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Wize High School Grade 12 Chemistry Textbook > Electrochemistry

Balancing Redox Reactions: Half-Reaction Method

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Balancing Redox Reactions: Half-Reaction Method

The "half-reaction method" of balancing redox reactions is the most commonly used method :)

Balance the following redox reaction in basic conditions and assign oxidation numbers to all species.

Ag(s)+ Zn2+(aq) → Ag2O(aq) +Zn(s){\small \text{Ag}\left(\text{s}\right)+\ \text{Zn}^{2+}\left(\text{aq}\right)\ \rightarrow\ \text{Ag}_2\text{O}\left(\text{aq}\right)\ +\text{Zn}\left(\text{s}\right)}


Wize Tip
Balancing redox equations:

1. Check the oxidation numbers of each species by inspection (change in oxidation # means this is a redox reaction!)

2. Write the reduction and oxidation half-reactions.

3. Balance each half-reaction: start with atoms that are neither O nor H, balance oxygens with H2O, hydrogens with H+, and remaining charges with e−.

4. Multiply the reduction and oxidation half-reactions as needed so that each reaction exchanges the same number of e−.

5. Add the oxidation and reduction half reactions together and simplify.

6. Additional step in basic solution only: there can't be any H+ ions, so neutralize the H+ by adding the same # of OH- to each side.
Watch Out!
Always solve in acidic conditions first!
Then, if the question is asking for basic conditions do step #6!

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1. Check the oxidation numbers of each species by inspection (change in oxidation # means this is a redox reaction!)

Here we are told it's a redox reaction already, but we can still check oxidation #s to see what is oxidized and what is reduced:

  • Ag goes from having an oxidation # of
    0
    to
    +1
    . Therefore Ag is
    oxidized
    .
  • Zn goes from having an oxidation # of
    2+
    to
    0
    . Therefore Zn is
    reduced
    .


2. Write the reduction and oxidation half-reactions.

Oxidation Half Reaction:
Ag(s) -> Ag2O (aq)

Reduction Half Reaction:
Zn2+(aq) -> Zn(s)

3. Balance each half-reaction: start with atoms that are neither O nor H, balance oxygens with H2O, hydrogens with H+, and remaining charges with e−.

a) Balance everything that is not O and H:
ox:
2Ag(s) -> Ag2O(aq)

red:
Zn2+(aq) -> Zn(s)
(stays the same)
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ox: 2Ag(s) -> Ag2O(aq)

red: Zn2+(aq) -> Zn(s)


b) Balance O with H2O:
ox:
2Ag(s) + H2O -> Ag2O(aq)

red:
Zn2+(aq) -> Zn(s)
(stays the same since there are no Os to balance!)

c) Balance H with H+
ox:
2Ag(s) + H2O -> Ag2O(aq) + 2H+(aq)

red:
Zn2+(aq) -> Zn(s)
(stays the same since there are no Hs to balance!)

d) Balance charges with electrons!
ox:
2Ag(s) + H2O -> Ag2O(aq) + 2H+(aq) + 2e-

red:
Zn2+(aq) + 2e- -> Zn(s)


PAGE BREAK
ox:
2Ag(s) + H2O -> Ag2O(aq) + 2H+(aq) + 2e-

red:
Zn2+(aq) + 2e- -> Zn(s)

4. Multiply the reduction and oxidation half-reactions as needed so that each reaction exchanges the same number of e−.

Here, both half reactions have the same amount of electrons (2) so we don't need to multiply anything further since they already have the same # of electrons!

5. Add the oxidation and reduction half reactions together and simplify.



2Ag(s)+ H2O+Zn2+(aq)+2e−→Zn(s)+Ag2O(aq) +2H+(aq)+2e−{\small 2\text{Ag}\left(\text{s}\right)+\ \text{H}_2\text{O} + \text{Zn}^{2+} (aq)+\sout{2e^-}\rightarrow\text{Zn(s)}+\text{Ag}_2\text{O}\left(\text{aq}\right)\ +2\text{H}^+(\text{aq})+\sout{2e^-}}

Answer for acidic conditions:

2Ag(s)+ H2O+Zn2+(aq)→Zn(s)+Ag2O(aq) +2H+(aq){\small 2\text{Ag}\left(\text{s}\right)+\ \text{H}_2\text{O} + \text{Zn}^{2+} (aq)\rightarrow\text{Zn(s)}+\text{Ag}_2\text{O}\left(\text{aq}\right)\ +2\text{H}^+(\text{aq})}
If we selected this as our answer we'd be incorrect because the question is asking us to balance the redox reaction in basic conditions!

PAGE BREAK
2Ag(s)+ H2O+Zn2+(aq)→Zn(s)+Ag2O(aq) +2H+(aq){\small 2\text{Ag}\left(\text{s}\right)+\ \text{H}_2\text{O} + \text{Zn}^{2+} (aq)\rightarrow\text{Zn(s)}+\text{Ag}_2\text{O}\left(\text{aq}\right)\ +2\text{H}^+(\text{aq})}

6. Additional step in basic solution only: there can't be any H+ ions, so neutralize the H+ by adding the same # of OH- to each side.



2OH−(aq)+2Ag(s)+ H2O+Zn2+(aq)→Zn(s)+Ag2O(aq) +2H+(aq)+2OH−(aq){\small 2\text{OH}^-\text{(aq)}+ 2\text{Ag}\left(\text{s}\right)+\ \text{H}_2\text{O} + \text{Zn}^{2+} (aq)\rightarrow\text{Zn(s)}+\text{Ag}_2\text{O}\left(\text{aq}\right)\ +\fbox{2\text{H}$^+(\text{aq})+2\text{OH}^-\text{(aq)}$}}

2OH−(aq)+2Ag(s)+ H2O+Zn2+(aq)→Zn(s)+Ag2O(aq) +2H2O{\small 2\text{OH}^-\text{(aq)}+ 2\text{Ag}\left(\text{s}\right)+\ \sout{\text{H}_2\text{O} }+ \text{Zn}^{2+} (aq)\rightarrow\text{Zn(s)}+\text{Ag}_2\text{O}\left(\text{aq}\right)\ +\sout{2}\text{H}_2\text{O}}

2OH−(aq)+2Ag(s)+ Zn2+(aq)→Zn(s)+Ag2O(aq) +H2O{\small 2\text{OH}^-\text{(aq)}+ 2\text{Ag}\left(\text{s}\right)+\ \text{Zn}^{2+} (aq)\rightarrow\text{Zn(s)}+\text{Ag}_2\text{O}\left(\text{aq}\right)\ +\text{H}_2\text{O}}
This last equation is the correct answer! We have fully balanced the redox reaction in BASIC conditions!


checklist
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Practice: Balancing Redox Reactions

Balance the following redox reaction in an acidic solution.

HCOOH + MnO4- --> CO2 + Mn2+
checklist
Mark Yourself Question
  1. Grab a piece of paper and try this problem yourself.
  2. When you're done, check the "I have answered this question" box below.
  3. View the solution and report whether you got it right or wrong.
Balance the following redox reaction in an acidic solution.
HCOOH + MnO4- --> CO2 + Mn2+