In this chapter of Redox Reaction, you will be learning what precisely the term ‘redox’ means. This topic is short, but a significant one because it is the foundation of more challenging topics later on. A clear understanding of this topic is needed to understand reaction mechanisms for future topics in O and A Levels.
What is redox reaction?
The term ‘redox’ is used to describe any type of chemical reactions. We can characterise any chemical reactions into two main groups: redox reaction and non-redox reaction.
Redox reaction is a chemical reaction where reduction and oxidation have occurred.
On the other hand, a non-redox reaction is where neither reduction or oxidation have occurred.
Important note: If a chemical reaction is a redox reaction, both oxidation and reduction must co-occur. In other words, whenever reduction occurs, oxidation must also occur in a same chemical reaction.
When both reduction and oxidation happens, we describe the overall reaction as a redox reaction.
Definition of reduction and oxidation
Reduction and oxidation are key terms in this chapter. So what exactly is the meaning of reduction and oxidation?
Let us first start with oxidation. Oxidation happens when a species either
(1) Gains oxygen
(2) Losses hydrogen
(3) Losses electrons
(4) Experience an increase in oxidation state.
In the previous statement, ‘species’ can mean any chemical ions, atoms or compounds. Anyone of the four requirements would indicate an oxidation reaction.
(1) Oxidation- Gain Oxygen
When a species gains oxygen, it is said to be oxidised. For example, when zinc reacts with atmospheric oxygen to form zinc oxide, zinc is oxidised to form zinc(II) ions, and we say that oxidation occurs.
E.g. 1: Fe2O3(s) + 3CO(g) –> 2Fe(s) + 3CO2(g)
Carbon gained oxygen and is oxidised.
E.g. 2: Fe2O3(s) + 2Al(s) –> 2Fe(s) + Al2O3(s)
Aluminium gained oxygen and is oxidised.
(2) Oxidation- Loss of Hydrogen
Oxidation is also said to happen when a species loses hydrogen. For instance, when hydrogen sulfide loses hydrogen to form sulfur, this reaction is called oxidation.
E.g. 3: H2S(g) + Cl2(g) –> S(s) + 2HCl(g)
Sulfur lost hydrogen and is oxidised.
E.g. 4: C2H4(g) + H2(g) –> C2H6(g)
H–H bond in H2 breaks. Hydrogen loses hydrogen and is therefore oxidised.
(3) Oxidation- Lose Electrons (LEO: lose electron oxidation)
Oxidation is also said to happen when a species loses electrons. Below is an example of oxidation through the loss of electrons.
When an iron atom loses two of its valence electrons (electrons from the outermost shells) to form positively charged iron(II) ions, oxidation has happened. We say that iron is oxidised to form iron(II) ions as it loses two electrons.
Remember a the start I mentioned when oxidation happen reduction must also happen simultaneously? From the example above, you may ask: Did any species undergo reduction by gaining electrons?
The answer is yes. Oxidation and reduction always happen simultaneously. In this reaction, reduction reaction does not seem to happen because the overall reaction is incomplete.
If you notice, I mentioned that iron atoms are oxidised because they lose electrons. Another species gain the electrons that were lost by iron, probably a non-metal. (Note: non-metal gain electrons to form a stable noble gas configuration.)
Let me extend this reaction by saying iron reacts with chlorine gas.
The complete reaction is as follows: iron atom loses two electrons to chlorine and chlorine gains the electrons. Iron forms positively charged iron(II) ions while chlorine forms negatively charged chloride ions (because they gained the electrons from iron). The above reaction is the complete redox reaction, where reduction of chlorine and oxidation of iron occurs.
(4) Oxidation- Increase in Oxidation state
There is another term that is very important in this chapter. That is the oxidation state.
Oxidation state is the hypothetical charge that an atom if all bonds to atoms of different elements were completely ionic, with no covalent character.
This definition is somewhat tricky to understand. I will further elaborate on this definition and using an example to illustrate.
Suppose you have a hydrogen chloride (HCl) molecule. From the chapter of bonding, you learn that the bond between H and Cl is a covalent bond because both are non-metal ions.
However, to find the oxidation of an atom, we will treat the bond as ionic. In that case, hydrogen will act as the ‘positively charged metal’ atom while chlorine will act as the ‘negatively charged non-metal’ atom. Once you have established this, you will realise that hydrogen atom loses an electron to the chlorine atom to form positively charged hydrogen ions of (1+). This is opposite for the chlorine atom; chlorine atom will acquire a charge of (1-) since it has gained an electron. Hypothetically, you will get H+ ions and Cl– ions in its structure.
Now, we can determine the oxidation state of atoms in HCl. Since H ‘has a charge of 1+’, the oxidation state is +1. Likewise, since Cl ‘has a charge of 1-‘, the oxidation state is -1.
Note: when we write charge, the sign is at the back of the magnitude (like ‘1+’ and not ‘+1’). However, writing oxidation state is the opposite. For oxidation states, the sign comes before the magnitude.
This is how we determine the oxidation state. The purpose of assigning oxidation states is to understand redox reaction later.
The oxidation state is a number given to an element. Therefore, you can only say the oxidation state of iron and not the oxidation state of iron(II).
Note: transition metals have variable oxidation states. Thus, when you are writing chemical symbols for the transition metal, do remember to write the oxidation states. For example, for CuCl2, write as copper(II) chloride and not copper chloride. The ‘(II)’ is to indicate the oxidation state of copper, since there is also copper(I) available.
Known oxidation states
- The oxidation state of pure elements is 0.
- The oxidation state of a simple ion is the charge of the ion. This means that the oxidation state of Group 1 elements in compounds is +1, Group 2 is +2…. This is because the charge of the ion formed is the same as the numerically same as the group number.
- The oxidation state of oxygen in any compounds is always +2 unless it is combined with fluorine.
- Oxidation of chlorine is always -1 unless it is combined with oxygen or fluorine (rare cases).
- Hydrogen is always +1 unless it is combined with reactive metals.
- The sum of the oxidation state of the elements in a compound is equal to the charge carried by the compound.
1. The oxidation state of oxygen in hydrogen peroxide is -1.
Note: there is no such thing as reduction state!
Oxidation state is the primary method to prove oxidation
All the reactions that indicate oxidation (gain in oxygen, loss in hydrogen and loss in electron) have something in common.
Whenever an element undergoes such reaction, the oxidation state always increase.
Take a look at the previous example, where iron atom loses electrons to form iron(II) ions. Before losing electrons, the oxidation state of iron is 0 (rule 1: the oxidation state of pure elements is 0). After the reaction has occurred, the oxidation state of iron is +2. Oxidation has , and there is an increase in oxidation state.
Oxidation always increases in oxidation state no matter what happens (lose an electron, gain oxygen or lose hydrogen).
Reduction is the direct opposite of oxidation. Reduction happens when there is a decrease in the oxidation state of the elements. This is due to either the following:
(1) Lose of oxygen
(2) Gain of electron (GER: Gain Electron Reduction)
(3) Gain hydrogen
(4) Decrease in oxidation state
Remember, looking at the oxidation state is the best way of determining if a reaction is oxidation or reduction reaction!
How to determine which reaction is redox?
Redox reactions happen when reduction and oxidation happen simultaneously. Hence, to see if redox occurs, we only need to check if either reduction or oxidation happens (because when one happens, the other always happens!).
To check if oxidation/reduction happens, we can check if there is any change of oxidation state of the same element.
Note: acid-base reaction is not a redox reaction, because there is no change of oxidation state.
Disproportionation reaction is a redox reaction where an element undergoes both oxidation and reduction simultaneously to form two different products in a same chemical reaction.
Oxidising and reducing agents
To understand the terms ‘oxidising agent’ and ‘reducing agent’, it is essential to understand what happens during oxidation or reduction.
Take, for example, the reaction between A and B. A is oxidised with B is reduced. Redox reaction occurs.
A is the reducing agent as A caused the reduction of B. Similarly, B is the oxidising agent as B caused the oxidation of A.
Therefore, the reducing agent is oxidised while an oxidising agent is reduced.
Test for reducing agent
To test for reducing/oxidising agent, you need something that can be oxidised. Mainly, you need something that can change colour after it can be oxidised so that you will know a redox reaction has occurred.
To test for reducing agent, you can use acidified aqueous potassium manganate(VII) solution ( KMnO4 (aq)). This solution is purple. The name of this solution is long, but it merely means adding acid (sulfuric acid) to a solution of potassium manganate(VII). The notation (VII) indicates that the oxidation state of manganese is +7.
When KMnO4 (aq) is added to a reducing agent, the reducing agent will reduce manganese in KMnO4 (aq) and the oxidation state of manganese decreases from +7 in KMnO4 (aq) to +2 in Mn2+.
When this process happens, Mn2+ will be formed which is colourless. Hence, when a redox reaction occurs involving KMnO4 (aq), the colour changes from purple to colourless.
MnO4–(aq) [purple] + 8H+(aq) + 5e– –> Mn2+(aq) [colourless] + 4H2O(l)
This method can be used to test for a reducing agent (a substance that is oxidised) when it is reacted with KMnO4 (aq) because a visible colour change is observed.
Note: hydrochloric acid cannot be used to acidify the KMnO4 (aq) solution because the chloride ions will be oxidised to form chlorine gas.
Another reagent acidified aqueous potassium dichromate(VI) can be used to identify the reducing agent. The oxidation state of chromium decreases from +6 in Cr2O72– to +3 in Cr3+. There is a colour change of orange Cr2O72– to green Cr3+.
Cr2O72–(aq) + 14H+(aq) + 6e– –> 2Cr3+(aq) + 7H2O(l)
Test for oxidising agent
Aqueous potassium iodide is used to test for oxidising agent. The iodide ions from the aqueous potassium iodide will be oxidised by the oxidising agent, and when this happens, the solution turns from colourless to brown.
2I–(aq) –> I2(aq) + 2e–