Why is Delta G negative for electrolysis reactions?

1 Answer
Jan 28, 2016

DeltaG^@>0 but after applying a potential E_(cell)>=2.06V from an external power source, DeltaG becomes negative and the reaction will be spontaneous.

Explanation:

Let us discuss the example of electrolysis of water.

In electrolysis of water, hydrogen and oxygen gases are produced.

The anode and the cathode half-reactions are the following:

Anode: 2H_2O->O_2+4H^(+)+4e^(-)" " "-E^@=-1.23V

Cathode: 4H_2O+4e^(-)->2H_2+4OH^-" "E^@=-0.83V

Net reaction: 6H_2O->2H_2+O_2+underbrace(4(H^(+)+OH^-))_(4H_2O)

2H_2O->2H_2+O_2" "E_(cell)^@=-2.06V

A negative cell potential implies non spontaneous process and therefore, DeltaG^@>0.

Note that the relationship between DeltaG^@ and E^@ is given by:

DeltaG^@=-nFE^@

where, n is the number of electrons transferred during redox, which is n=4 in this case,
and F=96485C/("mol "e^-) is Faraday's constant.

Therefore, since E^@<0 =>DeltaG^@>0

Because DeltaG^@>0, thus after applying a potential E_(cell)>=2.06V from an external power source, DeltaG becomes negative and the reaction will be spontaneous.

Note that, DeltaG=-nFE

Electrochemistry | Electrolysis, Electrolytic Cell & Electroplating.