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Redox

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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 introduce you to oxidising states.
  • We will then show you how to write redox equations and look at half equations.
  • Lastly, we will touch on disproportionation reactions and examples of redox reactions.

Redox Reaction Definition

Redox is a term used to describe reactions involving both oxidation and reduction. These reactions involve a movement of electrons, and are characterised by a change in oxidation states.

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.

Oxidation is the gain of oxygen.

You can think of reduction as the opposite of oxidation.

Reduction is the loss of oxygen.

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.

Oxidation is the loss of hydrogen, and reduction is the gain of hydrogen.

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.

Oxidation is the loss of electrons, and reduction is the gain of electrons.

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

Redox, OILRIG, StudySmarterThe 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 O2- 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 are species that oxidise another atom, ion, or compound. They are reduced in the process.

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.

There are several factors that affect the strength of an oxidising agent. These include electronegativity, electron affinity enthalpy, and oxidation state, which we’ll look at in just a second. This is because oxidising agents take electrons, so anything that increases the attraction between an atom or ion and an incoming electron will increase the agent’s oxidising strength. For example, fluorine is the most electronegative element in the periodic table, which makes it a very strong oxidising agent. Species with a high oxidation state also tend to be good oxidising agents.

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 are species that reduce another atom, ion, or compound. They are oxidised in the process.

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.

Redox, RADOAT, StudySmarterRAD 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.

Oxidation states are numbers assigned to ions that show how many electrons the ion has lost or gained, compared to the element in its uncombined state. A positive oxidation state shows that the element lost electrons, whilst a negative oxidation state shows that it gained electrons. They can also be referred to as oxidation numbers.

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:

What is the oxidation state of copper in copper oxide, CuO?

To solve this problem, look at the copper ions in copper oxide. They are ions with a charge of 2+. How do you turn an uncombined atom, which is just a neutral atom on its own, into an ion with a charge of 2+? By losing two electrons! To form a copper ion with a charge of 2+, each copper atom has to lose two electrons. This means that these copper ions have an oxidation state of +2.

When talking about charges, we put the number first: 2+. But when talking about oxidation states, we put the positive or negative symbol first: +2.

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

A redox equation is a way of representing a redox reaction.

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.

Spectator ions are ions that are present in both the reactants and products of a reaction. They are totally unchanged by the reaction - their physical state, oxidation state, and charge don’t change.

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.

Write a redox equation for the reaction between magnesium and copper sulfate. The overall equation is given below:

Mg(s) + CuSO4(aq) → MgSO4(aq) + Cu(s)

First of all, we need to find out the spectator ions in the equation. These don't change oxidation states. They aren't oxidised or reduced, so we don't need to worry about them.

To help you identify the spectator ions, split the ionic salts into their constituent ions:

Mg(s) + Cu2+(aq) + SO42-(aq) → Mg2+(aq) + SO42-(aq) + Cu(s)

The sulfate ion, SO42-, is present on both sides of the equation. It doesn’t change physical state, oxidation state, or charge. This means that it is a spectator ion. To write an overall redox equation, we simply omit this ion. Here's your final answer:

Mg(s) + Cu2+(aq) → Mg2+(aq) + Cu(s)

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.

Half equations are equations that show one half of a redox reaction. One half equation shows the oxidation process, whilst the other shows the reduction process.

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.

The only three things you can add into half equations, besides more of the reactant or product, are water, hydrogen ions, and electrons. You can't sneak in oxygen gas, for example.

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.

Disproportionation reactions are reactions where the same species is both oxidised and reduced.

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 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.

Redox - Key takeaways

  • Redox is a term used to describe reactions involving both oxidation and reduction. These reactions involve a movement of electrons, and are characterised by a change in oxidation states.
  • Oxidation and reduction have several definitions in chemistry. However, oxidation is generally taken to mean a loss of electrons, whilst reduction means a gain of electrons.
  • Reducing agents are electron donors that reduce another species and are oxidised themselves.
  • Oxidising agents are electron acceptors that oxidise another species and are reduced themselves.
  • Oxidation states are numbers assigned to ions that show how many electrons they have lost or gained, compared to the element in its uncombined state. A positive oxidation state shows that the element has lost electrons, whilst a negative oxidation state shows that it has gained electrons.
  • There are certain rules to follow when assigning oxidation states.
  • A redox equation is an equation showing the reduction and oxidation processes in a reaction, ignoring any spectator ions.
  • Redox equations can be split into half equations which show the oxidation and reduction processes separately. They also show the movement of electrons.
  • Disproportionation reactions are reactions in which the same species is both reduced and oxidised.

Frequently Asked Questions about Redox

A redox reaction is a reaction involving both oxidation and reduction.

A simple redox reaction is a displacement reaction between two metals, such as adding magnesium to a solution of Fe2+ ions. The magnesium is oxidised and loses electrons, whilst the iron is reduced and gains these electrons. Examples of more complicated redox reactions include respiration, combustion, and rusting.

We can identify redox reactions because they cause the oxidation state(s) of one or more species to change.

In electrolysis, reduction and oxidation both happen simultaneously. This makes it an example of a redox reaction. At the cathode, positive cations gain electrons and so are reduced, whilst at the anode, negative anions lose electrons and so are oxidised.

In redox reactions, one species is oxidised and another species is reduced. The oxidised species loses electrons, while the reduced species gains these electrons.

Final Redox Quiz

Question

What is a redox reaction?

Show answer

Answer

A reaction involving both oxidation and reduction.

Show question

Question

Which of the following are definitions of oxidation?

Show answer

Answer

Gain of oxygen

Show question

Question

Which of the following are definitions of reduction?


Show answer

Answer

Gain of hydrogen

Show question

Question

Which of the following are redox reactions?


Show answer

Answer

Mg + Cu2+ → Mg2+ + Cu

Show question

Question

What does the acronym OIL RIG stand for?


Show answer

Answer

Oxidation is loss, reduction is gain.

Show question

Question

What does the acronym RAD OAT stand for?


Show answer

Answer

 Reducing agents donate, oxidising agents take.

Show question

Question

What are oxidation states?


Show answer

Answer

Numbers assigned to ions that show how many electrons they have lost or gained, compared to the element in its uncombined state.

Show question

Question

What is the oxidation state of iron in iron(II) chloride?


Show answer

Answer

+2

Show question

Question

What is the oxidation state of oxygen in O2?


Show answer

Answer

0

Show question

Question

What is the oxidation state of S in SO42-?


Show answer

Answer

+6

Show question

Question

What is the oxidation state of S in H2S?


Show answer

Answer

-2

Show question

Question

What is a disproportionation reaction?


Show answer

Answer

A reaction where the same species is both reduced and oxidised.

Show question

Question

True or false? Group 1 metals always have an oxidation state of +1.


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Answer

True

Show question

Question

True or false? Chlorine always has an oxidation state of -1.


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Answer

False

Show question

Question

What is the oxidation state of Mg2+?


Show answer

Answer

+2

Show question

Question

Write half equations for the following redox reaction:

    

Cl2 + 2Br- → 2Cl- + Br2

Show answer

Answer

Cl2 + 2e- → 2Cl-

2Br- → Br2 + 2e-

Show question

Question

What three species can you add when balancing half equations?


Show answer

Answer

Oxygen

Show question

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