/*! 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 18 Use the localized electron model... [FREE SOLUTION] | 91影视

91影视

Log out

Learning Materials

Features

Discover

Chapter 9: Problem 18

Use the localized electron model to describe the bonding in \(\mathrm{CCl}_{4}\).

Short Answer

Expert verified
The bonding in carbon tetrachloride (CCl4) using the localized electron model can be described as follows: A central Carbon (C) atom covalently bonds to four surrounding Chlorine (Cl) atoms in a tetrahedral arrangement with 鈭燙l-C-Cl bond angles of 109.5掳. Each Carbon-Chlorine bond is a sigma bond formed by the overlapping of atomic orbitals between Carbon and Chlorine atoms.

Step by step solution

01

Determine the electron configuration of Carbon (C) and Chlorine (Cl) atoms

We first need to determine the electron configurations for both Carbon (C) and Chlorine (Cl) atoms: - Carbon (C) - Atomic number: 6 Electron configuration: \(1s^2 2s^2 2p^2\) - Chlorine (Cl) - Atomic number: 17 Electron configuration: \(1s^2 2s^2 2p^6 3s^2 3p^5\)
02

Calculate the number of valence electrons contributed by each atom

Next, we need to determine the valence electrons of Carbon and Chlorine atoms: - Carbon (C) - Valence electrons: 4 - Chlorine (Cl) - Valence electrons: 7
03

Find the total number of electrons in the molecule

We can now calculate the total number of electrons in the CCl4 molecule. Carbon contributes 4 electrons, and each Chlorine atom contributes 7 electrons: Total electrons = (1 x 4) + (4 x 7) = 4 + 28 = 32 electrons
04

Determine the structure of CCl4 using the localized electron model

Using the localized electron model, we can now determine the structure of the CCl4 molecule: 1. Place Carbon (C) in the center as it has the lowest electronegativity among the atoms. 2. Arrange the Chlorine atoms (Cl) around the Carbon atom (tetrahedral geometry). 3. Distribute the remaining electrons around the Chlorine atoms as lone pairs to complete their octet. Now, regarding the CCl4 molecule: - Carbon (C) forms 4 single covalent bonds with each Chlorine atoms. So, 鈭燙l-C-Cl bond angle = 109.5掳 (tetrahedral arrangement) - Each Carbon-Chlorine bond is a sigma bond formed by overlapping atomic orbitals between Carbon and Chlorine. In conclusion, the bonding in carbon tetrachloride (CCl4) using the localized electron model is characterized by a central carbon atom covalently bonded to four surrounding chlorine atoms in a tetrahedral geometry with 鈭燙l-C-Cl bond angles of 109.5掳.

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

As the head engineer of your starship in charge of the warp drive, you notice that the supply of dilithium is critically low. While searching for a replacement fuel, you discover some diboron, \(\mathrm{B}_{2}\). a. What is the bond order in \(\mathrm{Li}_{2}\) and \(\mathrm{B}_{2}\) ? b. How many electrons must be removed from \(\mathrm{B}_{2}\) to make it isoelectronic with \(\mathrm{Li}_{2}\) so that it might be used in the warp drive? c. The reaction to make \(\mathrm{B}_{2}\) isoelectronic with \(\mathrm{Li}_{2}\) is generalized (where \(n=\) number of electrons determined in part b) as follows: $$ \mathrm{B}_{2} \longrightarrow \mathrm{B}_{2}^{n+}+n \mathrm{e}^{-} \quad \Delta H=6455 \mathrm{~kJ} / \mathrm{mol} $$ How much energy is needed to ionize \(1.5 \mathrm{~kg} \mathrm{~B}_{2}\) to the desired isoelectronic species?

Explain the difference between the \(\sigma\) and \(\pi\) MOs for homonuclear diatomic molecules. How are bonding and antibonding orbitals different? Why are there two \(\pi\) MOs and one \(\sigma\) MO? Why are the \(\pi\) MOs degenerate?

In Exercise 89 in Chapter 8, the Lewis structures for benzene \(\left(\mathrm{C}_{6} \mathrm{H}_{6}\right)\) were drawn. Using one of the Lewis structures, estimate \(\Delta H_{\mathrm{f}}^{\circ}\) for \(\mathrm{C}_{6} \mathrm{H}_{6}(g)\) using bond energies and given that the standard enthalpy of formation of \(\mathrm{C}(g)\) is \(717 \mathrm{~kJ} / \mathrm{mol}\). The experimental \(\Delta H_{\mathrm{f}}^{\circ}\) value of \(\mathrm{C}_{6} \mathrm{H}_{6}(g)\) is \(83 \mathrm{~kJ} / \mathrm{mol} .\) Explain the discrepancy between the experimental value and the calculated \(\Delta H_{\mathrm{f}}^{\circ}\) value for \(\mathrm{C}_{6} \mathrm{H}_{6}(g)\)

Consider three molecules: \(\mathrm{A}, \mathrm{B}\), and \(\mathrm{C}\). Molecule A has a hybridization of \(s p^{3}\). Molecule B has two more effective pairs (electron pairs around the central atom) than molecule A. Molecule C consists of two \(\sigma\) bonds and two \(\pi\) bonds. Give the molecular structure, hybridization, bond angles, and an example for each molecule.

Describe the bonding in the first excited state of \(\mathrm{N}_{2}\) (the one closest in energy to the ground state) using the molecular orbital model. What differences do you expect in the properties of the molecule in the ground state as compared to the first excited state? (An excited state of a molecule corresponds to an electron arrangement other than that giving the lowest possible energy.)
See all solutions

Recommended explanations on Chemistry Textbooks

Chemical Reactions

Read Explanation

Chemical Analysis

Read Explanation

Nuclear Chemistry

Read Explanation

Inorganic Chemistry

Read Explanation

Ionic and Molecular Compounds

Read Explanation

Kinetics

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.