1.3 Recognizing Redox Reactions
Oxidation numbers are a convenient way of identifying redox reactions and also indicating which element is oxidized and which is reduced. Here's an example - the reaction between sodium metal and chlorine gas:
2 Na |
+ |
Cl2 |
→ |
2 NaCl |
It is often useful to write the oxidation number for every element, in every compound, above the element in the equation. Thus for our reaction we have:
0 |
0 |
+1 -1 | ||
2 Na |
+ |
Cl2 |
→ |
2 NaCl |
Be sure to note that the balancing coefficients in the equation (the "2" in front of Na and in front of NaCl) do not affect the value of the oxidation numbers. We'll return to these coefficients soon.
A chart is a useful way for us to summarize the changes in oxidation number for each element:
element |
initial ox no |
final ox no |
change in electrons (e-) |
oxidized or reduced |
|
Na |
0 |
→ |
+1 |
lost 1 e- |
oxidized |
Cl |
0 |
→ |
-1 |
gain 1 e- |
reduced |
We see several important things in our table -
Loss of electrons is Oxidation (LEO)
Gain of electrons is Reduction (GER) |
An increase in oxidation number indicates oxidation A decrease in oxidation number indicates reduction |
Let's try another example. Identifying what is oxidized and what is reduced - and how many electrons are involved - is very important to master now for we will expand upon this skill shortly.
Consider the reaction
2 Mg |
+ |
O2 |
→ |
2 MgO |
Determine oxidation numbers for all elements in every compound:
0 |
0 |
+2 -2 | ||
2 Mg |
+ |
O2 |
→ |
2 MgO |
Summarize the changes, determine the number of electrons transferred per atom, and identify what is oxidized and what is reduced:
element |
initial ox no |
final ox no |
change in electrons (e-) |
oxidized or reduced |
|
Mg |
0 |
→ |
+2 |
lost 2e- |
oxidized |
O |
0 |
→ |
-2 |
gain 2e- |
reduced |
Two new terms before we continue. You will recall that we mentioned in the first section of this unit that oxidation cannot occur without reduction, and vice versa. The substance losing electrons (undergoing oxidation) gives its electrons to the substance gaining electrons (undergoing reduction). If the reduced substance will not accept electrons, the oxidized substance could not give away electrons. Thus, one allows for the other to occur.
Reducing agent the substance that is oxidized. Oxidizing agent the substance that is reduced. |
In our last example above, magnesium was oxidized, therefore it was the reducing agent.
Oxygen was reduced, therefore it was the oxidizing agent.By convention we often refer to the oxidizing agent and reducing agents as the entire compound the element is in, not just individual element. Consider the following reaction. Oxidation numbers are shown only for substances whose oxidation numbers undergo a change:
0 |
+5 (N) |
+2 (Zn) | +4 (N) | |||||
Zn |
+ |
HNO3 |
→ |
Zn(NO3)2 |
+ | NO2 | + | H2O |
Summarize:
element |
initial ox no |
final ox no |
e- |
oxidized or reduced |
Agent |
|
Zn |
0 |
→ |
+2 |
lost 2e- |
oxidized |
reducing agent: Zn |
N |
+5 |
→ |
+4 |
gain 1e- |
reduced |
oxidizing agent: HNO3 |
We see that HNO3 is referred to as the oxidizing agent, not just N.
One final thing to note in this example - notice that the nitrogen in Zn(NO3)2 did not undergo a change in oxidation number.
Here is a final example.
Consider the reaction
N2 |
+ |
2H2 |
→ |
2 NH3 |
Determine oxidation numbers for all elements in every compound:
0 |
0 |
-3 +1 | ||
N2 |
+ |
2H2 |
→ |
2 NH3 |
Summarize the changes, determine the number of electrons transferred per atom, and identify what is oxidized and what is reduced, and identify the oxidizing agent and reducing agent:
element |
initial ox no |
final ox no |
e- |
oxidized or reduced |
Agent |
|
N |
0 |
→ |
-3 |
gain 3e- |
reduced |
oxidizing agent: N2 |
H |
0 |
→ |
+1 |
lose 1e- |
oxidized |
reducing agent: H2 |
Check your understanding with the practice questions, set 2, then complete Assignment 1.