Question

For each of the following oxidation-reduction reactions of metals with nonmetals, identify which element is oxidized and which is reduced. a. \(2 \mathrm{K}(s)+\mathrm{O}_{2}(g) \rightarrow \mathrm{K}_{2} \mathrm{O}(s)\) b. \(\mathrm{H}_{2}(g)+\mathrm{S}(s) \rightarrow \mathrm{H}_{2} \mathrm{S}(g)\) c. \(\mathrm{N}_{2}(g)+3 \mathrm{H}_{2}(g) \rightarrow 2 \mathrm{NH}_{3}(g)\) d. \(2 \mathrm{Hg}(l)+\mathrm{Cl}_{2}(g) \rightarrow \mathrm{Hg}_{2} \mathrm{Cl}_{2}(s)\)

Short answer

In these four reactions: a. K is oxidized and O is reduced. b. H is oxidized and S is reduced. c. N is reduced and H is oxidized. d. Hg is oxidized and Cl is reduced.

Step-by-step solution

Reaction a: Identify oxidized and reduced element

First, let's write down the oxidation states of the elements before and after the reaction: 2K(s) + O₂(g) → K₂O(s) Before: K(0) + O(0) After: K(+1) + O(-2) Here, we can see that potassium (K) goes from an oxidation state of 0 to +1, losing one electron, so it is oxidized. Oxygen (O) goes from 0 to -2, gaining two electrons, so it is reduced.

Reaction b: Identify oxidized and reduced element

Next, let's look at the oxidation states of the elements in Reaction b: H₂(g) + S(s) → H₂S(g) Before: H(0) + S(0) After: H(+1) + S(-2) In this reaction, hydrogen (H) goes from an oxidation state of 0 to +1, losing one electron, so it is oxidized. Sulfur (S) goes from 0 to -2, gaining two electrons, so it is reduced.

Reaction c: Identify oxidized and reduced element

Now let's look at the oxidation states of the elements in Reaction c: N₂(g) + 3H₂(g) → 2NH₃(g) Before: N(0) + H(0) After: N(-3) + H(+1) In this reaction, nitrogen (N) goes from an oxidation state of 0 to -3, gaining three electrons, so it is reduced. Hydrogen (H) goes from 0 to +1, losing one electron, so it is oxidized.

Reaction d: Identify oxidized and reduced element

Finally, let's look at the oxidation states of the elements in Reaction d: 2Hg(l) + Cl₂(g) → Hg₂Cl₂(s) Before: Hg(0) + Cl(0) After: Hg(+1) + Cl(-1) In this reaction, mercury (Hg) goes from an oxidation state of 0 to +1, losing one electron, so it is oxidized. Chlorine (Cl) goes from 0 to -1, gaining one electron, so it is reduced.

Key concepts

Oxidation States

Oxidation states, also known as oxidation numbers, are numerical values assigned to atoms to reflect their degree of oxidation or reduction. These numbers help track electron flow in chemical reactions. Knowing the oxidation state is crucial in identifying what happens during oxidation-reduction, or redox, reactions, where one element is oxidized and another is reduced.
- **Assigning Oxidation States**: - For pure elements in their natural form, the oxidation state is always 0. - Common rules include: - Oxygen is usually (-2) in compounds. - Hydrogen is typically (+1) . - Alkali metals, like potassium, have an oxidation state of (+1) . Let's see how knowing oxidation states aids our understanding of a reaction. In the reaction where 2K(s) + O₂(g) → K₂O(s), potassium changes from 0 to +1 , while oxygen goes from 0 to -2 . These changes signal a transfer of electrons, showing that potassium is oxidized, and oxygen is reduced.

Electron Transfer

In the heart of all oxidation-reduction reactions lies the transfer of electrons from one substance to another. Understanding electron transfer is crucial because it directly impacts the oxidation states of the elements involved.
- **Electron Loss and Gain**: - Oxidation refers to the loss of electrons, which increases an element's oxidation state. - Reduction involves the gain of electrons, resulting in a decrease in oxidation state. Considering the reaction H₂(g) + S(s) → H₂S(g) : - **Hydrogen** sheds electrons, shifting from an oxidation state of 0 to +1 , signifying it's been oxidized. - **Sulfur**, on the other hand, gains those electrons, moving from 0 to -2 , hence it's reduced.
Such electron movement is essential for energy transfer in many chemical and biological processes. By knowing which elements donate and receive electrons, we deeply understand how compounds form and breakdown.

Redox Reactions

Redox, short for reduction-oxidation reactions, involves simultaneous processes where one substance undergoes oxidation, while another undergoes reduction. These reactions are ubiquitous in chemistry and biology, affecting everyday life from rusting metals to energy metabolism in cells.
- **Identifying Redox Reactions**: - The key to redox reactions is finding the changes in oxidation states across a chemical equation. - Each reaction consists of two half-reactions: one for oxidation and one for reduction. Take the reaction N₂(g) + 3H₂(g) → 2NH₃(g) : - **Nitrogen** is reduced, accepting electrons to transition from an oxidation state of 0 to -3 . - **Hydrogen** is oxidized, its state increasing from 0 to +1 .
Understanding how redox reactions function helps in comprehending numerous processes, including combustion and photosynthesis. Mastering these concepts means recognizing the electron exchange and the interlinked nature of substance oxidation and reduction.