/*! This file is auto-generated */ .wp-block-button__link{color:#fff;background-color:#32373c;border-radius:9999px;box-shadow:none;text-decoration:none;padding:calc(.667em + 2px) calc(1.333em + 2px);font-size:1.125em}.wp-block-file__button{background:#32373c;color:#fff;text-decoration:none} Problem 52 The nitrogen atoms in \(\mathrm{... [FREE SOLUTION] | 91影视

91影视

Log out

Learning Materials

Features

Discover

Chapter 9: Problem 52

The nitrogen atoms in \(\mathrm{N}_{2}\) participate in multiple bonding, whereas those in hydrazine, \(\mathrm{N}_{2} \mathrm{H}_{4}\), do not. (a) Draw Lewis structures for both molecules. (b) What is the hybridization of the nitrogen atoms in each molecule? (c) Which molecule has a stronger \(\mathrm{N}-\mathrm{N}\) bond?

Short Answer

Expert verified
The Lewis structures are N鈮 for N2 and H2N-NH2 for N2H4. The nitrogen atoms in N2 are sp hybridized, while those in N2H4 are sp3 hybridized. The N-N bond in N2 (triple bond) is stronger than the N-N bond in N2H4 (single bond).

Step by step solution

01

Draw Lewis Structures

To draw the Lewis structures for N2 and N2H4, start by counting the total number of valence electrons in each molecule. Nitrogen has 5 valence electrons, and hydrogen has 1. N2: 2 nitrogen atoms * 5 electrons each = 10 valence electrons N2H4: 2 nitrogen atoms * 5 electrons each + 4 hydrogen atoms * 1 electron each = 14 valence electrons Now, draw the Lewis structures. For N2: 1. Place two nitrogen atoms together, with one single bond. 2. Add 3 more pairs of electrons to each nitrogen atom to complete the octet. 3. To minimize formal charge, turn one lone pair into a bonding pair, resulting in a triple bond. Lewis Structure for N2: N鈮 For N2H4: 1. Place two nitrogen atoms together with a single bond. 2. Add 2 hydrogen atoms to each nitrogen atom, with single bonds. 3. Add 2 remaining electrons (1 lone pair) to each nitrogen atom. Lewis Structure for N2H4: H2N-NH2
02

Determine Hybridization

To find the hybridization, count the electron groups around each nitrogen atom. Electron groups include single bonds, double bonds, triple bonds, and lone pairs. For N2: Each nitrogen atom has 1 triple bond and 1 lone pair, which corresponds to 2 electron groups. This suggests sp hybridization. For N2H4 (hydrazine): Each nitrogen atom has 2 single bonds to hydrogen atoms, 1 single bond to another nitrogen atom, and 1 lone pair, which corresponds to 4 electron groups. This suggests sp3 hybridization. Nitrogen in N2: sp hybridization Nitrogen in N2H4: sp3 hybridization
03

Compare N-N Bond Strength

The N-N bond strength generally increases in the order single bond < double bond < triple bond. As a result, the N-N bond in N2 (triple bond) is stronger than the N-N bond in N2H4 (single bond). The molecule with a stronger N-N bond is N2.

Unlock Step-by-Step Solutions & Ace Your Exams!

  • Full Textbook Solutions

    Get detailed explanations and key concepts

  • Unlimited Al creation

    Al flashcards, explanations, exams and more...

  • Ads-free access

    To over 500 millions flashcards

  • Money-back guarantee

    We refund you if you fail your exam.

Over 30 million students worldwide already upgrade their learning with 91影视!

One App. One Place for Learning.

All the tools & learning materials you need for study success - in one app.

Get started for free

Most popular questions from this chapter

The lactic acid molecule, \(\mathrm{CH}_{3} \mathrm{CH}(\mathrm{OH}) \mathrm{COOH}\), gives sour milk its unpleasant, sour taste. (a) Draw the Lewis structure for the molecule, assuming that carbon always forms four bonds in its stable compounds. (b) How many \(\pi\) and how many \(\sigma\) bonds are in the molecule? (c) Which \(\mathrm{CO}\) bond is shortest in the molecule? (d) What is the hybridization of atomic orbitals around each carbon atom associated with that short bond? (e) What are the approximate bond angles around each carbon atom in the molecule?

Antibonding molecular orbitals can be used to make bonds to other atoms in a molecule. For example, metal atoms can use appropriate \(d\) orbitals to overlap with the \(\pi_{2 p}^{*}\) orbitals of the carbon monoxide molecule. This is called \(d-\pi\) backbonding. (a) Draw a coordinate axis system in which the \(y\) -axis is vertical in the plane of the paper and the \(x\) -axis horizontal. Write \(^{\prime} \mathrm{M}^{\prime \prime}\) at the origin to denote a metal atom. (b) Now, on the \(x\) -axis to the right of \(M\), draw the Lewis structure of a CO molecule, with the carbon nearest the \(\mathrm{M}\). The \(\mathrm{CO}\) bond axis should be on the \(x\) -axis. (c) Draw the CO \(\pi_{2 p}^{*}\) orbital, with phases (see the Closer Look box on phases) in the plane of the paper. Two lobes should be pointing toward \(\mathrm{M}\). (d) Now draw the \(\mathrm{d}_{x y}\) orbital of \(\mathrm{M}\), with phases. Can you see how they will overlap with the \(\pi_{2 p}^{*}\) orbital of CO? (e) What kind of bond is being made with the orbitals between \(\mathrm{M}\) and \(\mathrm{C}, \sigma\) or \(\pi ?\) (f) Predict what will happen to the strength of the CO bond in a metal-CO complex compared to CO alone.

Draw sketches illustrating the overlap between the following orbitals on two atoms: (a) the 2 s orbital on each atom, (b) the \(2 p_{z}\) orbital on each atom (assume both atoms are on the z-axis), (c) the \(2 s\) orbital on one atom and the \(2 p_{z}\) orbital on the other atom.

Consider the \(\mathrm{H}_{2}{ }^{+}\) ion. (a) Sketch the molecular orbitals of the ion, and draw its energy-level diagram. (b) How many electrons are there in the \(\mathrm{H}_{2}{ }^{+}\) ion? (c) Draw the electron configuration of the ion in terms of its MOs (d) What is the bond order in \(\mathrm{H}_{2}{ }^{+}\) ? (e) Suppose that the ion is excited by light so that an electron moves from a lower-energy to a higherenergy MO. Would you expect the excitedstate \(\mathrm{H}_{2}{ }^{+}\) ion to be stable or to fall apart? Explain.

(a) What is the probability of finding an electron on the internuclear axis if the electron occupies a \(\pi\) molecular orbital? (b) For a homonuclear diatomic molecule, what similarities and differences are there between the \(\pi_{2 p}\) MO made from the \(2 p_{x}\) atomic orbitals and the \(\pi_{2 p}\) MO made from the \(2 p_{y}\) atomic orbitals? (c) Why are the \(\pi_{2 p}\) MOs lower in energy than the \(\pi_{2 p}^{*}\) MOs?
See all solutions

Recommended explanations on Chemistry Textbooks

Chemical Analysis

Read Explanation

The Earths Atmosphere

Read Explanation

Organic Chemistry

Read Explanation

Kinetics

Read Explanation

Chemical Reactions

Read Explanation

Inorganic Chemistry

Read Explanation
View all explanations

What do you think about this solution?

We value your feedback to improve our textbook solutions.

Study anywhere. Anytime. Across all devices.