1. Electrochemistry

2. In 1983, the US Mint decided that they could no longer afford making pennies out of pure copper.
Zinc was much cheaper, and the chemists at the mint knew that they could easily cover the zinc with a coating of copper:
Zn(s) + Cu2+(aq) → Zn2+(aq) + Cu(s)
By placing zinc in a solution of copper ions, a redox reaction occurs spontaneously, and the zinc becomes copper plated.
Zn is oxidized to Zn2+
Cu2+ is reduced to Cu0

3. What would happen if you placed a copper penny in a zinc solution?
Zn2+(aq) + Cu(s) → ??
Not much! Zinc ions will not oxidize copper metal.
This reaction is non-spontaneous.
So although it is easy to plate copper on zinc, the reverse reaction is much more difficult to achieve.

4. Electrochemistry
The reaction of copper ions with zinc metal is an electrochemical process.
An electrochemical process is the conversion of chemical energy into electrical energy, or visa versa.
All electrochemical processes are redox reactions. In this case, electrons flow from the zinc metal to the copper, precipitating the copper.

5. Electrochemical Cells
These reactions can be used as sources of electrical energy.
To do this, you need to separate the oxidation and reduction half-reactions.
The electrons are forced to travel through and external circuit to reach their destination. These devices are known as electrochemical cells. More on these later.

6. The Activity Series

7. The Activity Series
Table J indicates the range of chemical activity for many metals and a few non-metals.
The most active metals will donate valence electrons to the ion of the metals below, which are less active.
The less active metals (lower) will not react with the ions of the metals above them.

8. Energy States of Valence Electrons
Each electron in each orbital of atoms has a specific energy
Valence electrons always have the highest energy;
In metals, valence electrons have very high energy, non-metals, less so.
The metals have low ionization energy because their valence electrons are in high energy states.

9. The Activity Series The most active metals are on top
The least active metals are on bottom
The metal above will donate electron(s) spontaneously to the ion of the metal below.
The metal below will not sponateously give up its valence electrons to the metal above.

10. Active Metals
The most active metals will spontaneously give electrons to the ions of less active metals
The valence electrons are “looking” for an orbital of lower energy; they will give off a photon and “drop” into an orbital around the less active metal.

11. Using the Activity Series
For example,
Li + Ba2+ 
Barium is below lithium, the reaction happens sponaneously.
Mg + K+ 
Potassium ion is above Mg, and will not receive electrons; the reaction is non-spontaneous

12. What would you expect to happen when a strip of lead is placed in an aqueous solution of magnesium nitrate?

13. Non-Metals and Table J
The most active non-metals are found at the top right corner of the periodic table..
The most active non-metals gain electrons – they don’t lose them!

14. Halogens and Table J
The more active non-metal will take electrons from the ion of the less active metal:
F2 + 2Cl- 
For the opposite reaction, nothing happens:
Cl2 + 2F- 

15. Half Reactions and Table J
Write oxidation and reduction half reactions for the following metals and ions.
You need to look up the oxidation numbers on your periodic table for each element to see how many electrons are lost.
For the reduction half-reactions, each metal will be reduced to a neutral atom.

16. Metal
Oxidation Half-Reaction
Ion
Reduction Half-Reaction Li
Li → Li+ + e-
Na+
Na+ + e- → Na Ca K+ Mg
Ca2+ Al
Ba2+ Cu
Cu2+ Ag
Ag+ Zn
Au3+

17. Given the balanced ionic equation: Zn(s) + Cu2+(aq) → Zn2+(aq) + Cu(s)
Which equation represents the oxidation half-reaction?
Zn(s) + 2e- → Zn2+(aq)
Zn(s) → Zn2+(aq) + 2e-
Cu2+ (aq) →Cu(s) + 2e-
Cu2+ (aq) + 2e- → Cu(s)

18. Assign oxidation numbers to all atoms in the following equations.
Write them directly above the element.
Determine which elements are oxidized and reduced, and write the half-reactions.

19. 3Ag + Au3+→ 3Ag+ + Au 2H2 + O2 → 2H2O 2KCl → 2K + Cl2
Reaction
Ox. ½ rxn
Red. ½ rxn
3Ag + Au3+→ 3Ag+ + Au
2H2 + O2 → 2H2O
2KCl → 2K + Cl2
3Mg +2Al3+→3Mg2++ 2Al
F2 + 2Cl- → 2F- + Cl2
Ca + Pb2+ → Ca2+ + Pb
2I- + Cl2 → I2 + 2Cl-

20. Li + Ba2+ →
Mg + K+ →
Mg + H+ →
Ag + H+ →
Zn + Ca2+ →
F2 + KBr →
Br2 + F- →
Sr + Cu2+ →
Pb + Au3+ →
Cs + Fe3+ →
Au + Ag+ →
Cu + Au3+ →

21. Acids are considered “corrosive” because most of them react with metals to produce hydrogen gas.
How many metals here will react with acid?
How many will not?

22. Activity Series of Metals
Part 1 The effect of adding acid to metals
Materials
Well plate, dropper bottles of HCl, Mg(NO3)2 , FeCl3, Zn(NO3)2, CuSO4 magnesium, copper, iron, and zinc filings

23. Procedure Add 1 “squirt” of HCl into a well plate.
Add a small scoop of Zn filings to the well and observe any reaction that occurs. Record your observations in Table 1 on your answer sheet.
Repeat step 2 with each of the other metals.

24. Part 2 Single Replacement Reactions of Various Metals
You have been supplied with solutions of metal compounds and small grains of the metals as a pure element.
For example, you have a solution of a copper salt that will supply you with Cu2+(aq) in one test tube and 1 ml of Zn2+(aq) in another.
Put a piece of zinc in the Cu2+(aq) and a piece of copper in the Zn2+(aq). Note that the Zn begins to discolor.
This is because metallic copper is being deposited on the zinc. What you cannot see is the zinc going into solution.

25. As the copper is deposited: Cu2+(aq) + 2e-→Cuo(s)
The zinc is dissolved: Zno(s) → Zn2+(aq) + 2e-
Look at the other test tube. There doesn’t seem to be anything happening. There isn’t.
This is because reactions that are spontaneous in one direction are never spontaneous in the reverse direction.

26. Part 1: Reduction of hydrogen by metals
Table 1. Reaction of metals with acid. Write the oxidation half reaction for each reaction.
If there is no reaction, you leave the box empty.
Note that for the polyvalent metals, Fe and Cu, the highest oxidation number is more common and is assumed to have occurred here.

27. Oxidation half-reaction
Metal
Reaction (y/n)
Oxidation half-reaction Fe Cu Zn Mg

28. What metals react with the acid and which did not? How could you tell?
Write the reduction half reaction that is common for all of the metals that reacted in the box below.

29. Write balanced chemical equations for each of the three reactions that occurred.

30. Which metals react most vigorously with the acid
Which metals react most vigorously with the acid? What happened to the metals that seemed to disappear?
What do you think will happen to the pH of the solution as the reaction proceeds? How could you test your prediction?
For any metal that does not react with the acid, explain why using Table J to justify your answer.

31. Part 2. Single Replacement Reactions with Two Metals
Mg2+
Cu2+
Fe3+
Pb2+
Zn2+
Mgo *
Cuo NR
Feo
Pbo
Zno R

32. What is a spectator ion? What happens to them in chemical reactions?
Write the oxidation half reaction and the reduction half reaction for all spontaneous reactions in the table below.
Write balanced net ionic equations for all of the reactions.
The first reaction is done for you.
Write the name of the spectator ion in the box next to the reaction.

33. Which metal seemed to be the least reactive
Which metal seemed to be the least reactive? Explain each of your “No Reactions” using Table J.
What happens when you add a metal to a solution containing its ion?
Each of the solutions used in this lab are electrolytes. Using your knowledge of atomic structure and properties, explain why this is so.

34. Zn → Zn2+ + 2e- Cu2+ + 2e- → Cu Zn + Cu2+ → Zn2+ + Cu
Oxidation half reaction
Reduction half-reaction
Balanced net ionic equation
Zn → Zn2+ + 2e-
Cu2+ + 2e- → Cu
Zn + Cu2+ → Zn2+ + Cu