Electrochemistry Tutorial
Sections 20.1-20.2
Detailed solutions will be posted on the notice board.
1. a. Define oxidation and reduction. Define oxidizing and reducing agent
b. Identify the oxidizing agent in the equation below.
Fe(s) + 2H+(aq)
Fe2+(aq) + H2(g)
2. Determine the oxidation number for every element in each of the following:
a) CaH2
b) K2CrO4
3. Which element is oxidized, reduced or remains unchanged in the following reaction? Which
compound serves as the oxidizing agent?
a) H2S(g) + 4H2O2(aq)
H2SO4(aq) + 4H2O(l)
4. Balance these half-reactions.
a) C2O42-(aq)
b) Cr3+(aq)
CO2(s)
HCrO4-(aq) (acidic)
c) MnO4-(aq)
MnO2(s) (basic)
5. Balance the following oxidation-reduction equation using the method of half-reactions.
CrO42-(aq) + Br2(l)
Cr(OH)3(s) + BrO-(g)
(basic)
Electrochemistry Tutorial
Sections 20.3-20.5
You will need Appendix E from the textbook to complete these questions.
1. The two half-reactions in a voltaic cell are:
Co3+(aq) + eClO3-(aq) + 6H+(aq) + 6e-
Co2+(aq)
Cl-(aq) + 3H2O(l)
a) Indicate which reaction occurs at the anode and which at the cathode.
b) Which species will be consumed in the reaction?
c) Which reaction takes place at the positive electrode?
2. Write the equation for calculating cell emf at standard conditions. Define “standard conditions”.
3. Using Appendix E, calculate the following and tell whether each reaction is spontaneous:
a) Standard reduction potential for the manganese half reaction when Eºcell = -0.777V
Mn2+(aq) + Cd(s)
Mn(s) + Cd2+(aq)
b) Standard cell potential for
Sn(s) + Sn4+(aq)
2Sn2+(aq)
4. Using Appendix E, rank the following from strongest to weakest reducing agent: I-(aq), Mg(s),
Co(s), N2H4(aq).
5. Consider the following reaction:
3Ni2+(aq) + 2Cr(OH)3(s) + 10OH-(aq)
3Ni(s) + 2CrO42-(aq) + 8H2O(l)
a) What is the value of n (in ΔG˚ = -nFE˚) for this reaction?
b) Is the reaction spontaneous?
Electrochemistry Tutorial
Sections 20.6-20.9
Solutions will be posted on the notice board.
1. Define primary and secondary cells. Give an example of each.
2. During a period of discharge, 81.4g of Pb from the anode of a lead-acid battery is converted into
PbSO4(s). How much PbSO4(s) is formed?
3. Describe how iron corrosion can be prevented. Which of the following could provide anticorrosion protection to Fe: Al, Cu, Ni, Zn?
4. The electrolysis of molten LiCl(l) produces Li(s) and Cl2(g). What is the minimum external emf
needed to drive this electrolysis? Assume standard conditions.
5. Calculate the number of grams of Li produced in question #4 in 0.5 hours if the applied electrical
current is 12.0A.
Chemistry of the Environment Tutorial
Chapter 18
Solutions will be posted on the notice board.
1. Name the regions and boundaries of the atmosphere, indicating the altitude intervals for each one.
2. Which is most reactive: oxygen, nitrogen, or argon? Explain using Lewis Structures.
3. Describe the difference between photodissociation and photoionization. What is the chemical
formula for ozone? Does photodissociation or photoionization play a larger role in the destruction
of ozone in the stratosphere? Show how this occurs naturally using equations.
4. Draw the Lewis Dot Structure for CFCl3. Suggest a structure for a CFC that could have been
replaced by CH2FCF3.
5. Define desalination and describe the process of reverse osmosis. Why is reverse osmosis a better
option than simple distillation?
6. Name three greenhouse gases and give a source for each. (Note: there are more answers for this
question than are given on the notice board.)