Chemistry
Objectives - Chapter 20 and 21
Goals : To gain an understanding of :
1.Redox reactions.
2.The role of oxygen in redox reactions.
3.Electrochemistry.
4.Electrical cells.
5.Electrolytic cells.
Objectives : The student will be able to :
1.Describe oxidation-reduction reactions.
- Define "oxidation" and "reduction" in terms of :
-gain or loss of oxygen
-gain or loss of electrons
-gain or loss of hydrogen
-and in the shifting of electrons in covalent bonds.
3.Give two examples of common redox reactions.
4.Define "oxidizing agent" and "reducing agent" and be able to identify oxidizing and reducing agents in redox reactions.
5.Identify the oxidized and reduced substances in redox reactions.
6.Describe the corrosion of iron and name and explain two methods for preventing iron corrosion.
7. Explain why corrosion is faster in the presence of acids or salts.
8.Determine the oxidation numbers of elements in their combined or uncombined states.
9.Define "oxidation" and "reduction" in terms of a change in oxidation number.
10.Use the changes in oxidation state to identify substance being reduced or oxidized.
11.Use the oxidation number change method and/or half reaction method to balance redox reactions
12.Identify redox and non-redox reactions.
13.Define "electrochemistry" and state three uses of electrochemistry.
14.Describe the electrochemical process and relate electrochemical processes to redox reactions.
15.Define "electrochemical cell," "electrode," "anode" and "cathode."
16.Explain, using a diagram which includes cathode and anode, how a voltaic cell produces a current.
17.Use the shorthand method to represent electrochemical cells.
18.Distinguish between dry cells, batteries and fuel cells and give examples of each.
19.Explain how lead battery functions and describe the reaction in which a lead battery is recharged.
20.Explain how a fuel cell produces electricity and state the advantages and disadvantages of fuel cells.
21.Given a voltaic cell, identify the half-cell in which oxidation occurs and the half-cell in which reduction occurs.
22.Define "cell potential" and calculate cell potentials.
23.Distinguish between standard cell potential and cell potential and state how standard cell potentials are determined.
24.Describe the process of electrolysis and explain the use of electrolytic cells.
25.Describe the electrolysis of water.
26.Describe the electrolysis of brine and state its importance.
27. Describe the electrolysis of molten sodium chloride and state its importance.
28.Describe and explain the processes of electroplating, electroforming, electrowinning, electrorefining and electromachining.
Rules for assigning oxidation states :
1. Monatomic ions = charge (e.g. Na+ = +1)
2. Hydrogen in compounds = +1 (except hydrides = -1 e.g. LiH)
3. Oxygen in compounds = -2 (except peroxides = -1, e.g. H2O2)
4. Uncombined elements = 0
5. Sum of charges of a neutral compound = 0
6. Sum of charges on polyatomic ions = charge on ion.
1. Determine in which half cell oxidation occurs, in which half cell reduction occurs, and to write the half reactions and the shorthand notation of the reaction. (Hint : More reactive metal is oxidized - see Table 21-1 on p. 664).
a. half cell : copper and copper(II)nitrate - Cu(NO3) 2
half cell : magnesium and magnesium nitrate - Mg(NO3) 2
anode - oxidation half reaction :
cathode - reduction half reaction :
shorthand notation
b. half cell : tin and tin(II)sulfate - SnSO4
half cell : aluminum and aluminum sulfate - Al2(SO4) 3
anode -
cathode -
shorthand notation :
c. half cell : lead and lead(II) nitrate - Pb(NO3) 2
half cell : iron and iron(III) nitrate – Fe(NO3)3
anode -
cathode -
shorthand notation :
d. half cell : nickel and nickel(II) sulfate - NiSO4
half cell : zinc and zinc sulfate - ZnSO4
anode -
cathode -
shorthand notation :
e. half cell : lead and lead(II)nitrate - Pb(NO3) 2
half cell : copper and copper(II) nitrate - Cu(NO3) 2
anode -
cathode -
shorthand notation :