When studying chemical reactions, all the substances involved and the processes that occur in them must be taken into account. Among the most important are the oxidation-reduction reactions, also called redox, which involve the transfer or transfer of electrons between two or more chemical species.
Two substances interact in these reactions: the reducing agent and the oxidizing agent. Some of the oxidizing agents that can be observed more frequently are oxygen, hydrogen, ozone, potassium nitrate, sodium perborate, peroxides, halogens and permanganate compounds, among others.
What are oxidizing agents?
In the oxidation half-reaction, the oxidizing agent is reduced because, upon receiving electrons from the reducing agent, a decrease is induced in the value of the charge or oxidation number of one of the atoms of the oxidizing agent.
This can be explained by the following equation:
2Mg (s) + O 2 (g) → 2MgO (s)
It can be seen that magnesium (Mg) reacts with oxygen (O2), and that oxygen is the oxidizing agent because it removes electrons from magnesium —that is, it is being reduced— and magnesium, in turn, becomes in the reducing agent of this reaction.
Similarly, the reaction between a strong oxidizing agent and a strong reducing agent can be very dangerous because they can interact violently, so they must be stored in separate locations.
What factors define the strength of an oxidizing agent?
These species are distinguished according to their “strength”. That is, the weakest are those that have a lower capacity to subtract electrons from other substances.
It is known as the half of the distance that separates the nuclei of two atoms of neighboring or “neighboring” metallic elements.
Atomic radii are generally determined by the force with which the most superficial electrons are attracted to the nucleus of the atom.
Therefore, the atomic radius of an element decreases in the periodic table from bottom to top and from left to right. This implies that, for example, lithium has a significantly larger atomic radius than fluorine.
Electronegativity is defined as the ability of an atom to capture electrons belonging to a chemical bond towards itself . As electronegativity increases, elements show an increasing tendency to attract electrons.
Generally speaking, electronegativity increases from left to right in the periodic table and decreases as the metallic character increases, with fluorine being the most electronegative element.
It is said that it is the variation of the energy that is registered when an atom receives an electron to generate an anion; that is, it is the ability of a substance to receive one or more electrons.
As electron affinity increases, the oxidative capacity of a chemical species increases.
It is the minimum amount of energy that is needed to remove an electron from an atom or, in other words, it is a measure of the “force” with which an electron is bound to an atom.
The greater the value of this energy, the more difficult it is to detach an electron. Thus, the ionization energy enlarges from left to right and decreases from top to bottom in the periodic table. In this case, the noble gases have large values of ionization energies.
The strongest oxidizing agents
Taking into account these parameters of the chemical elements, it is possible to determine which are the characteristics that the best oxidizing agents should have: high electronegativity, low atomic radius and high ionization energy.
That said, the best oxidizing agents are considered to be the elemental forms of the most electronegative atoms, and it is noted that the weakest oxidizing agent is metallic sodium (Na +) and the strongest is the elemental fluorine molecule (F2), which is capable of oxidizing a large number of substances.
Examples of reactions with oxidizing agents
In some oxide-reduction reactions it is easier to visualize electron transfer than in others. Some of the most representative examples will be explained below:
The decomposition reaction of mercury oxide:
2HgO (s) → 2Hg (l) + O 2 (g)
In this reaction, mercury (oxidizing agent) is distinguished as the receptor for oxygen electrons (reducing agent), decomposing into liquid mercury and gaseous oxygen when heated.
Another reaction that exemplifies oxidation is that of sulfur burning in the presence of oxygen to form sulfur dioxide:
S (s) + O 2 (g) → SO 2 (g)
Here it can be seen that the oxygen molecule is oxidized (reducing agent), while elemental sulfur is reduced (oxidizing agent).
Finally, the combustion reaction of propane (used in gas for heating and cooking):
C 3 H 8 (g) + 5O 2 (g) → 3CO 2 (g) + 2H 2 O (l)
In this formula the reduction of oxygen (oxidizing agent) can be observed.