Oxidation-Reduction Reactions
Part 1: For each of the compounds listed below, write the oxidation number of each element in the compound.
1) H2O 11) MnO2
2) NiO2 12) NH2OH
3) Cd(OH) 2 13) H2O2
4) C2O42- 14) PBr3
5) Cr2O72- 15) ClO3-
6) MnO4- 16) HClO2
7) CN- 17) Pb(NO3)2
8) Cl2 18) CuSO4
9) H2CO3 19) N2O7
10) Fe(NO2)3 20) Pb3(PO4)4
Part 2: For each of the equations below, assign oxidation numbers to each element on both the reactant and product sides. Determine if the reaction is “redox” or “nonredox”, and list the reducing agent (RA) and oxidizing agent (OA) for all of the redox reactions.
1) 2H2(g) + O2(g)è 2H2O(g)
2) Zn(s) + 2H+(aq) è Zn2+(aq) + H2(g)
3) Cd(s) + NiO2(s) + 2H2O(l) è Cd(OH)2(s) + Ni(OH)2(s)
4) 2H2O(l) + Al(s) + MnO4-(aq) è Al(OH) 4-(aq) + MnO2(s)
5) 16H+(aq) + 2MnO4-(aq) + 5C2O42-(aq) è 2Mn2+(aq) + 8H2O(l) + 10CO2(g)
6) Cu(s) + 4H+(aq) + 2NO3-(aq) è Cu2+(aq) + 2NO2(g) + 2H2O(l)
7) I2O5(s) + 5CO(g) è I2(s) + 5CO2(g)
8) 2Hg2+(aq) + N2H4(aq) è 2Hg(l) + N2(g) + 4H+(aq)
9) 3H2S(aq) + 2H+(aq) + 2NO3-(aq) è 3S(s) + 2NO(g) + 4H2O(l)
10) Ba2+(aq) + 2OH-(aq) + H2O2(aq) + 2ClO2(aq) è Ba(ClO2)2(s) + 2H2O(l) + O2(g)
Wkst 2: Balancing Redox
Balance the following using the half-rxn method…
In acidic sol’n:
a) Cu + NO3- ® Cu2+ + NO
b) Cr2O72- + Cl- ® Cr3+ + Cl2
c) Pb + PbO2 + H2SO4 ® PbSO4
In basic sol’n:
a) Al + MnO4- ® MnO2 + Al(OH)4-
b) Cl2 ® Cl- + OCl-
c) NO2- + Al ® NH3 + AlO2-
Wkst 3: Electrochemical Cells – Practice Problems
1. Write half-cell reactions and give standard reduction potentials for these reduction reactions. (use table)
a) Cu+ è Cu
b) Cl2 è 2Cl-
c) Cu2+ è Cu
d) Ba2+ è Ba
e) Ag+ è Ag
2. Determine which metal will be the anode, and E°cell for the following sets of half-reactions:
a) Al3+ + 3e- ===> Al E° =-1.66 V
Au3+ + 3e- ===> Au E° =+1.50 V
b) Li+ + e- ===> Li E° = -3.05 V
Ag+ + e- ===> Ag E° = +0.80 V
c) Mg2+ +2e- ===> Mg E° = -2.37 V
Fe3+ + 3e- ===> Fe E° = -0.036 V
3. Two half-cells, one containing Fe2+ and Fe and the other containing Ag+ and Ag, are connnected to form a voltaic cell. Use a Standard Reduction Potential table to determine the direction of spontaneous reaction and the value for E°cell. Diagram the cell and label its parts. Give equations for the half reactions.
4. Use a Standard Reduction Potential table to determine the E°cell value for the spontaneous reaction of each pair of half-cells listed below:
a) Ag ===> Ag+ + 1e-; Fe2+ + 2e- ===> Fe
b) Mg ===> Mg2+ + 2e-; Sn2+ + 2e- ===> Sn
c) Li ===> Li+ + 1e-; Mn2+ + 2e- ===> Mn
d) Cr ===> Cr3+ + 3e-; Hg2+ + 2e- ===> Hg
e) Ni ===> Ni2+ + 2e-; Cu2+ + 2e- ===> Cu
5. Two half cells, one containing Ca2+ and Ca and the other containing Ag+ and Ag, are connected to form a voltaic cell. Use a Standard Reduction Potential table to determine the direction of spontaneous reaction and the value for E° cell. Diagram the cell and label its parts. Give equations for the half reactions.
6. Explain why a rechargeable battery can be considered a combination voltaic-electrolytic cell.
7. A voltaic cell is composed of the following half-cells:
Ca2+ + 2e- è Ca E° = -2.87 V
Fe3+ + e- è Fe2+ E° = +0.77 V
Write the reaction that takes place at the anode (oxidation) and the reaction that takes place at the cathode (reduction). Calculate the standard cell potential (E° cell).
8. What is the standard cell potential for a voltaic cell composed of the following half-cells:
Cu2+ + 2e- è Cu E° = +0.34 V
Ag+ + e- è Ag E° = +0.80 V
Write the reaction that occurs at the anode (oxidation) and the cell reaction that occurs at the cathode (reduction).
9. Is the following redox reaction spontaneous as written? (Hint: write oxidation and reduction half-reactions; look up the standard reduction potentials, and find E° cell; if E° is positive, the reaction is spontaneous!)
2Ag+ + Ni è 2Ag + Ni2+
10. Decide if the following redox reaction is spontaneous as written: (see hint in #9)
Cr3+ + Al è Cr + Al3+