Reduction-oxidation (Redox)

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Reduction -oxidation reaction (also, oxidation-reduction reaction or simply redox reaction) is any chemical reaction in which one or more electrons are transferred between the reactants, causing a change in their oxidation states.

For a reduction-oxidation reaction to exist, there must be an element in the system that gives up electrons, and another that accepts them:

The oxidizing agent is that chemical element that tends to capture those electrons, leaving it with a lower oxidation state than it had, that is, being reduced.
The reducing agent is that chemical element that supplies electrons from its chemical structure to the medium, increasing its oxidation state, that is, being oxidized.

When a reducing chemical element gives up electrons to the medium, it becomes an oxidized element, and its relationship with its precursor is established by what is called a “redox couple”. Similarly, it is said that when a chemical element captures electrons from the medium, it becomes a reduced element, and also forms a redox couple with its oxidized precursor. When a species can be oxidized and reduced at the same time, it is called ampholyte, and the oxidation-reduction process of this species is called ampholation or dismutation.

Principle of electroneutrality

Pauling's principle of electroneutrality corresponds to an approximation method to estimate the charge in molecules or complex ions; this principle assumes that the charge is always distributed at values close to 0 (ie -1, 0, +1).

Within a global redox reaction, there is a series of particular reactions called half-reactions or partial reactions.

Reduction half-reaction: Cu}}}">
Oxidation half-reaction: Fe^2+ + 2 e^-}}}">

or more commonly, also called the general equation:

{ displaystyle { ce {Fe + Cu ^ 2 + -> Fe ^ 2 + + Cu}}}

The tendency to reduce or oxidize other chemical elements is quantified by the reduction potential, also called the redox potential. A redox titration is one in which a chemical indicator indicates the percentage change in the redox reaction by the color change between the oxidant and the reductant.

Oxidation

Oxidation is a chemical reaction where an element loses electrons, thereby increasing its oxidation state.It should be noted that an oxidation or a reduction is actually a process by which the oxidation state of a compound changes. This change does not necessarily mean an exchange of ions. It implies that all the compounds formed by a redox process are ionic, since it is in these compounds where an ionic bond occurs, product of the transfer of electrons. For example, in the reaction to form hydrogen chloride from dihydrogen and dichloro gases, a redox process occurs and yet a covalent compound is formed. These two reactions always occur together; that is, when a substance is oxidized, it is always due to the action of another that is reduced. One gives up electrons and the other accepts. For this reason, the general term redox reactions is preferred.

Life itself is a redox phenomenon. Oxygen is the best oxidant that exists because the molecule is not very reactive (because of its double bond), and yet it is very electronegative, almost as much as fluorine. The most oxidizing substance that exists is the KrF cation.+because it easily forms Kr and F+. Among others, there is potassium permanganate (KMnO4), potassium dichromate (K2Cr2O7), hydrogen peroxide (H2O2), nitric acid (HNO3), hypohalites and halates (for example, sodium hypochlorite (NaClO), which is highly oxidizing in alkaline medium, and potassium bromate (KBrO3)). Ozone (O3) is a very energetic oxidant:

{displaystyle {ce {Br^- + O3 -> BrO3^-}}}

The name "oxidation" comes from the fact that, in most of these reactions, the transfer of electrons occurs through the acquisition of oxygen atoms (transfer of electrons) or vice versa. However, oxidation and reduction can occur without oxygen exchange involved: for example, the oxidation of sodium iodide to iodine by the reduction of chlorine to sodium chloride:

{displaystyle {ce {2 NaI + Cl2 -> I2 + 2 NaCl}}}

This can be broken down into its two corresponding half-reactions:

Reduction half-reaction:

2 Cl^-}}}">

Oxidation half-reaction:

I_2 + 2 e^-}}}">Example

Iron can present two oxidized forms:

Iron (II) oxide: FeO.
Iron (III) oxide: Fe2O3.

Reduction

In chemistry, reduction is the electrochemical process by which an atom or an ion gains electrons. It implies the decrease of its oxidation state. This process is the opposite of oxidation.

When an ion or an atom is reduced it has these characteristics:

It acts as an oxidizing agent.
It is reduced by a reducing agent.
Decreases its state or oxidation number.

Example

The iron(III) ion can be reduced to iron(II):

{displaystyle {ce {Fe^3+ + 1e^- -> Fe^2+}}}

In organic chemistry, decreasing bonds of oxygen atoms to carbon atoms or increasing hydrogen bonds to carbon atoms is interpreted as reduction. For example:

CH_2=CH_2}}}">(Ethyne is reduced to give ethene).
CH_3-CH_2OH}}}">(ethanal is reduced to ethanol).

Oxidation number

The quantification of a chemical element can be done by its oxidation number. During the oxidation process, the oxidation number or also called the oxidation state of the element increases. Instead, during reduction, the oxidation number of the species being reduced decreases. The oxidation number is an integer that represents the number of electrons an atom brings into play when it forms a given bond. In a pure element, all atoms are neutral, since they have no charge and are assigned the oxidation state 0.

The oxidation number:

It increases if the atom loses electrons (the chemical element that is oxidized), or shares them with an atom that has a tendency to capture them.
It decreases when the atom gains electrons (the chemical element that is reduced), or shares them with an atom that has a tendency to give them up.

Rules for assigning the oxidation number

The oxidation number of all uncombined elements is zero. Regardless of how those numbers are represented.
The oxidation number of monatomic ionic species matches the charge of the ion.
The oxidation number of combined hydrogen is +1, except in metal hydrides, where its oxidation number is –1 (ex: AlH3, LiH).
The oxidation number of combined oxygen is –2, except in peroxides, where its oxidation number is –1 (ex: Na2O2, H2O2).
The oxidation number in metallic elements, when combined, is always positive and numerically equal to the charge on the ion.
The oxidation number of halogens in hydracids and their respective salts is –1, whereas the oxidation number of sulfur in their hydracid and respective salts is –2.
The oxidation number of a neutral molecule is zero, so the sum of the oxidation numbers of the atoms that make up a neutral molecule is zero.
The total electrical charge of a non-neutral (non-zero) molecule corresponds to the algebraic sum of the oxidation numbers of all the atomic species that constitute it. (ex: MnO−4 = (1)*(+7) + (4)*(-2) = -1).

Relationship Adjustment

All redox processes require the stoichiometric adjustment of the components of the half-reactions for oxidation and reduction.

For reactions in aqueous medium, generally add:

in acid medium hydrogen ions (H+), water molecules (H2O), and electrons
in basic medium hydroxyls (OH−), water molecules (H2O), and electrons to compensate for changes in oxidation numbers.

Acid medium

In an acid medium, hydroniums (cations) are added (H+) and water (H2O) to the half-reactions to balance the final equation. On the side of the equation that lacks oxygen, water molecules will be added, and on the side of the equation that lacks hydrogen, hydronium will be added. For example, when manganese (II) reacts with sodium bismuthate.

Unbalanced equation: Bi^{3+}(aq) + MnO4^{-}(aq)}}}">Oxidation: MnO4^{-}(aq) + 5 e^-}}}">Reduction: Bi ^ {3 +} (aq)}}}">

Now we have to add the hydroniums and the water molecules where hydrogens are needed and where oxygens are needed, respectively.Oxidation: MnO4^{-}(aq)}}+color {Blue}{ce {8 H^{+}(aq)}}color {Black}+{ce {5 e^-}}}">Reduction: Bi^{3+}(aq)}}+color {Blue}{ce {3 H2O}}color {Black}}">

The reactions will balance when equalizing the number of electrons involved in both half-reactions. This will be achieved by multiplying the reaction in one half-reaction by the number of electrons in the other half-reaction (and, if necessary, vice versa), so that the number of electrons is constant.Oxidation: MnO4^{-}(aq) + 8 H^{+}(aq)}}+color {OliveGreen}{ce {5 e^-}}color {Black})color {Orange}times 2color {Black}}">Reduction:{Bi^{3+}(aq)}+{3H2O}}})color {OliveGreen}times 5color {Black}}">

At the end we will have:Oxidation: 2 MnO4^{-}(aq) + 16 H^{+}(aq) + 10 e^-}}}">Reduction: 5 Bi^{3+}(aq) + 15 H2O}}}">

As you can see, the electrons are balanced, so we proceed to add the two half-reactions, to finally obtain the balanced equation. ${displaystyle {underline {left.{begin{array}{rcl}8H_{2}O+2Mn_{(aq)}^{2+}to 2MnO_{4(aq)}^{-}+16H_{(aq)}^{+}+10e^{-}\10e^{-}+30H^{+}+5BiO_{3(s)}^{-}to 5Bi_{(aq)}^{3+}+15H_{2}Oend{array}}rightDownarrow +}}}$ ${displaystyle 14H_{(aq)}^{+}+2Mn_{(aq)}^{2+}+5NaBiO_{3(s)}to 7H_{2}O+2MnO_{4(aq)}^{ -}+5Bi_{(aq)}^{3+}+5Na_{(aq)}^{+}}$

Basic medium

Hydroxide ions (anions) (OH−) and water (H2O) to the half-reactions to balance the final equation. For example, we have the reaction between Potassium Permanganate and Sodium Sulfite.

Unbalanced equation: ${ displaystyle KMnO_ {4}+Na_ {2} SO_ {3}+H_ {2} O to MnO_ {2}+Na_ {2} SO_ {4}+KOH}$

We separate the half-reactions intoOxidation: ${displaystyle SO_{3}^{2-}to SO_{4}^{2-}+2e^{-}}$ Reduction: ${displaystyle 3e^{-}+MnO_{4}^{-}to MnO_{2}}$

We add the appropriate amount of Hydroxides and Water (the water molecules are located where there is a greater amount of oxygen).Oxidation: ${displaystyle color {Blue}2OH^{-}color {Black}+SO_{3}^{2-}to SO_{4}^{2-}+color {Blue}H_{2}Ocolor {Black}+2e-}$ Reduction: ${displaystyle 3e^{-}+color {Blue}2H_{2}Ocolor {Black}+MnO_{4}^{-}to MnO_{2}+color {Blue}4OH^{-}color {Black}}$

We balance the number of electrons as in the previous example.Oxidation: ${displaystyle (2OH^{-}+SO_{3}^{2-}to SO_{4}^{2-}+H_{2}O+color {OliveGreen}2e^{-}color {Black});color {Orange}times 3color {Black}}$ Reduction: ${displaystyle (color {Orange}3e^{-}color {Black}+2H_{2}O+MnO_{4}^{-}to MnO_{2}+4OH^{-});color {OliveGreen}times 2color {Black}}$

We obtain:Oxidation: ${displaystyle 6OH^{-}+3SO_{3}^{2-}to 3SO_{4}^{2-}+3H_{2}O+6e^{-}}$ Reduction: ${displaystyle 6e^{-}+4H_{2}O+2MnO_{4}^{-}to 2MnO_{2}+8OH^{-}}$

As you can see, the electrons are balanced, so we proceed to add the two half-reactions, to finally obtain the balanced equation. ${displaystyle {underline {left.{begin{array}{rcl}6OH^{-}+3SO_{3}^{2-}to 3SO_{4}^{2-}+3H_{2}O+6e^{-}\6e^{-}+4H_{2}O+2MnO_{4}^{-}to 2MnO_{2}+8OH^{-}end{array}}rightDownarrow +}}}$ ${displaystyle 2KMnO_{4}+3Na_{2}SO_{3}+H_{2}Oto 2MnO_{2}+3Na_{2}SO_{4}+2KOH}$

Applications

In industry, redox processes are also very important, both for their productive use (for example, the reduction of minerals to obtain aluminum or iron) and for their prevention (for example, in corrosion). The reverse reaction of the redox reaction (which produces energy) is electrolysis, in which energy is supplied to dissociate elements from their molecules.

Biological oxidations and reductions

In the metabolism of all living beings, redox processes are of paramount importance, since they are involved in the chain of chemical reactions of photosynthesis and aerobic respiration. In both reactions there is an electron transport chain formed by a series of enzymatic complexes, among which the cytochromes stand out; These enzymatic complexes accept (are reduced) and donate (oxidize) pairs of electrons in a sequential manner, in such a way that the first one gives up electrons to the second, this one to the third, etc., until a final acceptor that is definitively reduced; During their journey, the electrons release energy that is used to synthesize high-energy bonds in the form of ATP.

Another type of fundamental redox reaction in metabolic processes is dehydrogenation, in which an enzyme (dehydrogenase) removes a pair of hydrogen atoms from a substrate; Since the hydrogen atom consists of one proton and one electron, said substrate is oxidized (since it loses electrons). These electrons are captured by specialized molecules, mainly the coenzymes NAD+, NADP+and FAD that when gaining electrons are reduced, and lead them to the aforementioned electron transport chains. Metabolism involves hundreds of redox reactions. Thus, catabolism is constituted by reactions in which substrates are oxidized and coenzymes are reduced. In contrast, the reactions of anabolism are reactions in which substrates are reduced and coenzymes are oxidized. Taken together, catabolism and anabolism constitute metabolism.

Combustion

Combustion is a chemical reduction-oxidation reaction, in which a large amount of energy is generally released, in the form of heat and light, manifesting itself visually as fire. In all combustion there is an element that burns (fuel) and another that produces combustion (oxidant), generally oxygen in the form of O2gaseous. Explosives have chemically bound oxygen, so they do not need the oxygen in the air for combustion. The most frequent types of fuel are organic materials containing carbon and hydrogen (see hydrocarbons). In a complete reaction all elements have the highest oxidation state. The products formed are carbon dioxide (CO2) and water, sulfur dioxide (SO2) (if the fuel contains sulfur) and nitrogen oxides (NOx), depending on the temperature and the amount of oxygen in the reaction.

Impact

In metals, a very important consequence of oxidation is corrosion, a very negative structural impact phenomenon, given that materials acquire or modify their properties depending on the agents that are exposed, and how they act on them. Combining oxidation-reduction (redox) reactions in a galvanic cell, electrochemical batteries are achieved. These reactions can be used to avoid undesired corrosion phenomena by means of the sacrificial anode technique and to obtain direct electric current.

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