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Balancing Redox Reactions
19.2
A. Identify the change in oxidation number of the designated element and determine if the change
is due to oxidation or reduction.
1. ! ! !
Sulfur in S2O8 ---> SO42-
!
The oxidation number changes from ! !
!
This is an example of: !
2. ! ! !
oxidation!!
Phosphorus in P4 ---> H2PO42-!!
!
The oxidation number changes from
!
This is an example of: !
3. ! ! !
(note: P4 is the elemental state for phosphorus)
!!
oxidation!!
!
Nitrogen in N2O4 ---> NO !
!
This is an example of: !
oxidation!!
to ! !
!
reduction
Technetium in TcO42- ---> Tc2+
!
The oxidation number changes from
!
This is an example of: !
!
!
to ! !
reduction
The oxidation number changes from ! !
5. ! ! !
!
reduction
!
4. ! ! !
to ! !
oxidation!!
!!
to ! !
!
reduction
Tungsten in WO3 ---> W3O8 (note: oxidation numbers are not always whole numbers)
The oxidation number changes from ! !
to ! !
This is an example of: ! oxidation!! reduction
!
1
6. ! ! !
Chromium in Cr2O72- ---> Cr3+
!
The oxidation number changes from ! !
!
This is an example of: !
7. ! ! !
oxidation!!
Manganese in MnO4- ---> MnO2
The oxidation number changes from ! !
!
This is an example of: !
oxidation!!
!
Chlorine in Cl- ---> ClO3-
The oxidation number changes from ! !
!
This is an example of: !
oxidation!!
to ! !
!
reduction
Sulfur in HSO3- ---> S2O62-
!
The oxidation number changes from ! !
!
This is an example of: !
10. !! !
to ! !
reduction
!
9. ! ! !
!
reduction
!
8. ! ! !
to ! !
oxidation!!
to ! !
!
reduction
Nitrogen in NO3- ---> NH4+
!
The oxidation number changes from ! !
!
This is an example of: !
oxidation!!
to ! !
!
reduction
2
B. Balancing Redox Equations
Use the oxidation-number method to balance the following redox equations:
11. !! !
!
CO + I2O5 !
12. !! !
!
Cl2 + NaOH ! -----> NaCl
13. !! !
!
SO2 + Br2 !+ H2O! ----->
14. !! !
!
HBrO3 !
15. !
!
-----> I2
-----> Br2 +
MoCl5 + S2- !
+
CO2
+
HClO
HBr +
H2SO4
H2O + O2
-----> MoS2
+
Cl- !
+ !
S
3
16. !! !
!
Al + OH- ! + H2O -----> H2
+
17. !! !
!
TiCl62- + Zn ! -----> Ti3+
Cl- +
+
AlO2 -
Zn2+
C. Half-Reactions
Write unbalanced oxidation and reduction half-reactions represented in each of the following
redox equations. Write the half-reactions in net ionic form if they occur in aqueous solutions,
removing any spectator ions.
18. !
PbO(s) !
+
NH3(g)! ----->
N2(g)!+!
H2O(l)!+!
Pb(s)
!
Oxidation half-reaction:!
!
!
!
!
!
!
!
!
Reduction half-reaction:!
!
!
!
!
!
!
!
19. !
I2(s)! +!
+!
NaI(aq)
!
Oxidation half-reaction:!
!
!
!
!
!
!
!
!
Reduction half-reaction:!
!
!
!
!
!
!
!
20. !
Sn(s)! +!
----->! SnCl2(aq)!
+!
H2(g)
!
Oxidation half-reaction:!
!
!
!
!
!
!
!
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Reduction half-reaction:!
!
!
!
!
!
!
!
Na2S2O3(aq)! -----> Na2S2O4(aq)!!
2HCl(aq)!
4
D. Acidic Solutions
!
Use the half-reaction method to balance the following redox equations that
! !
take place in an acidic solution.
! !
When complete, box the oxidizing agent and underline the reducing agent.
21.!
H2S(g)
+
22.!
Ce4+(aq)
23.!
MnO4-(aq)
Cl2(g)
+
+
----->!
S(s)! +!
ClO3-(aq)
----->!Ce3+(aq)!
Cl-(aq)
----->! Cl2(g)! +!
Cl-(aq)
+!
ClO4-(aq)
Mn2+(aq)
5
24.!
HgS(s)
25.!
BiO3 -(aq)
NO3-(aq)
+
+
Mn2+(aq)
+
Cl-(aq)
----->
----->
MnO4-(aq)
HgCl42-(aq)
+
+
NO(g)
+
S(s)
Bi3+(aq)
6