Redox

Picture the scene. It’s a chilly November night and you’re standing outside in the dark. Beside you, a crackling bonfire greedily works its way through old pallets and branches. A few jagged nails stick out of the half-rotten wood, tinged dull orange-brown with flaking rust. Believe it or not, the burning wood and rusting metal have something in common - they’re both examples of redox reactions.

• This article is about redox reactions in chemistry.
• We will begin by explaining what is meant by the terms redox, oxidation, and reduction. We’ll then look at some common oxidising and reducing agents.
• Next, we’ll take a look at oxidising states and how we determine the oxidation state of a particular species.
• We will then explain how to write redox equations and half equations and go through some worked examples.
• Lastly, we will touch on disproportionation reactions.

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 are characterised by a movement of electrons and a change in oxidation states.

Let’s start by looking 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.

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

The acronym OILRIG. Anna Brewer, 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:

• Gaining oxygen
• A loss of hydrogen
• A loss of electrons

Likewise, reduction can mean:

• A loss of oxygen
• Gaining hydrogen
• Gaining 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

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 should help you remember the actions of oxidising and reducing agents in terms of electrons: RAD OAT.

RAD OAT. Anna Brewer, 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.

Let’s go through an example.

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.

Calculating oxidation states

The sum of all the oxidation states in a neutral compound must add up to zero, and the sum of all the oxidation states in a complex ion must add up to the charge of the ion - we know this from our rules for assigning oxidation states. But how do we work out the oxidation states of the individual elements within the compound or ion? For this, we can apply our knowledge of fixed oxidation states and work out the unknown oxidation states by deduction.

It can help to follow this process:

1. Look at the charge of the ion or compound, if any. This will help you know what you are aiming for.
2. Identify any atoms with fixed oxidation states and write these above the atoms.
3. Deduce the oxidation states of the remaining atoms, making sure the sum of all the oxidation states adds up to the charge of the ion or compound.

Let’s look at some examples.

Have a go at this next example.

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.

But this overall equation makes it tricky to see the individual oxidation and reduction processes. To see them more clearly, we need to look at half equations. We'll explore them next.

Half Equations

Remember that redox reactions involve both oxidation and reduction. One species is oxidised whilst another is reduced. This means that there is movement of electrons. If we look at our overall redox equation, it isn’t always easy to see this movement. Instead, we can split our redox equation into two half equations that show the oxidation and reduction reactions separately. One half equation represents the process of reduction, while the other half equation represents the process of oxidation.

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.

Let’s look at a couple of examples.

Your two half equations are complete.

Now try a trickier example.

Well done! We made it - our two half equations are complete. But now we need to combine them into one overall equation.

Combining half equations

What if we want to go from half equations to a redox equation? You’ll notice that in redox equations, we don’t see any electrons. In order to form redox equations, we combine multiples of the two half equations so the electrons cancel out.

In our first example, we formed the following two half equations:

Br₂ + 2e^- → 2Br^-

2I^- → I₂ + 2e^-

Both sides of the equation have just two electrons. We can therefore simply add the two equations together to make one overall redox equation:

Br₂ + 2I^- + 2e^- → I₂ + 2Br^- + 2e^-

Because two electrons feature on both sides of the equation, they cancel out:

Br₂ + 2I^- + 2e^- → I₂ + 2Br^- + 2e^-

Here's the overall redox equation:

Br₂ + 2I^- → I₂ + 2Br^-

That example was fairly simple, but our second example from earlier is a little more tricky. We can see that there is only one electron in the first half equation, but five in the second half equation:

Fe²⁺ → Fe³⁺ + e^-

MnO₄^- + 8H⁺ + 5e^- → Mn²⁺ + 4H₂O

We need to multiply the first equation by 5 so that both equations feature the same number of electrons:

5Fe²⁺ → 5Fe³⁺ + 5e^-

We can now add the two equations together. The electrons on each side of the equation cancel out:

MnO₄^- + 8H⁺ + 5e^- + 5Fe²⁺ → Mn²⁺ + 4H₂O + 5Fe³⁺ + 5e^-

MnO₄^- + 8H⁺ + 5Fe²⁺ → Mn²⁺ + 4H₂O + 5Fe³⁺

This is our final answer.

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 can occur.

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

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 look at some more.

Redox Reactions Examples

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

Electrolysis

Either oxidation or reduction takes place at each electrode in electrolysis.

• 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.