Electrochemical Series

One thing that makes the periodic table so special is that each of the elements is different in one way or another. Although there are trends in some of their properties, they all differ ever so slightly. A consequence of this is that they all have different reactivities. We can see this in terms of how easily they give up their electrons - or in other words, how easily they are oxidised. The electrochemical series is a handy form of sorting species in this way, and is the basis behind redox reactions, fuel cells, and batteries.

• This article is about the electrochemical series in physical chemistry.
• We'll learn what the electrochemical series is, which will involve an introduction to half-cells, electrochemical cells, and standard electrode potential.
• You'll then see the electrochemical series in the form of a table.
• Lastly, we'll explore the applications of the electrochemical series.

Electrochemical series explanation

At the start, we put across the idea that some elements are better at giving up their electrons than others. This is a fundamental principle of chemistry, and is the basis behind all redox reactions. To explore this in more detail, we need to look at the electrochemical series. But first, you need a bit of background knowledge. For that, we'll start by learning about half-cells, electrochemical cells, and standard electrode potential.

Half-cells

Imagine putting a rod of zinc in a solution of zinc ions. You eventually form an equilibrium, in which some of the zinc atoms give up their electrons to form zinc ions. Here's the equation. You'll notice that it is conventional to write the ions and electrons on the left-hand side and the metal atom on the right:

${\mathrm{Zn}}^{2+}\left(\mathrm{aq}\right)+2{\mathrm{e}}^{-}\rightleftharpoons \mathrm{Zn}\left(\mathrm{s}\right)$

The zinc ions move into solution whilst the released electrons gather on the rod, giving the rod a negative charge. This creates a potential difference between the zinc rod and the zinc ion solution, known as an electrode potential. The exact position of this equilibrium and the exact potential difference of the system depends on the reactivity of zinc and how easily it gives up its electrons.

Now, consider what would happen if you put a rod of copper in a solution of copper ions. Some of the copper atoms react, giving up their electrons to form copper ions. Once again, you'll eventually form an equilibrium. But copper is less reactive than zinc and is not so good at giving up its electrons - we can say that copper is a worse reducing agent. It means that the potential difference of this copper system is less negative than the potential difference of the zinc system. We call the system created by putting a metal in a solution of its own ions a half -cell.

Electrochemical cells

Finally, take a moment to think about what would happen if you joined the zinc half-cell and the copper half-cell together with a wire and a salt bridge. Zinc is a better reducing agent than copper - it is more reactive and gives up its electrons more easily. This means that it has a more negative potential difference than copper does. It creates an overall potential difference between the two cells, which we can also call an electrode potential, showing the difference in how readily the two metals give up their electrons. The potential difference is measured by a voltmeter connected to the system.

We represent standard electrode potentials (E°) using half equations involving the element and its ions. Note the following:

• Although we've emphasised how standard electrode potentials tell you how easily an element gives up its electrons, they are written as reduction potentials - in other words, how easily an ion gains electrons. This means that we write the ions and electrons on the left-hand side and the combined element on the right-hand side.
• Hydrogen always has a value of zero, as it is the standard with which we measure all other electrode potentials.
• A more negative electrode potential means the element is more likely to give up its electrons. It is a better reducing agent and is more easily oxidised.
• A more positive electrode potential means the element is less likely to give up its electrons. It is a worse reducing agent and is less easily oxidised.

Here is the standard electrode potential for zinc, Zn:

${\mathrm{Zn}}^{2+}\left(\mathrm{aq}\right)+2{\mathrm{e}}^{-}\to \mathrm{Zn}\left(\mathrm{s}\right)\mathrm{E}°=-0.76\mathrm{V}$

The standard electrode potential value is negative. This means that zinc is a better reducing agent than hydrogen and is more easily oxidised.

Electrochemical series

Amassing the standard electrode potentials of different elements creates the electrochemical series.

The electrochemical series is the basis behind all kinds of modern fuel cells and batteries. But before we look at these applications, let's take a look at the electrochemical series itself in the form of a table.

Electrochemical series table

The moment you've been waiting for - here's the electrochemical series. We've shown it as a table of different half-cells and their standard electrode potentials.

Electrochemical series applications

We've learned what the electrochemical series is. Now let's consider some of its applications.

• Combining two half-cells with different electrode potentials creates an electrochemical cell. We can use this to generate electricity. Because of this, the electrochemical series is the basis of fuel cells and batteries.
• We can also use electrode potentials to calculate an electrochemical cell's overall cell potential.
• In addition, the electrochemical series allows us to predict the direction of a redox reaction and predict whether disproportionation reactions can occur. In general, electrons travel from a species with a more negative standard electrode potential to a species with a more positive standard electrode potential.
• We can also use it to identify strong and weak oxidising and reducing agents. In general, species with a negative electrode potential are good reducing agents and tend to lose electrons themselves.