Redox

Redox Reaction Definition

The term redox is short for reduction-oxidation, and is used to describe reactions involving - you guessed it - both oxidation reactions and reduction reactions. Redox reactions all feature a movement of electrons and a change in oxidation state.

Let’s look more closely at the definitions of oxidation and reduction.

Oxidation and reduction

The words oxidation and reduction have a few different meanings in chemistry. The first definition looks at them in terms of oxygen. Take a punt - you can probably guess what oxidation means.

You can think of reduction as the opposite of oxidation.

For example, when copper reacts with oxygen, it forms copper oxide. The copper is oxidised.

Cu(s) + ½ O₂(g) → CuO(s)

But reacting hydrogen with copper oxide separates the copper and the oxygen. The copper oxide is reduced.

CuO(s) + H₂(g) → Cu(s) + H₂O(l)

Did you notice that we added hydrogen to reduce copper oxide? This leads us to the second set of definitions for oxidation and reduction.

However, in chemistry, we tend to use a different definition. It refers to the movement of electrons between species in a reaction, and it is the definition we'll focus on for the rest of this topic.

There’s a handy acronym that will help you remember this third definition: OILRIG.

The acronym OILRIG. StudySmarter Originals

Let’s revisit the example from before. What happens when copper reacts with oxygen? It forms copper oxide, an ionic compound. Copper oxide is made up of copper ions and oxygen ions, Cu²⁺ and O^2- respectively. To form these ions from neutral atoms, we need to move some electrons around.

• To turn a copper atom into a copper ion, the atom must lose two electrons. Copper is therefore oxidised.
• To turn an oxygen atom into an oxygen ion, the atom must gain two electrons. Oxygen is therefore reduced.

Because both oxidation and reduction are happening side by side, this is an example of a redox reaction.

In summary, oxidation can mean:

• Gain of oxygen.
• Loss of hydrogen.
• Loss of electrons.

Likewise, reduction can mean:

• Loss of oxygen.
• Gain of hydrogen.
• Gain of electrons.

Oxidising and Reducing Agents in Redox

We know what oxidation and reduction reactions are. Now let’s look at the species that carry out these reactions.

Oxidising agents

Oxidising agents take electrons from another species - they oxidise it. They are also called oxidants. Some particularly strong oxidising agents are fluorine and, perhaps unsurprisingly, oxygen.

Reducing agents

We now know what oxidising agents are. What do you think reducing agents are? You can probably take a good guess.

Reducing agents donate electrons to another species - they reduce it. They are also called reductants. Many metals, such as lithium, aluminium, and zinc, are good reducing agents, and so is hydrogen gas (if it is in the presence of a nickel catalyst).

Once again, there is a handy acronym that will help you remember the actions of oxidising and reducing agents in terms of electrons: RAD OAT.

RAD OAT. StudySmarter Originals

Take the example of copper and oxygen again. We know that copper is oxidised and oxygen is reduced. Copper loses two electrons, which oxygen gains. This also means that copper acts as a reducing agent and oxygen acts as an oxidising agent.

You should now feel confident about the terms oxidation, reduction, oxidising agent, and reducing agent. Let's move on to our next topic.

Oxidation States in Redox

Now that we know what redox reactions are, we can look at how to work out which species is oxidised and which species is reduced in a reaction. To do this, we use oxidation states.

We can use changes in oxidation states to see whether species have been oxidised or reduced. Going from a positive oxidation state to a more negative oxidation state means that the species has gained electrons. It has therefore been reduced. On the other hand, going from a negative oxidation state to a more positive oxidation state means that the species has lost electrons. It has therefore been oxidised.

Assigning oxidation states

It is all well and good knowing what an oxidation state is, but how do you assign them? There are a few rules you can use to find out a species' oxidation state.

• All uncombined elements have an oxidation state of 0.
  • e.g. Cl₂, Zn, and O₂ all have oxidation states of 0.

• The oxidation states of all the atoms or ions in a neutral compound add up to 0.
  • e.g. In the neutral compound NaCl, Na has an oxidation state of +1, and Cl has an oxidation state of -1. These add up to make 0.

• The sum of the oxidation states in an ion equals the charge of the ion. This works for both monatomic ions made from one atom, and complex ions made from lots of atoms.
  • e.g. Cl^- has an oxidation state of -1 and Ca²⁺ has an oxidation state of +2.
  • e.g. in the negative nitrate ion NO₃^-, N has an oxidation state of +5 and the three O each have an oxidation state of -2. These add up to make -1, which is the charge of the ion.

• In an ion or a compound, the more electronegative atom generally has the negative oxidation state.
  • e.g. in F₂O the oxidation state of F is -1, and the oxidation state of O is +2.

• Some elements take certain oxidation states:
  • Group 1 metals always have an oxidation state of +1
  • Group 2 metals always have an oxidation state of +2
  • Al always has an oxidation state of +3
  • H usually has an oxidation state of +1 (except in metal hydrides)
  • F always has an oxidation state of -1
  • Cl usually has an oxidation state of -1 (except in compounds with O or F)
  • O usually has an oxidation state of -2 (except in peroxides and compounds with F)

We can use this knowledge to calculate oxidation states of unknown elements in compounds and ions. Here's an example:

If you want to see further examples of oxidation states in action, head over to "Oxidation Number" for more worked problems.

Roman numerals in redox

Although some elements only form ions with one oxidation state, some elements can form ions with multiple different oxidation states. To avoid confusion when talking about these ions, we show oxidation states using roman numerals.

For example, when talking about the copper ions in CuO with an oxidation state of +2, we would write copper(II).

Redox Equation

During a redox reaction, two simultaneous processes occur - reduction and oxidation. We can show these processes using one overall equation that ignores any ions that aren’t oxidised or reduced - that is to say, aren't taking part in the reaction. These are known as spectator ions.

Writing redox equations

We'll now focus on how you can write redox equations, using some worked examples. Here’s one such example: the displacement reaction between magnesium and copper sulfate.

Half Equations

Redox equations are useful for showing an overall redox reaction. One species is oxidised whilst another is reduced, meaning that there is an overall movement of electrons. However, they can make it tricky to identify the individual oxidation and reduction processes. To see those more clearly, we often use half equations.

Writing half equations

To write half equations, we consider each of the ions or atoms involved in the redox equation separately. We add in electrons to show the processes of oxidation and reduction, and might also have to add in water or hydrogen ions to balance the equation.

These steps should help you learn how to write half equations.

1. Pick an atom or ion involved in the redox reaction and write out the reactants and products involving it.
2. Balance the elements apart from oxygen and hydrogen. Like all equations, half equations must be balanced - you must have the same number of moles of each element on both sides of the equation.
3. Add in water molecules to balance the oxygen atoms on both sides of the equation.
4. Add in hydrogen ions to balance the hydrogen atoms on both sides of the equation.
5. Add in electrons to balance the charges.

This process may sound a little tricky, but don't worry - we've got lots of examples over at "Half Equations" to get you started.

Disproportionation Reactions in Redox

Let’s now touch on disproportionation reactions. Before, we only looked at equations where a species was either oxidised or reduced. In disproportionation reactions, both processes occur.

We can see whether a species has been reduced, oxidised, or both by looking at its oxidation states. Here’s an example:

${\mathrm{Cu}}_2\mathrm{O}+{\mathrm{H}}_2{\mathrm{SO}}_4\to {\mathrm{CuSO}}_4+\mathrm{Cu}+{\mathrm{H}}_2\mathrm{O}$

We can see the following:

• Copper goes from an oxidation state of +1 in Cu₂O to O in Cu and +2 in CuSO₄.
• To get from Cu₂O to Cu, copper must gain an electron. This means that it is reduced.
• To get from Cu₂O to CuSO₄, copper must lose an electron. This means that it is oxidised.

Because copper has been both oxidised and reduced, this is a disproportionation reaction. This is just one example of a type of redox reaction. Let’s now look at some more.

Redox Reactions Examples

There are a few common examples of redox reactions. One of these includes electrolysis.

Electrolysis

Electrolysis is a way of splitting ionic compounds into simpler substances by passing an electrical current through them. Either reduction or oxidation takes place at the system's two electrodes:

• At the cathode, positive cations gain electrons to form neutral atoms. They are reduced.
• At the anode, negative anions lose electrons to form neutral atoms. They are oxidised.

For example, electrolysis of sodium chloride produces chlorine gas at the anode and hydrogen gas at the cathode. Chloride ions are oxidised whilst hydrogen ions are reduced.

2Cl^- → Cl₂ + 2e^-

2H⁺ + 2e^- → H₂

Other examples of redox reactions

Further examples of redox reactions in everyday life are rusting, respiration, and combustion.

• In combustion, the fuel is oxidised and oxygen is reduced.
• In rusting, the iron is oxidised while oxygen is reduced.
• In respiration, electron carriers are oxidised while oxygen is reduced.