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Chapter 22: Problem 50

Write the chemical formula for each of the following compounds, and indicate the oxidation state of nitrogen in each: (a) dinitrogen trioxide, (b) dinitrogen pentoxide, (c) hydrogen cyanide, (d) ammonium nitrate, (e) iron(III) nitrite, \((\mathbf{f})\) nitrogen trichloride.

Short Answer

Expert verified
The chemical formulas and oxidation states of nitrogen for the given compounds are: a) \(N_2O_3\): +3 b) \(N_2O_5\): +5 c) \(HCN\): -3 d) \(NH_4NO_3\): -3 (in ammonium) and +5 (in nitrate) e) \(Fe(NO_2)_3\): +3 f) \(NCl_3\): +3

Step by step solution

01

Write the chemical formula

To write the chemical formula for each compound, we need to balance the charges of the elements involved. We'll do this for each compound listed: a) Dinitrogen trioxide: There are 2 nitrogen (N) atoms and 3 oxygen (O) atoms. Chemical formula: \(N_2O_3\) b) Dinitrogen pentoxide: There are 2 nitrogen (N) atoms and 5 oxygen (O) atoms. Chemical formula: \(N_2O_5\) c) Hydrogen cyanide: There is 1 hydrogen (H) atom, 1 carbon (C) atom, and 1 nitrogen (N) atom. Chemical formula: \(HCN\) d) Ammonium nitrate: This compound contains 1 ammonium ion (\(NH_4^+\)) and 1 nitrate ion (\(NO_3^-\)). Chemical formula: \(NH_4NO_3\) e) Iron(III) nitrite: This compound contains 1 iron (Fe) ion with a charge of +3 and 3 nitrite ions (\(NO_2^-\)) to balance the charge. Chemical formula: \(Fe(NO_2)_3\) f) Nitrogen trichloride: There is 1 nitrogen (N) atom and 3 chlorine (Cl) atoms. Chemical formula: \(NCl_3\)
02

Determine the oxidation state of nitrogen

We will apply the oxidation state rules to find the oxidation state of nitrogen in each compound: a) In \(N_2O_3\), the oxidation state of oxygen is typically -2. The overall charge of the molecule is 0, and there are two nitrogen atoms, so the combined oxidation state of the nitrogen atoms is +6 (3 x (-2) = -6). For each nitrogen atom, the oxidation state is +3. Oxidation state of N in \(N_2O_3\): +3 b) In \(N_2O_5\), the oxidation state of oxygen is typically -2. The overall charge of the molecule is 0, and there are two nitrogen atoms, so the combined oxidation state of the nitrogen atoms is +10 (5 x (-2) = -10). For each nitrogen atom, the oxidation state is +5. Oxidation state of N in \(N_2O_5\): +5 c) In \(HCN\), the oxidation state of hydrogen is typically +1, and carbon's oxidation state is +2 (since it's part of a cyanide ion). Nitrogen's oxidation state balances the charge in the cyanide ion. Oxidation state of N in \(HCN\): -3 d) In \(NH_4NO_3\), the ammonium ion (NH4+) has an oxidation state of +1 for hydrogen and -3 for nitrogen, and the nitrate ion (\(NO_3^-\)) has an oxidation state of -2 for oxygen and +5 for nitrogen. Oxidation states of N in \(NH_4NO_3\): -3 (in ammonium) and +5 (in nitrate) e) In \(Fe(NO_2)_3\), the iron ion has an oxidation state of +3, and the oxidation state of oxygen in nitrite ions is typically -2. Nitrogen balances the charge in the nitrite ion. Oxidation state of N in \(Fe(NO_2)_3\): +3 f) In \(NCl_3\), the oxidation state of chlorine is typically -1. Nitrogen balances the charge in the molecule. Oxidation state of N in \(NCl_3\): +3

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Key Concepts

These are the key concepts you need to understand to accurately answer the question.

Oxidation States
Oxidation states, also known as oxidation numbers, help us understand the distribution of electrons among the atoms in a chemical compound.
They indicate the electrical charge of an atom in a molecule or ion. Think of it as a bookkeeping method.
- **General Rules:** - Oxygen usually has an oxidation state of -2. - Hydrogen is typically +1. - The sum of oxidation states in a neutral compound must be zero. In ions, it must equal the charge of the ion.
For example, consider dinitrogen trioxide \( (N_2O_3) \): - Oxygen鈥檚 oxidation state is -2. To find nitrogen's oxidation state, we know the compound has no overall charge, so nitrogen atoms should have a combined oxidation state of +6 to balance three oxygen atoms \((-6)\). Since there are two nitrogen atoms, each must have an oxidation state of +3.
Dinitrogen Compounds
Dinitrogen compounds are molecules containing two nitrogen atoms bonded together. They can form various oxides depending on the number of oxygen atoms.
Such compounds play essential roles in both chemical manufacturing and in nature.- **Dinitrogen Trioxide (\(N_2O_3\))** - Has two nitrogen atoms and three oxygen atoms. - Commonly appears in discussions around nitrogen chemistry due to its unique properties.
- **Dinitrogen Pentoxide (\(N_2O_5\))** - Contains two nitrogen atoms and five oxygen atoms. - Often serves as a nitration agent in laboratory settings.
These compounds illustrate how varying numbers of oxygen atoms change properties and uses, while underpinning the importance of precise chemical formulas in describing their structures.
Ionic Compounds
Ionic compounds are formed from positively charged cations and negatively charged anions.
They are held together by an electrostatic force called ionic bonding, which creates a crystalline lattice structure.
- **Understanding Charges:** - Cations are typically metal ions, such as iron ( \(Fe^{3+}\)).
- Anions usually come from nonmetals, like the nitrite ion ( \(NO_2^-\)).
The stability of ionic compounds arises from the need to balance charges, which gives rise to neutral compounds overall.
For example, in iron(III) nitrite ( \(Fe(NO_2)_3\)), the +3 charge of iron is balanced by three -1 charges of nitrite ions.
Complex Ions
Complex ions are molecules consisting of a central metal atom bonded to surrounding ligands. These ligands are molecules or ions which donate electron pairs to the metal.
The result is a stable arrangement that characterizes many diverse chemical species.
- **Structure:** - Central metal atom (often a transition metal). - Ligands attach to this central atom, filling its electron orbitals and giving specific properties to the central ion.
The coordination number, which is the number of ligand attachment sites, and the oxidation state of the metal, greatly influence the complex ion's behavior.
This is evident in compounds like ammonia complexes or other nitrogen-based ligands, which impact the metal's reactivity and stability. In analysis, complex ions highlight the versatility and bonding intricacies in chemical compounds.

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Most popular questions from this chapter

The physical properties of \(\mathrm{D}_{2} \mathrm{O}\) differ from those of \(\mathrm{H}_{2} \mathrm{O}\) because (a) D has a different electron configuration than O. (b) \(\mathrm{D}\) is radioactive. (c) D forms stronger bonds with \(\mathrm{O}\) than \(\mathrm{H}\) does. (d) D is much more massive than H.

A sulfuric acid plant produces a considerable amount of heat. This heat is used to generate electricity, which helps reduce operating costs. The synthesis of \(\mathrm{H}_{2} \mathrm{SO}_{4}\) consists of three main chemical processes: (a) oxidation of \(S\) to \(\mathrm{SO}_{2}\), (b) oxidation of \(\mathrm{SO}_{2}\) to \(\mathrm{SO}_{3},\) (c) the dissolving of \(\mathrm{SO}_{3}\) in \(\mathrm{H}_{2} \mathrm{SO}_{4}\) and the subsequent reaction with water to form \(\mathrm{H}_{2} \mathrm{SO}_{4}\). If the third process produces \(130 \mathrm{~kJ} / \mathrm{mol}\), how much heat is produced in preparing a mole of \(\mathrm{H}_{2} \mathrm{SO}_{4}\) from a mole of \(S\) ? How much heat is produced in preparing \(2000 \mathrm{~kg}\) of \(\mathrm{H}_{2} \mathrm{SO}_{4} ?\)

Although the \(\mathrm{ClO}_{4}^{-}\) and \(\mathrm{IO}_{4}^{-}\) ions have been known for a long time, \(\mathrm{BrO}_{4}^{-}\) was not synthesized until \(1965 .\) The ion was synthesized by oxidizing the bromate ion with xenon difluoride, producing xenon, hydrofluoric acid, and the perbromate ion. (a) Write the balanced equation for this reaction. (b) What are the oxidation states of Br in the Br-containing species in this reaction?

Write a balanced equation for the reaction of each of the following compounds with water: (a) \(\mathrm{PCl}_{5}(s),\) (b) \(\mathrm{CO}_{2}(g)\) (c) \(\mathrm{K}_{2} \mathrm{O}_{2}(s)\) (d) \(\mathrm{Mg}_{3} \mathrm{P}_{2}(s)\) (e) \(\operatorname{LiAlH}_{4}(s)\) (f) \(\mathrm{Cl}_{2} \mathrm{O}(g)\) (g) \(\mathrm{NO}_{2}(g)\)

Name the following compounds and assign oxidation states to the halogens in them: (a) \(\mathrm{BCl}_{3}\), (b) \(\mathrm{Sr}\left(\mathrm{IO}_{4}\right)_{2}\), (c) \(\mathrm{LiOCl}\) (d) \(\mathrm{HClO}\) (e) \(\mathrm{CuClO}\) (f) \(\mathrm{Mg}\left(\mathrm{IO}_{2}\right)\).
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