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Wize University Chemistry Textbook > Chemical Reactions

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)
PAGE BREAK
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!

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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!


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Practice: Balancing Redox Reactions

Balance the following redox reaction in an acidic solution.

HCOOH + MnO4- --> CO2 + Mn2+