Download Name Redox and Electrochemistry

Name ______________________________________
Redox and Electrochemistry
Oxidation number
Assign the correct oxidation numbers to each atom in each of the following. Write the numbers directly
above the symbols in each formula, as in the samples. Show the math directly beneath each.
Samples:
+1
\
-2
/
-3
\
NH4+
H2O
2 ( x) + (-2) = 0
1.
K
2. K2Cr2O7
+1
/
X + (+4) = +1
+1 +6 -2
\
| /
Na2SO4
(+2) + X + (-8) = 0
RbCl
Na2O
H2O2
MgBr2
S8
NH4+
PO43-
NO2
1-
FeS
CrO4 2-
Identifying Redox reactions
Examine each reaction below, Write oxidation states above each atom. Write the type of reaction and
also write if it a redox reaction (Look for changes in oxidation state):
Examples:
0
0 +1 -2
H2 + O2  H2O
+1 -2
+2 -1
+2 –2
+1 -1
Na2S(aq) + MgCl2(aq)  MgS(s) + NaCl(aq)
synthesis
Double replacement
yes - Redox
Not redox
3. BaCl2 (aq) + 2 KIO3(aq)  Ba(IO3)2(s) + 2 KCl(aq)
4. H2CO3(aq)  H2O(l) + CO2(g)
5. Mg (s) + Br2(l)  MgBr2 (s)
6. NH4NO2 (s)  N2(g) + 2 H2O(l)
7. Zn(s) + AgNO3(aq)  Ag(s) + Zn(NO3)2(aq)
8. Short cut: What do all the redox reactions you identified above have in common? (i.e. what will you
look for in an equation in order to identify it as a Redox reaction? )
Oxidation and Reduction:
9. Define each term in your own words
Oxidation number OxidationReductionOxidizing agentReducing agent –
1) Redox in reactions: (DO IN PENCIL) For each of the following write in oxidation numbers for
each element above its symbol, then write the symbol for the species oxidized and species reduced, the
oxidizing agent and reducing agent to the right. (Hint: GIN – LIP)
Example:
0
0
+1
-2
H2
+
O2

H2O
|_____Lost 1 e-_|______________| |
|________Gain 2 e-_|
Then:
__ H2 +
10. ____H2
+
__O2 
____ Br2

Oxidized:
Reduced:
Oxidizing agent:
Reducing agent:
__ H2O
_____HBr
Ox ½:
Oxidized:
Reducing agent:
Reduced:
Oxidizing agent:
Red ½:
+____ HCl 
11. ____ Zn
____ ZnCl2
+
____ H2
Reduced:
Reducing agent:
Oxidizing agent:
Oxidized:
____ O2
Oxidizing agent:
Reduced:
Reducing agent:
Oxidized:
Ox ½:
Red ½:
12. ____ Al2O3

____ Al
+
Ox ½:
Red ½:
13. ____ HI
Ox ½:
Red ½:

_____ H2
+
____ I2
Reduced:
Oxidized:
Reducing agent:
Oxidizing agent:
H0
O0
O0
H0
14. ____Mg
+
____ N2

_____Mg3N2
Oxidizing agent:
Reducing agent:
Reduced:
Oxidized:
____ FeCl3 + ____ Br2
Oxidized:
Reduced:
Reducing agent:
Oxidizing agent:
Ox ½:
Red ½:
15. ____ FeBr3 + ____ Cl2 
Ox ½:
Red ½:
16. ____Al
+
____ Cl2

_____AlCl3
Ox ½:
Reducing agent:
Oxidized:
Reduced:
Oxidizing agent:
Red ½:
17. _____ K0 + ____ Cu2+  ____ K+ + ____ Cu0
Oxidized:
Reduced:
Ox agent:
Red agent:
Ox ½:
Red ½:
18. ____ Fe+2 + ____ Sn4+  ____ Fe 3+ + ____ Sn2+
Ox ½:
Reduced:
Red agent:
Oxidized:
Ox agent:
Red ½:
Writing half reactions: Write half reactions for the oxidation and reduction that occur in the equations above,
then balance: For simplicity, leave out subscripts where they occur. [Honors and other experts: Balance the half
reactions and use to balance the overall reaction]
Example:
Ox #s
Loss
And gain
0
0
+1
-2
H2
+
O2

H2 O
|_____Lost 1 e-_|______________| |
|________Gain 2 e-_|
Oxidation: (loss: electrons shown as product)
H0

H+
+
1e-
Reduction: (gain: electrons shown as reactant)
O0
+
2e-

O-2
notice balance of charge on both sides