Introduction

When elements changes oxidation steps during a redox reaction, elentrons are moved around in the process. This redistribution of electrons forms a current, that can be measured and calculated, and this is electrochemistry. An electrochemical process is therefore just a redox reaction.

The electrochemical reaction is a redox reaction not at equilibrium. This is an important detail, both in regards to the calculations, and in regards to understanding why batteries becomes flat and what happens when you recharge batteries. The redox reactionen, the driving force for obtaining equilibrium, is what provides the electrical current, and when the system is in equilibrium, there is no current. The batteries are flat. This also means that recharging is the same as bringing the reaction back out of equilibrium.

In practice, electrochemistry is not just about batteries. It is also the reason for galvanic corrosion of metals, e.g. in process equipment and plumbing that isn't done right. We will get back to this, when we start looking at the reactivity series.

First we take a look at how to describe the electrochemical elements.

If we start with a simple redox reaction for zinc and copper, which was the one the chemist John Frederic Daniell used for his invention called the Daniell cell:

Zn(s) + Cu²⁺(aq) Cu(s) + Zn²⁺(aq)

The Daniell cell is two solutions, separated by a salt bridge, so you have copper(II) ions with a copper rod and zinc(II) ions with a zinc rod. These two solutions are called half cells. In Daniell's original version, the copper and copper ions were in a pottery vase, surrounded by a zinc plate in a solution of zinc ions. The smart part about the setup, is that the two solutions are connected, to the electrons can move, but without the two solutions getting mixed. A salt bridge for connecting the two half cells are written as ║, and the two phases (the oxidized and the reduced) in each half cell are separated by │. A Daniell cell, written as a cell diagram, will look like this:

Zn(s) │ Zn²⁺(aq) ║ Cu²⁺(aq) │ Cu(s)

If you know the concentrations of the two ions, e.g. 1 M, you write it like this:

Zn(s) │ Zn²⁺(aq, 1 M) ║ Cu²⁺(aq, 1 M) │ Cu(s)


The general way to write a cell diagram looks like this:

red1 │ ox1 ║ ox2 │ red2

red is the reduced version of the metal, usually the metal in oxidation step 0, and ox is the oxidized version of the metal, which is usually the ionic form. The writing mirrors the reaction:

red1 + ox2 ox1 + red2


The orientation of the cell diagram dictates the orientation of the reaction equation. As the cell diagram is written, for the Daniell cell this specifies that Zn(s) and Cu(s) are the two poles on a battery and the redox reaction for the element is:

Zn(s) + Cu²⁺(aq) Cu(s) + Zn²⁺(aq)

If you reverse the cell diagram,

Cu(s) │ Cu²⁺(aq) ║ Zn²⁺(aq) │ Zn(s)

you also reverse the reaction equation, i.e.:

Cu(s) + Zn²⁺(aq) Zn(s) + Cu²⁺(aq)

This orientation is important when we are going to do calculations on which way the reaction is going, in order to produce a current. This will be shown in details in "Electromotive force (EMF)".

The two halves of the cell diagram, separated by ║, are called half cells. In the same manner as when splitting the total reaction into partial reactions, to work with them effectively, you split up the electrochemical reactions. The partial reactions are the respective oxidations and reductions for the atoms. For the general reaction

red1 + ox2 ox1 + red2

the reaction is divided into

1. The componen giving off electrons, thereby getting oxidized:
   red1 ox1 + e⁻
2. The component taking up electrons, thereby getting reduced:
   ox2 + e⁻ red2

For the Daniell cell

Zn(s) │ Zn²⁺(aq) ║ Cu²⁺(aq) │ Cu(s)

Zn(s) + Cu²⁺(aq) Cu(s) + Zn²⁺(aq)

the partial reactions are

1. Zn(s) Zn²⁺(aq) + 2e⁻    (oxidation)
2. Cu²⁺(aq) + 2e⁻ Cu(s)   (reduction)

The half cells we have looked at so far, has consisted of a solution of ions and a piece of metal. A natural question would be, whether you could have more components, e.g. in multiple phases, and you can. The standard hydrogen electrode consists of H⁺ and H₂ as the redox pair with a platinum wire as the solid phase where the redox reaction takes place. I writing it looks like this:

Pt(s) │ H₂(g) │ H⁺(aq)

If you have a half cell containing multiple ions, e.g. Fe(II) and Fe(III), both at 0.5 M concentrations, you write the half cell like this:

Fe(s) │ Fe²⁺(aq, 0,5 M), Fe³⁺(aq, 0,5 M)


We now have the electrochemical cells/reactions broken down in pieces we can work with in practice.