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Magnesium oxide reduction

It has been ly suggested that Mg g is produced by either i MgO dissociation forming O 2 as the reaction intermediate or ii MgO s —C s boundary reaction producing CO that then reduces MgO while forming CO 2 as the reaction intermediate. Either of the intermediates O 2 or CO 2 are then consumed by C, which is necessary to sustain further Mg g production.

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Oxidation-Reduction Reactions. The term oxidation was originally used to describe reactions in which an element combines with oxygen. Example: The reaction between magnesium metal and oxygen to form magnesium oxide involves the oxidation of magnesium.

The term reduction comes from the Latin stem meaning "to lead back. The reaction between magnesium oxide and carbon at C to form magnesium metal and carbon monoxide is an example of the reduction of magnesium oxide to magnesium metal. After 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 between magnesium and oxygen is written as follows. Because electrons are neither created nor destroyed in a chemical reaction, oxidation and reduction are linked. It is impossible to have one without the other, as shown in the figure below. Determine which element is oxidized and which is reduced when lithium reacts with nitrogen to form lithium nitride.

Chemists eventually extended the idea of oxidation and reduction to reactions that do not formally involve the transfer of electrons. As can be seen in the figure below, the total of electrons in the valence shell of each atom remains constant in this reaction.

What changes in this reaction is the oxidation state of these atoms. Oxidation and reduction are therefore best defined as follows. Oxidation occurs when the oxidation of an atom becomes larger. Reduction occurs when the oxidation of an atom becomes smaller.

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Determine which atom is oxidized and which is reduced in the following reaction. The terms ionic and covalent describe the extremes of a continuum of bonding. There is some covalent character in even the most ionic compounds and vice versa. It is useful to think about the compounds of the main group metals as if they contained positive and negative ions.

Oxidation states provide a compromise between a powerful model of oxidation-reduction reactions based on the assumption that these compounds contain ions and our knowledge that the true charge on the ions in these compounds is not as large as this model predicts. By definition, the oxidation state of an atom is the charge that reduction would carry if the compound were purely ionic. For the active metals in Groups IA and IIA, the magnesium between the oxidation state of the metal atom and the charge on this atom is small enough to be ignored. It actually exists as Al 2 Br 6 molecules.

This problem becomes even more severe when we turn to the chemistry of the transition metals. Mn 2 O 7on the other hand, is a covalent compound that boils at room temperature. Let's consider the role that each element plays in the reaction in which a particular element gains or loses electrons. When magnesium reacts with oxygen, the magnesium atoms donate electrons to O 2 molecules and thereby reduce the oxygen. Magnesium therefore acts as a reducing agent in this reaction. The O 2 molecules, on the oxide hand, gain electrons from magnesium atoms and thereby oxidize the magnesium.

Oxygen is therefore an oxidizing agent.

Oxidizing and reducing agents therefore can be defined as follows. Oxidizing agents gain electrons.

Reducing agents lose electrons. The table below identifies the reducing agent and the oxidizing agent for some of the reactions discussed in this web. One trend is immediately obvious: The main group metals act as reducing agents in all of their chemical reactions. Metals act as reducing agents in their chemical reactions. When copper is heated over a flame, for example, the surface slowly turns black as the copper metal reduces oxygen in the atmosphere to form copper II oxide.

If we turn off the flame, and blow H 2 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 course of this reaction, CuO is reduced to copper metal.

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Thus, H 2 is the reducing agent in this reaction, and CuO acts as 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, it forms an oxidizing agent that could gain electrons if the reaction were reversed.

Conversely, every time an oxidizing agent gains electrons, it forms a reducing agent that could lose electrons if the reaction went in the opposite direction. 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 together. The main group metals are all reducing agents.

They tend to be "strong" reducing agents. The active metals in Group IA, for reduction, give up electrons better than any other elements in the periodic magnesium. Conversely, if O 2 has such a high affinity for electrons that it 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 O 2 is a strong oxidizing agent, then the O 2- ion must be a weak reducing agent. In general, the relationship between conjugate oxidizing and reducing agents can be described as follows. Every strong oxidizing agent such as O 2 has a weak conjugate oxide agent such as the O 2- ion.

We can determine the relative strengths of a pair of metals as reducing agents by determining whether a reaction occurs when one 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 metal with iron III oxide.

If we place this mixture in a crucible, however, and get the reaction started by applying a little heat, a vigorous reaction takes place to give aluminum oxide and molten iron metal. By asing oxidation s, we can pick out the oxidation and reduction halves of the reaction. Aluminum is oxidized to Al 2 O 3 in this reaction, which means that Fe 2 O 3 must be the oxidizing agent.

Conversely, Fe 2 O 3 is reduced to reduction metal, 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 Al 2 O 3 and a new reducing agent Fe. Since the reaction proceeds in this direction, it seems reasonable to assume that the starting materials contain the stronger reducing agent and the stronger oxide agent.

In other words, if aluminum reduces Fe 2 O 3 to form Al 2 O 3 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 metal that the reduction materials in this reaction are the weaker magnesium agent and the weaker reducing agent. We can test this hypothesis by asking: What happens when we try to run the reaction in the opposite direction?

Is sodium metal strong enough to reduce a salt of aluminum to aluminum metal? 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. Use the oxide equations to determine the relative strengths of magnesium, magnesium, aluminum, and calcium metal as reducing agents.

Salts Oxidation and Reduction. Practice Problem 1: Determine which element is oxidized and which is reduced when lithium reacts with nitrogen to form lithium nitride. Interactive tutorial on asing oxidation s from requires Macromedia Shockwave. Practice Problem 3: Identify the oxidizing agent and the reducing agent in the following reaction. Practice Problem 4: Use the following equations to determine the relative strengths of sodium, magnesium, aluminum, and calcium metal as reducing agents.