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Electrochemistry: Free Energy and Cell Potential Energy

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  • 0:03 The Galvanic Cell
  • 2:35 Measuring Cell…
  • 5:36 Free Energy & Cell…
  • 7:33 Lesson Summary
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Lesson Transcript
Instructor: Nicola McDougal

Nicky has taught a variety of chemistry courses at college level. Nicky has a PhD in Physical Chemistry.

Our modern lives are totally dependent on electricity. In this lesson, we learn about electricity spontaneously produced by electrochemical cells or batteries. We make the link between the potential energy they produce and Gibbs free energy.

The Galvanic Cell

Our lives would be unimaginable without electricity or batteries. Our lives literally depend on them. A battery is an electrochemical cell, which converts chemical energy into electrical energy. Here, we will focus on the galvanic cell, where a spontaneous chemical reaction produces electrical energy. This energy powers our cell phones, our laptops, our cars, pretty much everything!

A galvanic cell, also called a voltaic cell, makes use of reduction-oxidation chemical reactions. A redox reaction is divided into two half-reactions. One half-reaction involves the loss of electrons, and we say this is oxidized. The other half-reaction involves the gain of electrons, and we say this is reduced. I like to use the word OILRIG to help me remember which way this goes: oxidation is loss, reduction is gain. The number of electrons lost must be the same as the number of electrons gained. We must make sure we are balanced.

Here is an example of two half-reactions:

two chemical reactions

We have copper two plus gaining two electrons forming solid copper. From OILRIG, you know this is the reduction half-reaction. The second half-reaction is solid zinc losing two electrons forming zinc two plus. Again, OILRIG tells you this is the oxidation half-reaction. This is a balanced equation, as the number of electrons gained and lost is the same.

In a galvanic cell, the two half-reactions occur at two different electrodes, often metal or wire in a solution. Let us now look at a typical galvanic cell:

Galvanic Cell
diagram of galvanic cell

On the left, we are at the anode, and we have a zinc electrode dipped in zinc two plus solution. An Ox lives here because electrons are produced at the anode during oxidation. These two electrons move through the wire to where the Red Cat lives. Here a reduction takes place at the copper cathode.

A salt bridge completes the circuit, and we have a voltage being produced. Different half cells produce different voltages. This voltage is known as the cell potential energy, or E cell, and is a measure of the spontaneity of the reaction.

Measuring the Cell Potential Energy

The driving force behind a galvanic cell is the cell potential energy. The larger the cell potential, the more work we can get out of the cell. So how do we know how much energy we are going to get out? In each different cell you put together you will have two half-reactions. One is the oxidation, An Ox, and one is the reduction, Red Cat.

Each will have a standard voltage associated with it. These are Eox (or E oxidation) and Ered (or E reduction). The value depends on the reaction taking place there. We can simply look up our reaction in a standard potential table and calculate the overall cell potential, Ecell. Before we go ahead and do that, let us take a closer look at the standard potential table.

Standard Potential Table
standard potential table

The first thing to notice is that standard potentials are all shown as reduction reactions. This is just convention. In reality, you will have one reduction and one oxidation reaction. Please do not change the reaction around or the sign of the number. You just find your reaction in the reduction form and use the number as it is written. You don't even need to worry about the number of electrons transferred. Just use the number given to you. For our cell, we will use the zinc two plus and the copper two plus reactions.

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