**COMPLETE ALL WORK ON SEPARATE PAPER**
B1 Net Ionic Equations
For each of the following reactions, write and label the half reactions and the net ionic equation.
1. Fe(s) + Cu(NO3)2(aq) ® Cu(s) + Fe(NO3)2(aq)
2. Mg(s) + 2 NaCl(aq) ® 2 Na(s) + MgCl2(aq)
3. 3 Ag(s) + Al(NO3)3(aq) ® Al(s) + 3 AgNO3(aq)
4. 2 Al(CH3COO)3(aq) + 3 Sn(s) ® 3 Sn(CH3COO)2(aq) + 2 Al(s)
5. Cl2(g) + CaI2(aq) ® I2(s) + CaCl2(aq)
B1b Predicting Redox Reactions
For each of the following situations, determine the net redox reaction and state the spontaneity:
1. Aqueous solutions of tin (II) bromide and iron (III) nitrate are mixed.
2. A laboratory technician stores an aqueous solution of iron (III) chloride in a nickel plated container.
3. A chemistry teacher demonstrates the test for bromide ions by bubbling some chlorine gas cautiously through a sodium bromide solution.
4. Acidified potassium dichromate is added to a solution of tin(II) sulphate.
5. A solution of nickel(II) nitrate is stored in a copper container.
6. Solid sulphur is placed in sulphuric acid.
7. Hydrogen peroxide is slowly poured into a beaker of cobalt(II) bromide.
8. Solutions of sodium dichromate, sodium hypochlorite and sodium hydroxide are mixed in a beaker.
9. Sodium metal is added to some water in a typical demonstration of the reactivity of the alkali metals.
10. Iron is used in an environment (eg. ocean water) containing aqueous magnesium chloride.
11. A student uses hydrobromic acid to acidify a potassium dichromate solution for later use as an oxidizing solution.
12. Two students attempt to etch their initials on a copper plate using hydrochloric acid.
13. An iron bolt is exposed to air and water, a reaction which causes millions of dollars of damage each year.
14. A student uses copper electrodes to test the conductivity of a nitric acid solution.
15. An acidic solution of potassium dichromate is mixed with and aqueous solution of hydrogen peroxide.
B2 Generating Redox Tables
Use the following information to answer the next question.
1. Generate a table of relative strengths of oxidizing and reducing agents for the metals and metal ions in the data chart. Write all half-reaction equations as reductions and label the strongest oxidizing agent and the strongest reducing agent.
2. From the evidence given, set up a table of relative strengths of oxidizing and reducing agents. Write all half-reaction equations as reductions and label the strongest oxidizing agent and the strongest reducing agent.
Co(s) + Pd2+(aq) ® Co2+(aq) + Pd(s)
Pd(s) + Pt2+(aq) ® Pd2+(aq) + Pt(s)
Mg(s) + Co2+(aq) ® Mg2+(aq) + Co(s)
3. The following equations are interpretations of the evidence from the reactions of four metals with various cation solutions. Make a table of redox half-reactions and arrange the four metallic ions and the hydrogen ion in order of their decreasing tendency to react.
Cd(s) + 2 H+(aq) ® Cd2+(aq) + H2(g)
Hg(l) + 2 H+(aq) ® Hg2+(aq) + H2(g)
Be(s) + Cd2+(aq) ® Be2+(aq) + Cd(s)
Ca2+(aq) + Be(s) ® Ca(s) + Be2+(aq)
4. In an experiment, four metals were placed into test tubes containing various ion solutions. Their resulting behaviour is communicated by the equations below. List the oxidizing agents from strongest to weakest.
Pt(s) + 2 H+(aq) ® Pt2+(aq) + H2(g)
2 Ce(s) + 3 Ni2+(aq) ® 2 Ce3+(aq) + 3 Ni(s)
3 Sr(s) + 2 Ce3+(aq) ® 3 Sr2+(aq) + 2 Ce(s)
Ni(s) + 2 H+(aq) ® Ni2+(aq) + H2(g)
Use the following information to answer the next question.
5. Given the list of observations above, list the reducing agents from most reactive to least reactive.
B3 Oxidation Numbers
1. Determine the oxidation number for carbon in each of the following substances:
a) CO2 b) C12H22O11 c) CO32- d) HCO3-
2. Determine the oxidation number of chlorine in each of the following substances:
a) NaCl b) NaClO c) NaClO2 d) NaClO3
3. Determine the oxidation number of manganese in each of the following substances:
a) Na2MnO4 b) MnO2 c) KMnO4 d) Mn2O7
4. Determine the oxidation number of the element in bold type in each of the following compounds or polyatomic ions:
a) C2H6O b) OsO4 c) K2TaF7 d) UO2
e) Na2Cr3O10 f) As2S5 g) S2O3 h) S4O6
i) Ca(CrO2)2 j) H5IO6 k) H2C2O4 l) POCl3
5. For each of the following reactions, determine whether the element in bold type has been oxidized or reduced:
a) Mg(s) + Fe2O3(s) ® Fe(s) + MgO(s)
b) Na2CO3(s) ® Na2O(s) + CO2(g)
c) C2H5OH(l) + O2(g) ® CO2(g) + H2O(g)
d) H2SO4(aq) + KOH(aq) ® HOH(l) + K2SO4(aq)
e) CH4(g) + O2(g) ® CO2(g) + H2O(g)
B4a Writing Complex Half Reactions
Balance the following half reactions:
1. IO3-(aq) ® I2(s)
2. Cl2O2-(aq) ® Cl-(aq)
3. AsO2(s) ® AsO43-(aq)
4. Cl2(g) ® ClO-(aq)
5. P4(s) ® H2PO4-(aq)
B4b Balancing Redox Reactions – Half Reaction Method
Complete and balance the following redox reactions using the half-reaction method. Include the net ionic equation in its simplified form.
1. H2C2O4(aq) + MnO4-(aq) ® Mn2+(aq) + CO2(g)
2. AsO33-(aq) + BrO3-(aq) ® Br-(aq) + AsO43-(aq)
3. NH3(aq) + Cu2+(aq) ® NO3-(aq) + Cu+(aq)
4. Cr2O72-(aq) + Sn2+(aq) ® Cr3+(aq) + Sn4+(aq) + H2O(l)
5. H2SO3(aq) + HIO3(aq) ® H2SO4(aq) + HI(aq)
B4c Balancing Redox Reactions – Oxidation Number Method
Complete and balance the following redox reactions using the oxidation number method. Please use the lowest whole number ratio.
1. ___H+(aq) + ___S2-(aq) + ___NO3-(aq) ® ___S(s) + ___NO(g) + ___H2O(l)
2. ___H+(aq) + ___MoO3(s) + ___Zn(s) ® ___Mo2O3(s) + ___Zn2+(aq) + ___H2O(l)
3. ___H2O(l) + ___Al(s) + ___NO3-(aq) ® ___AlO2-(aq) + ___NH3(aq) + ___H+(aq)
4. ___H+(aq) + ___Zn(s) + ___NO3-(aq) ® ___Zn2+(aq) + ___N2O(g) + ___H2O(l)
5. ___H+(aq) + ___Cl-(aq) + ___Cr2O72-(aq) ® ___Cl2(g) + ___Cr3+(aq) + ___H2O(l)
B5 Redox Stoichiometry
1. A 2.75 g piece of aluminum is placed in 250 mL of iron(III) nitrate solution. Assuming that the reaction reaches endpoint, calculate the concentration of the Fe3+(aq) ions.
2. If 6.00 mol/L nitric acid is poured into a beaker containing 50.0 mL of 1.50 mol/L hydrogen peroxide, what volume of acid is needed to reach endpoint?
3. If 30.0 mL of acidic dichromate ion solution is poured into a beaker containing 50.0 mL of 0.400 mol/L tin(II) nitrate, calculate the dichromate ion concentration and the Sn4+(aq) concentration.
4. Bromine can be obtained by bubbling chlorine gas through sea water. The concentration of bromide ions in sea water is 0.00020 mol/L. What mass of chlorine gas is needed to oxidize all the bromide ions in 1000 L of water?
5. Copper (II) nitrate can be produced by reacting copper metal with concentrated nitric acid. What volume of 15 mol/L nitric acid is needed to react with 12.7 g of copper?
6. The copper (II) ions in a solution can be converted to copper metal by trickling the solution over iron. The reaction produces iron (II) ions from the scrap iron. If the process produces 25 L of solution containing 0.0020 mol/L of Fe2+(aq) ions, what mass of copper is produced?
7. In an experiment to analyze the iron in an iron ore sample, 0.05000 mol/L K2Cr2O7(aq) was used in an acidic solution to oxidize Fe2+(aq) ions to Fe3+(aq) ions. Use the following data to calculate the concentration of Fe2+(aq) in the solution:
volume of Fe2+(aq) solution………………………….25.0 mL
final buret reading (K2Cr2O7(aq))……………………48.7 mL
initial buret reading………………………………….3.7 mL
8. Another experiment was used to analyze the tin in a tin ore sample. The Sn2+(aq) ions in an acidic solution were oxidized to Sn4+(aq) by a 0.200 mol/L KMnO4(aq) solution. Use the following information to calculate the concentration of Sn2+(aq) in the solution.
volume of Sn2+(aq) solution………………………….10.0 mL
final buret reading (KMnO4(aq))…………………….39.3 mL
initial buret reading………………………………….1.8 mL
9. The following data were obtained from the titration of 24.0 mL of an acidified 0.200 mol/L SnCl2(aq) solution with a KMnO4(aq) solution. Calculate the molar concentration of the KMnO4(aq) solution:
Trial 1 Trial 2
final buret reading 20.41 mL 40.62 mL
initial buret reading 0.22 mL 20.41 mL
volume of KMnO4(aq) used
B6Electrochemical Cells
Draw the following electrochemical cells and calculate the cell potential. Label the following:
8
**COMPLETE ALL WORK ON SEPARATE PAPER**
- voltmeter
- anode
- cathode
- positive terminal
- negative terminal
- anode material
- cathode material
- electrolytes
- direction of electron flow
- direction of cation flow
- direction of anion flow
- reduction half-reaction
- oxidation half-reaction
- net reaction
- spontaneity
8
**COMPLETE ALL WORK ON SEPARATE PAPER**
1. Zn(s) / Zn2+(aq) // Cu2+(aq) / Cu(s)
2. Zn(s) / Zn2+(aq) // Cd2+(aq) / Cd(s)
3. Cd(s) / Cd2+(aq) // Cu2+(aq) / Cu(s)
4. Cd(s) / Cd2+(aq) // Cr2O72-(aq), Cr3+(aq), H+(aq) / C(s)
5. Zn(s) / Zn2+(aq) // MnO4-(aq), Mn2+(aq), H+(aq) / C(s)
6. Cu(s) / Cu2+(aq) // Cr2O72-(aq), Cr3+(aq), H+(aq) / C(s)
7. Zn(s) / Zn2+(aq) // Ag+(aq) / Ag(s)
8. Ni(s) / Ni2+(aq) // MnO4-(aq), Mn2+(aq), H+(aq) / C(s)
B7 Electrolytic Cells
For each of the following, determine the half-reactions and net reaction and the minimum voltage required.
1. An aqueous solution of potassium sulphate is electrolyzed.
2. An aqueous solution of lead (II) nitrate is electrolyzed.
3. A solution of aqueous sodium bromide and aqueous zinc chloride are mixed in an electrolytic cell using inert electrodes.
4. An aqueous solution of nickel (II) chloride is electrolyzed.
5. Electricity is passed through an aqueous solution containing sodium chloride, potassium bromide, and lithium iodide.
B8 Quantitative Study of Electrolysis
1. Over a period of 30.0 minutes, a nickel-cadmium battery supplied a current of 0.268 A to a calculator. Calculate the charge supplied by the battery.
2. How many moles of electrons were supplied by an electrochemical cell producing a current of 0.200 A for a period of one hour?
3. A transistor radio is turned on for two hours. The battery caused 0.500 mol of electrons to flow. What was the average current supplied to the radio?
4. How many moles of electrons pass through a bulb when 8600 C of charge are supplied?
5. If a 5.00 A current flows from an electrochemical cell and 0.280 mol of electrons leave the cell, how long was the cell in operation?
6. How long does it take a 0.500 A current to produce a charge of 5800 C?
7. Determine the amperage involved when 0.500 mol of electrons flow through a wire for 90.0 min.
8. An electrolytic cell containing molten chromium(III) chloride operated for 45.0 minutes. It was found that the mass of molten chromium metal formed was 1.56 g. Calculate the average current of the cell.
9. How long must a 0.250 A current run through an electrolytic cell containing a solution of iron(III) nitrate, so that 8.37 g of iron(II) ions are produced?
10. When a 7.50 A current is passed through molten nickel(II) chloride for 1.40 hours, what mass of solid nickel will collect on the cathode?
11. What is the average current required to produce 8.25 g of iodine at the anode of an electrolytic cell containing a solution of tin(II) iodide, if the cell operates for 150 minutes?
12. If a 2.80 A current runs for 4.00 hours through a solution of zinc sulphate, what mass of zinc solid will be produced?
13. If 9.72 g of magnesium was formed in an electrolytic cell using a current of 0.600 A, how long did the cell operate?
14. Determine the mass of magnesium deposited at the cathode of a molten MgCl2 electrolytic cell if 10.0 A flow through the cell for 9.65 h.
15. An electroplating firm wishes to plate 12.7 g of copper from a Cu(NO3)2(aq) solution onto a pair of baby shoes. If a 2.00 A current is used, calculate the time required. At which electrode would the shoes be attached?
16. If 76 g of fluorine are required, what current would have to flow for 10 h to produce the fluorine from molten NaF? At which electrode could this reaction occur?
17. If a current plates out 13.5 g of aluminum, what mass of magnesium would be plated out in the same time by the same current?
8