# Introduction to electrochemistry

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##### Intros
###### Lessons
1. What is electrochemistry?
2. The electrochemical cell
3. Redox reactions
4. Introducing half-equations
5. Reducing and oxidising agents
6. The activity series
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##### Examples
###### Lessons
1. Identify whether reduction or oxidation is taking place in the half-equations.
State whether the half-equations below show reduction or oxidation, giving a reason for your answer.
1. Au $\,$$\,$ Au+ + e-
2. Na $\,$$\,$ Na+ + e-
3. Cl2 + 2e- $\,$$\,$ 2Cl-
2. Write simple half-equations showing reduction and oxidation taking place.
Write a balanced half-equation for the following chemical processes.
1. Bromide ions being oxidised back to elemental bromine.
2. Sodium metal being oxidised.
3. Oxygen gas being reduced to oxide ions.
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###### Topic Notes

In this lesson, we will learn:

• The definition of electrochemical cell, reduction and oxidation.
• How to write half equations showing oxidation and reduction.
• How to identify oxidizing and reducing agents in reactions.
• The activity series as a predictor of relative oxidative strength.

Notes:

• Electrochemistry is similar to acid and base chemistry. Think about some of the key points we saw in in the Acid-Base theory and Solubility equilibria chapters:
• Acids and bases are ionic substances that make ionic solutions when dissolved.
• Acids and bases react with each other and with metals to produce salts, which are also ionic substances that make ionic solutions when dissolved.
• Ionic solutions conduct electricity.
• Reactions between some acids and bases or salts can release lot of energy.

The reactions that release energy can be set up in chemical ‘circuits’ called electrochemical cells. An electrochemical cell is a system that converts chemical potential energy into electrical energy. The electrical energy (electricity!) is produced by the movement of electrons from one of the reactants to another. See example diagram below.

Because electrons are negatively charged, reactions in electrochemical cells will always result in at least two of the atoms or ions in the reaction changing their individual charge. One is the atom losing the electrons and one is the atom gaining these electrons.

• Any reaction where an atom or ion changes its individual charge is called a redox reaction. The word redox comes from reduction and oxidation, the two opposite effects happening in a redox reaction.
• Reduction is when any atom or compound gains electrons, hydrogen, or loses oxygen.
• Oxidation is when any atom or compound loses electrons, hydrogen, or gains oxygen.
Reduction and oxidation are complementary – one cannot happen without the other. This means that if there is a chemical being reduced in a reaction, there will be a chemical getting oxidized (that’s why the overall reaction is called redox!).

There is a trick to remember what oxidation and reduction do in chemical reactions – OIL-RIG:
• Oxidation Is Loss (OIL) of electrons or hydrogen.
• Reduction Is Gain (RIG) of electrons or hydrogen.

• Any redox reaction can be split up to show reduction and oxidation separately in a half-equation. To do this, you need to use electrons to show the atom either gaining or losing electrons. We use the symbol e- to do this.
For example, the two-electron oxidation of copper can be written in a half-equation:

Cu $\,$$\,$ Cu2+ + $\,$2e-

The reduction of chlorine can be written in a half-equation:

Cl2 + 2e- $\,$$\,$ 2Cl-

In half-equations, electrons in the reactants show reduction, and electrons in the products show oxidation.

The e- symbol for electrons is ONLY used in half-equations; do not use it in full equations. In full equations showing the ‘real’ process, those electrons are contained in the other chemical that is being reacted!

• Chemicals that reduce other chemicals are called reducing agents, while chemicals that oxidize other chemicals are called oxidizing agents.
An agent or agency is someone/something that does things – a reducing agent does reduction to other chemicals, so the reducing agent itself is oxidized! Do not get them confused!
We learned in lessons like Electron configuration 2 and Forming ions that some atoms tend to gain or lose electrons to obtain a full outer shell of electrons. If you know this already, then you already have a good idea of what elements are reducing and oxidizing agents!
• Most reducing agents are metals because metals, as we know, normally lose electrons in a reaction, becoming positive ions while the other species in the reaction gains those electrons.
• Most oxidizing agents are electronegative non-metals that, as we know, tend to gain electrons in reactions by ‘removing them’ from other elements. Fluorine is an excellent oxidizing agent, as is oxygen (that’s why it’s called oxidation!) and chlorine – these are the most electronegative elements in the table.

• The activity series is a basic list that shows the relative tendencies of metal species to be oxidised – that is, to go from the elemental metal to metal ion plus electron(s). This has a correlation with electronegativity, as group 1 and 2 metals are amongst the highest in the series, like calcium, potassium and lithium. Near the bottom are the heavier transition metals like mercury, platinum and gold.
The activity series is based on standard electrode potential (AKA standard reduction potential) which we will use in calculations in the later lesson Calculating cell potential (Voltaic cells).