Question

For each of the following oxidation-reduction reactions of metals with nonmetals, identify which element is oxidized and which is reduced. a. \(6 \mathrm{Na}(s)+\mathrm{N}_{2}(g) \rightarrow 2 \mathrm{Na}_{3} \mathrm{N}(s)\) b. \(\mathrm{Mg}(s)+\mathrm{Cl}_{2}(g) \rightarrow \mathrm{MgCl}_{2}(s)\) c. \(2 \mathrm{Al}(s)+3 \mathrm{Br}_{2}(l) \rightarrow 2 \mathrm{AlBr}_{3}(s)\) d. \(4 \mathrm{Fe}(s)+3 \mathrm{O}_{2}(g) \rightarrow 2 \mathrm{Fe}_{2} \mathrm{O}_{3}(s)\)

Short answer

In conclusion: a. N is reduced. b. Mg is oxidized, and Cl is reduced. c. Al is oxidized, and Br is reduced. d. Fe is oxidized, and O is reduced.

Step-by-step solution

Determine the initial oxidation states

In order to determine the change in oxidation states, we first need to identify the initial oxidation states of the elements. a. For Na and N₂: Na is a metal (group 1), so its initial oxidation state is +1. N₂ is a diatomic molecule with an oxidation state of 0. b. For Mg and Cl₂: Mg is a metal (group 2), so its initial oxidation state is +2. Cl₂ is a diatomic molecule with an oxidation state of 0. c. For Al and Br₂: Al is a metal (group 3), so its initial oxidation state is +3. Br₂ is a diatomic molecule with an oxidation state of 0. d. For Fe and O₂: Fe is a metal (group 8), and its common oxidation states are +2 and +3. O₂ is a diatomic molecule with an oxidation state of 0.

Determine the final oxidation states

Next, we will determine the final oxidation states of the elements after the reactions. a. For Na₃N: Na has an oxidation state of +1. N has an oxidation state of -3. b. For MgCl₂: Mg has an oxidation state of +2. Cl has an oxidation state of -1. c. For AlBr₃: Al has an oxidation state of +3. Br has an oxidation state of -1. d. For Fe₂O₃: Fe has an oxidation state of +3. O has an oxidation state of -2.

Identify the oxidized and reduced elements

Now that we have the initial and final oxidation states, we can determine which elements are oxidized and which are reduced. a. Na: +1 to +1; no change, neither oxidation nor reduction N: 0 to -3; gain of electrons, reduction b. Mg: +2 to +2; loss of electrons, oxidation Cl: 0 to -1; gain of electrons, reduction c. Al: +3 to +3; loss of electrons, oxidation Br: 0 to -1; gain of electrons, reduction d. Fe: 0 to +3; loss of electrons, oxidation O: 0 to -2; gain of electrons, reduction In conclusion: a. N is reduced. b. Mg is oxidized, and Cl is reduced. c. Al is oxidized, and Br is reduced. d. Fe is oxidized, and O is reduced.

Key concepts

Oxidation States

Oxidation states, also known as oxidation numbers, are a fundamental concept in chemistry used to describe the degree of oxidation of an atom in a chemical compound. Essentially, they are hypothetical charges that an atom would have if all bonds to atoms of different elements were completely ionic.

When determining oxidation states, there are some basic rules to follow: elements in their elemental form have an oxidation state of 0; for monoatomic ions, the oxidation state is equal to the charge of the ion; and for a molecule, the sum of oxidation states of all atoms must equal the overall charge of the molecule or ion.

In metal-nonmetal reactions, metals typically have positive oxidation states because they tend to lose electrons (they are oxidized), whereas nonmetals generally have negative oxidation states as they gain electrons (they are reduced). Recognizing changes in oxidation states during a reaction is paramount to identifying redox processes, where the movement of electrons from one element to another is a key feature.

Redox Reactions in Chemistry

Redox reactions, or oxidation-reduction reactions, are processes that involve the transfer of electrons between two species. These reactions are crucial in both natural processes, such as respiration and photosynthesis, and industrial applications, like batteries and rusting.

There are two parts to a redox reaction: oxidation, where an element loses electrons, and reduction, where an element gains electrons. Importantly, oxidation and reduction always occur together; when one species is oxidized, another is simultaneously reduced, maintaining an overall balance of charge.

To identify redox reactions, we can observe changes in oxidation states. If the oxidation state of an element increases, it indicates oxidation; if the oxidation state decreases, it signals reduction. For example, in reaction 'b' from the exercise, Mg goes from an oxidation state of +2 to +2, losing electrons, and hence, is oxidized. Cl, on the other hand, goes from 0 to -1, gaining electrons, and is reduced.

Metal-Nonmetal Reactions

Metal-nonmetal reactions are a typical category of redox reactions where metals usually lose electrons to nonmetals, resulting in the formation of ionic compounds. Metals, having fewer valence electrons, are more inclined to give up those electrons, becoming positively charged ions, also known as cations. Nonmetals with a high electron affinity accept these electrons to become negatively charged ions, or anions.

In the given reactions, metals such as Na, Mg, Al, and Fe undergo oxidation by losing electrons and nonmetals like N, Cl, Br, and O are reduced by gaining electrons. These reactions are quintessential redox reactions because they clearly demonstrate the transfer of electrons from metals to nonmetals.

Important to Consider

• The type of metal and nonmetal involved can affect the reaction's outcome regarding the ionic compound produced.
• The concept of electronegativity is useful here, as it helps predict which atom will gain or lose electrons.
• In practical applications, metal-nonmetal reactions are instrumental in metal extraction and the formation of a vast array of useful products.

For students seeking to fully grasp redox processes, it is critical to not only memorize rules but also to understand the underlying principles such as electron transfer and energy changes which drive these reactions.