Standard cell potential | Applications of thermodynamics | AP Chemistry | Khan Academy

Standard cell potential, which is also called standard cell voltage, refers to the voltage of an electrochemical cell when reactants and products are in their standard states at a particular temperature. For a zinc-copper galvanic cell, solid zinc reacts with copper (II) ions to form solid copper and zinc (II) ions. Standard cell potential is symbolized by E° (not) of the cell, where the superscript “°” refers to the fact that reactants and products are in their standard states.

For a solid, the standard state refers to the pure solid under a pressure of one atmosphere. So, we're talking about pure solid zinc and pure solid copper. For a solution, the standard state refers to a one molar concentration. Therefore, our concentration of copper (II) ions in solution is one molar, and so is our concentration of zinc (II) ions. For this zinc-copper cell at 25°C, the standard cell potential is equal to positive 1.10 Volts.

Now that we've talked about the standard cell potential for a zinc-copper cell, let's look at a diagram of this galvanic cell. For this cell, the anode compartment contains solid zinc metal and a one molar concentration of zinc (II) ions in aqueous solution. The anode is where oxidation takes place. So, at the zinc electrode, the solid zinc is converted into zinc (II) ions, and two electrons are lost. This is the oxidation half-reaction for this cell.

The electrons that are lost in the oxidation half-reaction travel through the wire. So, there's a flow of electrons from the zinc electrode toward the copper electrode. The cathode compartment contains the solid copper electrode and in solution, a one molar concentration of copper (II) ions. The cathode is where reduction takes place. So, at the surface of the copper electrode, the copper (II) ions in solution gain two electrons and turn into solid copper. This is the reduction half-reaction for this cell.

If we were to attach a voltmeter to our two electrodes, the voltmeter would read 1.10 volts, and this voltage is the standard cell potential. So, E of the cell is equal to positive 1.10 Volts. The standard cell potential depends on the potentials for the two half-reactions that make up the cell.

If we know the potentials for the half-reactions that make up the cell, we can calculate the standard potential of any galvanic cell. Let's see how we could do that for this particular cell. To calculate the standard cell potential for our zinc-copper cell, we're going to use what are called standard reduction potentials. Standard reduction potentials are symbolized by E° (not) of reduction, and these refer to the voltages for half-reactions that are written as reduction half-reactions.

For example, if copper (II) is reduced by two electrons to form solid copper, the voltage for this half-reaction, or the standard reduction potential, is equal to positive 0.34 volts. For the reduction of zinc (II) ions by two electrons to form solid zinc, the standard reduction potential is equal to 0.76 Volts. We just saw from our zinc-copper cell diagram that copper (II) ions are reduced to form solid copper.

So, we're going to leave this half-reaction as it's written. However, in our diagram for our zinc-copper cell, we saw that solid zinc was actually being oxidized to form zinc (II) ions. So, we need to rewrite this second half-reaction as an oxidation half-reaction. That just means reversing everything.

If we're reversing our half-reaction, we have to change the sign on the voltage. So, if the potential is negative 0.76 Volts for the reduction half-reaction, it would be positive 0.76 Volts for the oxidation half-reaction. Now, I've changed it to show the oxidation half-reaction: solid zinc turning into zinc (II) plus ions in solution and the loss of two electrons. The standard oxidation potential for this half-reaction is positive 0.76 volts.

The next step is to add the two half-reactions together. Here are all of our reactants, and here are all of the products. Notice how two electrons would cancel out. That gives us copper (II) ions in aqueous solution plus solid zinc, forming solid copper and zinc (II) ions in solution. Since we were able to get the overall equation by adding the two half-reactions together, we should be able to get the overall standard cell potential, E° of the cell, by adding together the standard potentials for our two half-reactions.

So, E° of the cell is equal to E° of the reduction half-reaction plus E° of the oxidation half-reaction, which is equal to 0.34 + 0.76 or positive 1.10 volts. So, the standard cell potential of our zinc-copper cell is equal to positive 1.10 Volts.

At 25°C, we can also calculate standard cell potential using a slightly different form of the equation that we just learned. Remember that we flipped the sign of the standard reduction potential to get the standard oxidation potential, which we plugged into this equation, and we're able to calculate the standard potential of the cell. Instead of flipping the sign ourselves and getting a standard oxidation potential, we could use this new equation to just use reduction potentials to calculate the standard potential of the cell.

Notice that this new equation has a negative sign that essentially flips the sign of the standard reduction potential for the oxidation process. So, it accomplishes the same thing that we did ourselves in the first equation. To find the standard potential of the cell, we take the standard reduction potential for the reduction process, and from that, we subtract the standard reduction potential for the oxidation process.

So, this equation is a bit of a shortcut. When we plug in our standard reduction potentials of positive 0.34 Volts and 0.76 Volts, the negative sign means we end up with the same answer we got before: positive 1.10 volts.