Log out Go to App
Go to App

Learning Materials

Features

Discover

Chapter 10: Problem 105

Draw three Lewis structures for compounds with the formula \(\mathrm{C}_{2} \mathrm{H}_{2} \mathrm{~F}_{2} .\) Indicate which of the compound(s) are polar.

Short Answer

Expert verified
The three possible Lewis structures for C2H2F2 can have the F and H either on the same C atom or on different ones, with all atoms having a formal charge of zero. Out of the three structures, only the one where each C atom is bonded to one H atom and one F atom results in a polar molecule, the other two are nonpolar as the dipoles cancel out.

Step by step solution

01

Identify the central atom

Identify the central atom in the molecule. In this case, there are two carbon atoms \(C\) and they are the central atoms.
02

Calculate valence electrons

Determine the total number of valence electrons. Carbon (C) has 4 valence electrons, Hydrogen (H) has 1 valence electron, and Fluorine (F) has 7 valence electrons. Hence, total valence electrons = 2*4(C) + 2*1(H) + 2*7(F) = 24 valence electrons.
03

Draw base structure

Draw a base structure that shows how the atoms are bonded.In the first structure, Carbon atoms are in the middle with two fluorines bonded to one carbon and two hydrogens bonded to the other carbon. In the second structure, each carbon is bonded to one hydrogen and one fluorine. In the third structure, there's alternating single and double bonds between carbon atoms, and each carbon atom is bonded to one hydrogen and one fluorine.
04

Draw Lone pairs

Place the remaining valence electrons as lone pairs on the atoms. Each fluorine atom will have 6 non-bonding electrons or 3 lone pairs after forming a bond.
05

Check formal charges

Check formal charges to ensure you have the best Lewis structure. Formal charge is the charge assigned to an atom in a molecule, assuming that electrons in all chemical bonds are shared equally between atoms, regardless of relative electronegativity. Each structure has a formal charge of zero, which is what we want for the most stable structure.
06

Determine polarity

A molecule is polar if the bond dipoles do not cancel each other out. In the first structure, the molecule is nonpolar as the bond dipoles cancel. In the second structure, the molecule is polar as the bond dipoles do not cancel. The third structure is nonpolar as the bond dipoles cancel.

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.

Start your free trial

Over 30 million students worldwide already upgrade their learning with Vaia!

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

Given that the order of molecular orbitals for \(\mathrm{NO}\) is similar to that for \(\mathrm{O}_{2}\), arrange the following species in increasing bond orders: \(\mathrm{NO}^{2-}, \mathrm{NO}^{-}, \mathrm{NO}, \mathrm{NO}^{+}\) \(\mathrm{NO}^{2+}\)

Which of the following molecules and ions are linear? \(\mathrm{ICl}_{2}^{-}, \mathrm{IF}_{2}^{+}, \mathrm{OF}_{2}, \mathrm{SnI}_{2}, \mathrm{CdBr}_{2}\)

Carbon monoxide (CO) is a poisonous compound due to its ability to bind strongly to \(\mathrm{Fe}^{2+}\) in the hemoglobin molecule. The molecular orbitals of CO have the same energy order as those of the \(\mathrm{N}_{2} \mathrm{~mol}-\) ecule. (a) Draw a Lewis structure of \(\mathrm{CO}\) and assign formal charges. Explain why CO has a rather small dipole moment of 0.12 D. (b) Compare the bond order of CO with that from molecular orbital theory. (c) Which of the atoms (C or O) is more likely to form bonds with the \(\mathrm{Fe}^{2+}\) ion in hemoglobin?

For each pair listed here, state which one has a higher first ionization energy and explain your choice: (a) \(\mathrm{H}\) or \(\mathrm{H}_{2},\) (b) \(\mathrm{N}\) or \(\mathrm{N}_{2},\) (c) \(\mathrm{O}\) or \(\mathrm{O}_{2},\) (d) \(\mathrm{F}\) or \(\mathrm{F}_{2}\).

What is the state of hybridization of the central \(\mathrm{O}\) atom in \(\mathrm{O}_{3} ?\) Describe the bonding in \(\mathrm{O}_{3}\) in terms of delocalized molecular orbitals.
See all solutions

Recommended explanations on Chemistry Textbooks

Kinetics

Read Explanation

Physical Chemistry

Read Explanation

Making Measurements

Read Explanation

Ionic and Molecular Compounds

Read Explanation

Organic Chemistry

Read Explanation

Chemical Reactions

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.

Sign-up for free