Chemistry is the branch of science where interactions and reactions between elements, atoms, electrons, molecules, and many other particles, is mainly considered and studied.
One of the most famous reactions widely abundant and occurring in chemistry, and in our day-to-day life, is the topic of this article: the process of oxidation and reduction and their reactions.
Here, we will discuss the nature of these reactions, their definitions, the difference between oxidation potential and reduction potential, and the difference between oxidizing agents and reducing agents, in addition to giving oxidation vs. reduction examples to clearly comprehend the process… Let’s start!
Table of Contents
Oxidation and Its Historical Definition
In case you are a fan of oxygen-based cleansers, or you are thankful for the sterilizing power of hydrogen peroxide products, then you must thank “oxidation.” On the other hand, you can blame oxidation, if you ever have dealt with a rusty car or a toss out browned fruit.
Oxidation is a process that can start spontaneously or artificially; it is helpful at times and very destructive at others.
The very classical example of oxidation is when iron interacts and combines with oxygen to form the iron oxide, or the famous reddish or orangish component: rust.
In this example, iron is said to be oxidized into rust, where the chemical reaction is given by:
2 Fe + O2 → Fe2O3
The older definition of the oxidation process was defined when oxygen or any other electronegative element was added to a compound. This very basic understanding was originating because oxygen (O2) was the first oxidizing agent to be known.
Later on, oxidation’s definition was widened to include other types of chemical reactions that do not necessarily contain oxygen. Nevertheless, oxygen’s addition to a compound meets the criteria of electron loss perfectly, in parallel with an increase in the oxidation state.
Nowadays, the term oxidation is not looked at to be only oxygen-related, but it is looked at as a whole process that does not have to include oxygen. Rather, oxidation is defined as:
“the loss of electrons during a reaction by a molecule, atom, or ion.”
Not all elements have the same tendency to lose or gain electrons. Some elements, i.e., metals, including magnesium, iron, and sodium, are easily oxidized, whereas the nonmetals like, chlorine, nitrogen, and oxygen, are not easily oxidized, and they are more reluctant to lose their electrons.
Oxidation reactions may involve many forms, as follows:
- Addition of oxygen
C + O2 → CO2 (oxidation of carbon)
- Addition of an electronegative element
Fe + S → FeS (oxidation of iron)
- Removal of an electropositive element
2 KI + H2O2 → I2 + 2 KOH (oxidation of iodide)
- Removal of hydrogen
H2S + Br2 → 2 HBr + S (oxidation of sulfide)
Another important definition is the oxidizing agent, which is:
“a substance that accepts, gains, or receives an electron from a reducing agent.”
Or, more briefly,
“Any substance that oxidizes another substance.”
In the four previously mentioned examples, the oxidizing agents are: O2, S, Br2, H2O2.
Oxidation occurs in stages, and it results in a change in the properties of the atom or the compound that is being oxidized. For example, when iron experiences oxidation, it is transformed into a brittle, reddish powder, and it loses its stiffness and structurally-sound metal nature. This happens as a consequence of the process of losing electrons.
To visually imagine what happens clearly, this following diagram illustrates what happens to an iron atom as it is oxidized. It starts to carry a charge once it is oxidized. This charge is:
- Not only one, but 3 charges, and
- Positive, as it lost 3 electrons.
This is chemically denoted by (3+) sign written as a superscript to the right of the iron (Fe) symbol. Iron is an effortlessly oxidized element; that is, iron’s exposure to oxygen and moisture is important and is always preferred. As long as oxygen is abundant, iron will keep losing its electrons.
Examples of Oxidation Reactions
Many reactions are considered as great examples of oxidation reactions, such as:
- The reaction between hydrogen and fluorine to give out hydrofluoric acid, that is given by:
H2 + F2 → 2 HFThis reaction can be better understood if it is written in terms of two half reactions, like:
H2 → 2 H++ 2 e–
F2 + 2e– → 2 F–
It is easily noticed that this reaction does not have any oxygen atoms in any of its parts.
- The interaction between copper and silver is a great example of electrochemical reactions, where:
Cu (s) + 2 Ag+ (aq) → Cu2+ (aq) + 2 Ag (s)Here, a wire of copper is placed into a silver ions solution, where electrons transfer from the copper metal to the silver ions. As a consequence, copper is oxidized by releasing its ions into the solution, and silver whiskers grow onto the copper wire.
- The reaction between magnesium and oxygen to give out magnesium oxide is one example of oxidation where oxygen is conspicuous in the equation. That is to say:
2 Mg (s) + O2 (g) → 2 MgO (s)
The counter process to that of oxidation is known as reduction, and rationally speaking, it is the process of gaining electrons. Reduction is defined as:
“The gain of electrons during a reaction by a molecule, atom, or ion.”
The historical perspective of reduction was viewed as if it is a process where hydrogen, or any electropositive element, is added.
Reduction mechanism is not any different than that of the oxidation. Here, and instead of losing electrons as it is the case with oxidation, elements gain electrons, and the compound is said to be reduced.
Also, changes happen to the properties of the atom or the compound that is being reduced.
Examples of Reduction Reactions
This may involve many forms, as follows:
- Addition of hydrogen
N2 + 3 H2 → 2 NH3 (reduction of nitrogen)
- Addition of electropositive element
SnCl2 + 2 HgCl2 → SnCl4 + HgCl2 (reduction of mercuric oxygen)
- Removal of oxygen
ZnO + C → Zn + CO (reduction of zinc oxide)
- Removal of electronegative element
2 FeCl3 + H2 → 2 FeCl2 + 2 HCl (reduction of ferric chloride)
Again, the reduction agent is important to be defined as follows:
“a substance that gives, loses, or donates an electron from an oxidizing agent.”
Or, more briefly,
“Any substance that oxidizes another substance.”
The reduction agents in the previous examples are: H2, HgCl2, and C.
Oxidation and reduction are two simultaneous processes that occur together, and they result in a very famous reaction, known as: redox reaction. In which, the atoms or the compounds gain and lose electrons simultaneously at the same time.