1. Electrochemistry Electrons in Chemical Reactions

2. Examples: Car Battery Electrolysis of water Rusting of an old Chevette Cellular respiration What do all these things have in common? Electrochemistry is the study of how electrons are transferred in chemical reactions The chemical changes that occur when electrons are transferred between reactants are known as reduction-oxidation (REDOX) reactions The # of e- gained by one entity = the # of e- lost by another Example: 2Na (s) + Cl 2(g) 2NaCl (s) REDOX REACTIONS

3. Writing Half- Reactions Redox reactions can be separated into an oxidation reaction and a reduction reaction which are called half-reactions because they involve only one-half of the redox reaction taking place The half-reaction must be balanced by adding electrons to the side with the more positive charge Elements have a charge of zero The charge of an ion is written as part of the symbol ! eg. charge on N 3- is -3

4. Reduction Half-Reactions Electrons are GAINED in a reduction reaction and are written on the REACTANT side of the equation Reduction: Fe 3+ + 3 e - Fe (s) Cl 2 + 2 e - 2Cl - REDUCTION is a GAIN in electrons(RIG)

5. Oxidation Half-Reactions Electrons LOST in oxidation are placed on the PRODUCT side of the equation Oxidation: Mg Mg 2+ + 2e - 2F 1- F 2 (g) + 2e - OXIDATION is a LOSS in electrons(OIL)

6. Remembering Half-Reactions Oxidation = OIL Reduction = RIG So... OIL RIG

7. Now try these: Sulfur to sulphide ions Sodium ion to sodium metal Nickel to nickel ions (2 answers) + O OX (OIL) RED (RIG)

8. Reactivity of Metals and Their Ions The more stable a metal atom is, the more reactive it is as an ion. The more stable a metal ion is, the more reactive it is as a metal. If you create a list that organizes metal ions from most reactive to least reactive, you end up creating a list that organizes metal atoms from least reactive to most reactive as shown in the table below. The resulting table is a table of reduction half-reactions based on the reactivity of atoms or ions, or an activity series.

9. Activity Series for Metals and Metal Ions (page 80 in text, page 5 in data booklet)

10. Using the Activity Series The ranking of the reactivity of metal ions is shown on the left-hand side of the table, with the most reactive metal ions placed higher than less reactive metal ions. Which is more reactive – tin ions or iron ions? The right side of the table lists metals in order of reactivity. Metals lower in the series are more reactive than the metals higher in the series. Which is more reactive – copper or zinc?

11. Using the Activity Series The table is most useful for determining if a redox reaction is: Spontaneous – a chemical reaction that occurs without the addition of external energy or Nonspontaneous – a chemical reaction that does not occur without the addition of external energy How do we tell if a reaction is spontaneous or nonspontaneous …

12. Using the Activity Series

13. Spontaneous Redox Reaction Example

14. Zinc – copper reaction in real life !

15. Two potatoes are pierced by zinc and copper electrodes. A small clock runs on the voltage produced. Potato Clock

16. Voltaic Cells Applying our understanding of Redox Reactions

17. Voltaic Cells also called GALVANIC cells. electrochemical cells that are used to convert chemical energy into electrical energy energy produced by SPONTANEOUS redox reactions In a spontaneous redox reaction, electrons are transferred from the substance oxidized to the substance reduced If reactants are arranged in a certain way, these electrons can be made to move through a wire The reactants must be separated and yet in contact with each other so that the reaction will occur s l o w l y They can be separated by a salt bridge or a porous cup

18. Salt Bridge containing unreactive sodium sulfate

19. Porous Cup Unglazed porcelain cup Allows for ion transfer just like a salt bridge does!

20. Parts of a Voltaic Cell Half-cell - one part of a voltaic cell in which either oxidation or reduction occurs – contains a solid (electrode) and a solution (electrolyte) Electrodes - solid conductors that connect the electrolyte solution to the external circuit Anode: oxidation occurs here, electrons are produced negative (-) terminal Cathode: reduction occurs here, electrons are used positive (+) terminal Electrolytes – solutions that conduct electricity – must contain ions! Salt bridge or porous cup - allows the passage of electrons without contamination of the two half-cells

21. How does the electricity flow? How does a voltaic cell work? The electrons move from the anode to the cathode through a wire. This is the external circuit. Another circuit keeps the charges moving – the internal circuit is the flow of ions in the solutions through the salt bridge or porous cup. Cations (positively-charged ions) move to the cathode Anions (negatively-charged ions) move to the anode

22. Voltaic cells consist of Anode: Zn(s)  Zn 2+ (aq) + 2e (loses mass) Cathode: Cu 2+ (aq) + 2e -  Cu(s) (gains mass) As oxidation occurs, Zn is converted to Zn 2+ and 2e -. The electrons flow towards the cathode where they are used in the reduction reaction.

23. Cell Notation

24. Identify the anode and cathode!

26. Batteries A BATTERY is a self-contained VOLTAIC CELL. 1. Primary Batteries converts stored chemical potential energy into electrical energy when the two half cells within the battery are connected by an external circuit. usually the container is the anode and the graphite center is the cathode. the two are separated by a thick moist paste, which is the electrolyte. are not rechargeable

29. 2. Secondary Batteries are rechargeable by passing a direct current through it. Hg and Pb are often used in these batteries, are an environmental concern. cathode grill is filled with lead (IV) oxide, the anode grill is filled with spongy lead.

30. Electrolytic Cells

31. The Electrolytic Plant, which is the size of four football fields, consumes the same amount of power as a city of 250,000 people.

32. Copper is electroplated onto a strip of silver.

33. Note the differences….