If an Element X Can Form an Oxide

Oxidation and Reduction

Oxidation-Reduction Reactions

The term oxidation was originally used to draw reactions in which an element combines with oxygen.

Example: The reaction between magnesium metallic and oxygen to class magnesium oxide involves the oxidation of magnesium.

The term reduction comes from the Latin stem meaning "to lead back." Anything that that leads back to magnesium metal therefore involves reduction.

The reaction between magnesium oxide and carbon at 2000C to grade magnesium metal and carbon monoxide is an example of the reduction of magnesium oxide to magnesium metallic.

Later electrons were discovered, chemists became convinced that oxidation-reduction reactions involved the transfer of electrons from one atom to another. From this perspective, the reaction betwixt magnesium and oxygen is written as follows.

2 Mg + O₂ ii [Mg²⁺][O^two-]

In the form of this reaction, each magnesium atom loses two electrons to form an Mg²⁺ ion.

Mg Mgⁱⁱ⁺ + two eastward^-

And, each O₂ molecule gains iv electrons to form a pair of O^2- ions.

O₂ + four e^- 2 O^two-

Because electrons are neither created nor destroyed in a chemic reaction, oxidation and reduction are linked. It is incommunicable to accept i without the other, equally shown in the figure below.

The Role of Oxidation Numbers in Oxidation-Reduction Reactions

Chemists eventually extended the idea of oxidation and reduction to reactions that exercise not formally involve the transfer of electrons.

Consider the following reaction.

CO(g) + H_iiO(g) CO₂(g) + H_ii(grand)

As tin can exist seen in the figure below, the total number of electrons in the valence shell of each cantlet remains constant in this reaction.

What changes in this reaction is the oxidation country of these atoms. The oxidation country of carbon increases from +ii to +four, while the oxidation state of the hydrogen decreases from +i to 0.

Oxidation and reduction are therefore best divers as follows. Oxidation occurs when the oxidation number of an atom becomes larger. Reduction occurs when the oxidation number of an atom becomes smaller.

Oxidation Numbers Versus the Truthful Charge on Ions

The terms ionic and covalent describe the extremes of a continuum of bonding. There is some covalent grapheme in fifty-fifty the well-nigh ionic compounds and vice versa.

It is useful to retrieve nigh the compounds of the main group metals every bit if they contained positive and negative ions. The chemistry of magnesium oxide, for example, is easy to understand if we assume that MgO contains Mg²⁺ and O^ii- ions. But no compounds are 100% ionic. There is experimental bear witness, for example, that the true charge on the magnesium and oxygen atoms in MgO is +1.5 and -1.5.

Oxidation states provide a compromise between a powerful model of oxidation-reduction reactions based on the assumption that these compounds comprise ions and our noesis that the true accuse on the ions in these compounds is not as large as this model predicts. By definition, the oxidation land of an atom is the accuse that atom would carry if the compound were purely ionic.

For the agile metals in Groups IA and IIA, the departure betwixt the oxidation land of the metal cantlet and the charge on this atom is modest enough to be ignored. The master group metals in Groups IIIA and IVA, yet, form compounds that accept a significant amount of covalent character. Information technology is misleading, for case, to presume that aluminum bromide contains Al^three+ and Br^- ions. Information technology actually exists equally Al₂Br_vi molecules.

This problem becomes even more severe when we turn to the chemistry of the transition metals. MnO, for instance, is ionic enough to be considered a common salt that contains Mn²⁺ and O^2- ions. Mn_twoO_seven, on the other hand, is a covalent compound that boils at room temperature. It is therefore more useful to call up virtually this compound as if it independent manganese in a +seven oxidation land, not Mn⁷⁺ ions.

Oxidizing Agents and Reducing Agents

Let's consider the role that each element plays in the reaction in which a particular chemical element gains or loses electrons..

When magnesium reacts with oxygen, the magnesium atoms donate electrons to O_two molecules and thereby reduce the oxygen. Magnesium therefore acts equally a reducing agent in this reaction.

two Mg + O₂ 2 MgO

reducing
agent

The O_two molecules, on the other hand, gain electrons from magnesium atoms and thereby oxidize the magnesium. Oxygen is therefore an oxidizing amanuensis.

2 Mg + O₂ ii MgO

oxidizing
amanuensis

Oxidizing and reducing agents therefore tin can be defined as follows. Oxidizing agents proceeds electrons. Reducing agents lose electrons.

The table below identifies the reducing amanuensis and the oxidizing agent for some of the reactions discussed in this web page. Ane tendency is immediately obvious: The chief group metals act equally reducing agents in all of their chemical reactions.

Typical Reactions of Main Group Metals

Conjugate Oxidizing Amanuensis/Reducing Agent Pairs

Metals deed as reducing agents in their chemical reactions. When copper is heated over a flame, for instance, the surface slowly turns black as the copper metallic reduces oxygen in the atmosphere to form copper(II) oxide.

If nosotros plow off the flame, and blow H₂ gas over the hot metal surface, the black CuO that formed on the surface of the metal is slowly converted back to copper metal. In the grade of this reaction, CuO is reduced to copper metal. Thus, H₂ is the reducing agent in this reaction, and CuO acts equally an oxidizing agent.

An important feature of oxidation-reduction reactions can be recognized by examining what happens to the copper in this pair of reactions. The first reaction converts copper metal into CuO, thereby transforming a reducing agent (Cu) into an oxidizing agent (CuO). The second reaction converts an oxidizing agent (CuO) into a reducing agent (Cu). Every reducing agent is therefore linked, or coupled, to a conjugate oxidizing agent, and vice versa.

Every time a reducing agent loses electrons, information technology forms an oxidizing amanuensis that could proceeds electrons if the reaction were reversed.

Conversely, every time an oxidizing agent gains electrons, information technology forms a reducing amanuensis that could lose electrons if the reaction went in the opposite management.

The idea that oxidizing agents and reducing agents are linked, or coupled, is why they are called conjugate oxidizing agents and reducing agents. Conjugate comes from the Latin stem meaning "to join together." It is therefore used to describe things that are linked or coupled, such every bit oxidizing agents and reducing agents.

The main grouping metals are all reducing agents. They tend to exist "strong" reducing agents. The active metals in Group IA, for case, give upwards electrons better than whatever other elements in the periodic table.

The fact that an active metal such as sodium is a strong reducing agent should tell us something about the relative strength of the Na⁺ ion every bit an oxidizing amanuensis. If sodium metal is relatively good at giving upwardly electrons, Na⁺ ions must be unusually bad at picking upwards electrons. If Na is a strong reducing agent, the Na⁺ ion must be a weak oxidizing agent.

Conversely, if O_two has such a high affinity for electrons that information technology is unusually good at accepting them from other elements, it should be able to hang onto these electrons once it picks them up. In other words, if O₂ is a strong oxidizing agent, then the O^two- ion must be a weak reducing agent.

In full general, the human relationship betwixt cohabit oxidizing and reducing agents can be described as follows. Every strong reducing agent (such equally Na) has a weak conjugate oxidizing agent (such as the Na ⁺ ion). Every strong oxidizing agent (such as O _two ) has a weak cohabit reducing amanuensis (such as the O ^two- ion).

The Relative Force of Metals as Reducing Agents

We tin determine the relative strengths of a pair of metals as reducing agents past determining whether a reaction occurs when ane of these metals is mixed with a salt of the other. Consider the relative strength of iron and aluminum, for example. Nothing happens when we mix powdered aluminum metallic with iron(III) oxide. If nosotros place this mixture in a crucible, however, and get the reaction started by applying a trivial heat, a vigorous reaction takes place to requite aluminum oxide and molten atomic number 26 metal.

ii Al(south) + Iron_iiO_three(s) Al₂O₃(s) + 2 Fe(l)

Past assigning oxidation numbers, we can pick out the oxidation and reduction halves of the reaction.

Aluminum is oxidized to Al_twoO₃ in this reaction, which means that Fe_iiO₃ must exist the oxidizing agent. Conversely, Fe_iiO₃ is reduced to atomic number 26 metallic, which means that aluminum must be the reducing agent. Because a reducing agent is always transformed into its conjugate oxidizing agent in an oxidation-reduction reaction, the products of this reaction include a new oxidizing agent (Al₂O_iii) and a new reducing amanuensis (Atomic number 26).

Since the reaction proceeds in this direction, information technology seems reasonable to presume that the starting materials comprise the stronger reducing agent and the stronger oxidizing amanuensis.

In other words, if aluminum reduces Iron₂O_three to class Al₂O₃ and iron metal, aluminum must be a stronger reducing agent than iron.

We can conclude from the fact that aluminum cannot reduce sodium chloride to form sodium metallic that the starting materials in this reaction are the weaker oxidizing agent and the weaker reducing amanuensis.

Nosotros can test this hypothesis past request: What happens when we try to run the reaction in the opposite direction? (Is sodium metal strong enough to reduce a table salt of aluminum to aluminum metallic?) When this reaction is run, we find that sodium metal can, in fact, reduce aluminum chloride to aluminum metal and sodium chloride when the reaction is run at temperatures hot enough to melt the reactants.

3 Na(fifty) + AlCl₃(50) iii NaCl(l) + Al(l)

If sodium is strong enough to reduce Al³⁺ salts to aluminum metal and aluminum is strong enough to reduce Fe³⁺ salts to atomic number 26 metal, the relative strengths of these reducing agents can be summarized as follows.

Na > Al > Iron