an electrolytic solution that contains the same metal ions as the electrode. Zinc is in a zinc nitrate solution Copper is in a copper (II) nitrate solution

9. Galvanic cell Half cells are connected by a salt bridge. Concentrated solution of electrolyte, it should not react with the other chemical in the reaction.

10. How does a salt bridge work? How does a salt bridge work? (do not look at text, they used the wrong metals) The purpose of the salt bridge is to provide ions to prevent charge from building up. In a way it is much like simple diffusion.

11. How does a salt bridge work? Every time a zinc atom is oxidized to an ion it would make the solution more positive, which would then stop the reaction. The nitrate from the salt bridge moves in and balances it. On the other side, every time a copper ion is reduced it would make the solution more negative and stop the reaction. A sodium on then moves in a negative nitrate ion leaves. This allows the circuit to continue without a build up of charge.

12. Questions Page 400 # 1-7

14. Cell Reactions The chemical equation for this reaction can be broken down into 2 parts, called half cell reactions. Anode half –reaction = Zn (s)  Zn 2+ (aq) + 2e - (oxidation) Cathode half –reaction = Cu 2+ (aq) + 2e -  Cu (s) (reduction)

15. Cell Reactions Therefore, as the cell operates the mass of the zinc electrode decreases and the mass of the copper electrode increases. Anode half –reaction Zn (s)  Zn 2+ (aq) + 2e - Cathode half –reaction Cu 2+ (aq) + 2e -  Cu (s) Overall cell reaction = Cu 2+ (aq) + Zn (s)  Cu (s) + Zn 2+ (aq)

16. Writing Cell Reactions Step 1) Establish elements oxidized and reduced Step 2) Balance charges You may need to multiply the half reactions by a common multiple to balance out the number of electrons.

17. Example: A strip of silver in a solution of silver nitrate, and a copper strip in a solution of copper (II) nitrate Step 1) Copper is higher on the activity series so it will be oxidized. Anode half-reaction Cu (s)  Cu 2+ (aq) + 2e - Cathode half-reaction Ag + (aq) + e -  Ag (s) -note: the number of e - lost by copper is not equal to the number of e - gained by silver.

18. Step 2) Multiply both sides of the cathode reaction by 2 so that the number of e - is equal to the e - gained. Anode half-reaction Cu (s)  Cu 2+ (aq) + 2e - Cathode half-reaction 2 Ag + (aq) + 2e -  2Ag (s) _ Overall cell reaction Cu (s) + 2 Ag + (aq)  Cu 2+ (aq) + 2Ag (s)

19. Practice – page 397 Activity 5.8 – Designer Cells

20. Corrosion

21. Corrosion – The deterioration of metals as a result of oxidation.

22. With the exception of a few un-reactive metals, most metals are found as minerals and not in elemental form. Without protection against corrosion, most metals are oxidized by their environment.

23. A few metals such as copper, zinc and aluminum form protective coatings when they oxidize. This makes them more corrosion resistant than other metals that are lower on the activity series. This is why they are commonly used to coat and protect other metals.

24. Example Aluminum reacts with oxygen to form aluminum oxide. It is one of the hardest compounds known; it is chemically stable and heat also resistant. It is a great protectant because it forms as quickly as it is scraped off

25. Rusting Iron Rust is produced when iron reacts with oxygen to form rust oxide. This new compound does not stick well to the existing metal and flakes off leaving the metal underneath it to rust. This process continues until all the metal is gone.

26. The Redox Reaction of Rust A corroding metal is a galvanic cell in which the anode and the cathode are found at different points on the same metal surface. The metal itself is the conducting material that allows the electrons to flow from the anode to the cathode. The anode is normally starts due to a scratch dent or impurity.

27. The Redox Reaction of Rust

28. Cathode O 2(aq) + 2H 2 O (l) + 4e -  4OH - (aq) Anode Fe (s)  Fe 2+ (aq) + 2e -

29. The Redox Reaction of Rust In the presence of oxygen and moisture, iron oxidizes releasing 2 e - The electrons then travel through the metal to the cathode, where they are used to reduce oxygen molecules. Water has two purposes in this reaction; it acts as a salt bridge for ions to flow and it also takes part in the reaction with hydroxide ions and the iron (III) ions to form rust.

30. The Redox Reaction of Rust

31. Factors that Affect the Rate of Corrosion Moisture Since water takes part in the reaction, it must be present for the reaction to occur. A relative humidity of at least 40% is needed for the reaction to take place.

32. Electrolytes When salts dissolve in water they become ions which increase the conductivity of water. The chlorine ions also act in a similar manner to a salt bridge in the way they offset the increase of Fe 2+ ions at the anode. The sodium ions play a similar role at the cathode as they help to offset the negative charge build up from the hydroxide ions.

33. Contact with Less Reactive Metals When two different metals come in contact with each other, the more reactive metal becomes oxidized. This is why metal fabricators must use the same type of metal when fabricating materials to avoid corrosion.

34. Mechanical Stress Bending, shaping, or cutting metal, stresses the structure of the metal which creates weak points. The weak points are then prone to corrosion.

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36. Preventing Corrosion Protective Coatings The simplest method of corrosion resistance is to cover the metal with a protective coating. Once the coating is scratched or exposed the metal will rust, even though the rust may appear to only be at the surface it can actually spread deep into the metal.

37. Protective Coatings

38. Galvanizing The process of coating iron or steel with a thin layer of zinc. This can be done by dipping the metal in a hot vat of molten zinc or by electroplating. When the zinc oxidizes it forms a tough, protective coating.

39. Cathodic Protection A form of metal corrosion prevention in which the metal being protected is forced to be the cathode of a cell, using either impressed current or a sacrificial anode.

40. Sacrificial Anode A form of protection where a metal that is more easily oxidized is attached to another metal to protect it. The more reactive metal acts as the anode, thus protecting the other metal by making it the cathode.

41. Sacrificial Anode This method does not require complete covering of the metal; all it needs is some sort of conductive connection that allows it to pass electrons to the metal that needs protecting. The sacrificial anode will need periodic replacement.

43. Impressed Current In this method, the metal needing protection is attached to the negative terminal of a power source making it the cathode. Continually pumping electrons into the cathode prevents corrosion

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