Imagine this: it's a chilly November night, and you're standing outside in the dark, next to a crackling bonfire devouring old pallets and branches. The wood and rusty nails sticking out of it might not seem like they have anything in common, but they actually do - they're both part of redox reactions in chemistry.

In this article, we'll explain what redox reactions are, along with oxidation and reduction. We'll also introduce you to some common oxidising and reducing agents, and the concept of oxidising states. Then, we'll show you how to write redox equations and half equations. Finally, we'll touch on disproportionation reactions and provide some examples of redox reactions. If you're interested in chemistry and want to learn more about redox reactions, keep reading. You'll discover how seemingly unrelated things can be connected through the fascinating world of chemistry. And if you're looking for more information on redox reactions, type in "redox" in your search engine to find more articles and resources.

Redox Reaction Definition

Redox reactions involve both oxidation and reduction. Oxidation is the process of losing electrons, while reduction is the process of gaining electrons. In other words, oxidation involves an increase in oxidation state, while reduction involves a decrease in oxidation state.

For example, when iron rusts, it undergoes oxidation because it loses electrons and increases its oxidation state. On the other hand, when chlorine gas reacts with sodium metal to form sodium chloride, chlorine undergoes reduction because it gains electrons and decreases its oxidation state. Understanding oxidation and reduction is crucial in understanding redox reactions. In the next section, we'll explore some common oxidising and reducing agents that are involved in these reactions.

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.

The acronym OILRIG
The acronym OILRIG

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 substances that can oxidise other atoms, ions, or compounds by taking electrons from them. As a result, they themselves are reduced in the process. These agents are also known as oxidants, and some examples of strong oxidising agents include fluorine and oxygen.

The strength of an oxidising agent is influenced by several factors, including electronegativity, electron affinity enthalpy, and oxidation state. This is because oxidising agents require electrons, so anything that increases the attraction between an atom or ion and an incoming electron will increase the agent's oxidising strength. For instance, fluorine is the most electronegative element in the periodic table, making it a potent oxidising agent. Species with a high oxidation state also tend to be good oxidising agents. In the next section, we'll explore reducing agents, which are the opposite of oxidising agents and are involved in the reduction half of redox reactions.

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.


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

Oxidation states, also known as oxidation numbers, are crucial in determining which species is oxidised and which is reduced in a redox reaction. These numbers represent the number of electrons that have been gained or lost by an ion, compared to the element in its uncombined state. A positive oxidation state indicates that the element has lost electrons, while a negative oxidation state indicates that the element has gained electrons.

Changes in oxidation states can be used to determine whether a species has been oxidised or reduced. If a species goes from a positive oxidation state to a more negative oxidation state, it means that the species has gained electrons and has been reduced. Conversely, if a species goes from a negative oxidation state to a more positive oxidation state, it means that the species has lost electrons and has been oxidised. By using oxidation states, we can identify the oxidising and reducing agents in a redox reaction and determine the overall direction of electron flow. This allows us to balance the equation and predict the outcome of the reaction.

Assigning oxidation states

Assigning oxidation states to elements in compounds and ions is a crucial step in understanding redox reactions. There are several rules to follow when assigning oxidation states.

Firstly, all uncombined elements have an oxidation state of 0. Secondly, the oxidation states of all atoms or ions in a neutral compound add up to 0. Thirdly, the sum of oxidation states in an ion equals the charge of the ion. Additionally, in an ion or compound, the more electronegative atom generally has the negative oxidation state. Finally, certain elements have specific oxidation states such as group 1 metals having an oxidation state of +1 and oxygen usually having an oxidation state of -2.

Using these rules, we can determine the oxidation state of an unknown element in a compound or ion. For example, in copper oxide (CuO), copper ions have a charge of 2+. To form a copper ion with a charge of 2+, each copper atom has to lose two electrons, resulting in a copper ion with an oxidation state of +2.

By understanding oxidation states, we can accurately balance redox reactions and determine the direction of electron flow. It is a fundamental concept in chemistry that is essential for understanding a wide range of chemical reactions.

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

When writing a redox equation, we first identify the reactants and products and then assign oxidation states to each element. We then balance the equation by making sure the number of atoms on both sides of the equation is equal and that the overall charge is balanced.

Spectator ions are not involved in the redox reaction and do not change their oxidation state. Therefore, they can be left out of the redox equation as they do not affect the overall reaction. For example, in the reaction between zinc and hydrochloric acid to produce zinc chloride and hydrogen gas:

Zn + 2HCl → ZnCl2 + H2

The hydrogen and chlorine ions are spectator ions as they are present on both sides of the equation and do not undergo any oxidation or reduction. Therefore, we can eliminate them from the equation and simplify it to:

Zn → Zn2+ + 2e- (oxidation)

2H+ + 2e- → H2 (reduction)

Overall equation: Zn + 2H+ → Zn2+ + H2

This equation shows the oxidation of zinc and the reduction of hydrogen ions to form hydrogen gas. It is a simplified representation of the redox reaction that occurs.

Redox equations are important in understanding the basics of electrochemistry and chemical reactions. They provide a simple way to represent complex chemical processes and allow us to predict the outcome of a reaction. By identifying the species being oxidised and reduced, we can determine the direction of electron flow and the overall outcome of the reaction.

Writing redox equations

Now, let's break down the redox equation into half-reactions to see the oxidation and reduction processes that occur:

Oxidation half-reaction: Mg(s) → Mg2+(aq) + 2e- (Mg loses 2 electrons and is oxidised)

Reduction half-reaction: Cu2+(aq) + 2e- → Cu(s) (Cu gains 2 electrons and is reduced)

Overall redox equation: Mg(s) + Cu2+(aq) → Mg2+(aq) + Cu(s)

This redox equation shows that magnesium is oxidised to form magnesium ions, while copper ions are reduced to form solid copper. The electrons lost by magnesium are gained by copper ions, resulting in the formation of a solid copper.

Redox equations are essential in understanding chemical reactions that involve electron transfer. By breaking down the equation into half-reactions, we can identify the species being oxidised and reduced, and determine the flow of electrons during the reaction.

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

As an AI language model, I don't have access to websites or external resources. However, I can provide you with an example of how to write a half equation.

Let's consider the reaction between potassium permanganate and iron(II) ions in acidic solution:

MnO4-(aq) + Fe2+(aq) → Mn2+(aq) + Fe3+(aq)

To write a half equation for the oxidation of Fe2+ to Fe3+, we follow these steps:

Fe2+(aq) → Fe3+(aq)

Balance the elements: Fe2+(aq) → Fe3+(aq)

Add water molecules to balance the oxygens: Fe2+(aq) → Fe3+(aq) + H2O(l)

Add hydrogen ions to balance the hydrogens: Fe2+(aq) → Fe3+(aq) + H2O(l) + 2H+(aq)

Add electrons to balance the charges: Fe2+(aq) → Fe3+(aq) + e-

So the half equation for the oxidation of Fe2+ to Fe3+ is:

Fe2+(aq) → Fe3+(aq) + e-

This half equation shows that Fe2+ loses an electron to form Fe3+. By following the same steps, we can also write a half equation for the reduction of MnO4- to Mn2+, which involves the gain of 5 electrons:

MnO4-(aq) + 8H+(aq) + 5e- → Mn2+(aq) + 4H2O(l)

Overall, the redox equation for the reaction between potassium permanganate and iron(II) ions in acidic solution is:

MnO4-(aq) + 8H+(aq) + 5Fe2+(aq) → Mn2+(aq) + 5Fe3+(aq) + 4H2O(l)

Disproportionation Reactions in Redox

Electrolysis is a redox reaction that occurs when an electric current is passed through an electrolyte. An electrolyte is a solution containing ions that can be broken down by the electric current.

During electrolysis, the positively charged ions move towards the negative electrode (the cathode), while the negatively charged ions move towards the positive electrode (the anode). At the electrodes, the ions may undergo oxidation or reduction reactions, depending on their oxidation states.

For example, consider the electrolysis of aqueous copper(II) sulfate using copper electrodes. The overall equation is:

CuSO4(aq) + H2O(l) → Cu(s) + O2(g) + H2SO4(aq)

At the cathode (negative electrode), copper ions are reduced to form solid copper:

Cu2+(aq) + 2e- → Cu(s)

At the anode (positive electrode), water molecules are oxidised to form oxygen gas and hydrogen ions:

2H2O(l) → O2(g) + 4H+(aq) + 4e-

The hydrogen ions then react with sulfate ions in the solution to form sulfuric acid:

2H+(aq) + SO42-(aq) → H2SO4(aq)

Overall, the electrolysis of copper(II) sulfate produces solid copper, oxygen gas, and sulfuric acid.

Another common example of a redox reaction is the combustion of hydrocarbons, such as methane (CH4), in the presence of oxygen gas (O2). During combustion, the hydrocarbon is oxidised to form carbon dioxide (CO2) and water (H2O), while the oxygen is reduced to form water:

CH4(g) + 2O2(g) → CO2(g) + 2H2O(l)

In this reaction, methane is oxidised to carbon dioxide, and oxygen is reduced to water.


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

Overall, redox reactions are fundamental to many chemical processes, both in the laboratory and in everyday life. Understanding how to identify oxidation and reduction, assign oxidation states, write redox equations, and balance half equations is important for anyone studying chemistry. With practice and perseverance, anyone can become proficient in this essential aspect of chemistry.

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.


What is the definition of a redox reaction?

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

What is an example of a redox reaction?

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.

How do you identify redox reactions?

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

Why is electrolysis called a redox reaction?

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.

What happens in redox reactions?

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

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