1. Molecule: CF4indicate the number ofavailable electrons thatelectronsae =are in the molecule.in the space to the rightconnect all of the atomsto the central atom andthen make each atomfollow the octet ruleTrial Structure:(duet rule for hydrogen).How many electrons arenecessary in the trialne =structure?Circle the correctne = aene < aene>aerelationship between neand ae.Draw the correctedLewis Structure to theright.Add Later:e- geometry:molecular geOmHybridization: 1. Molecule: CF4indicate the number ofavailable electrons thatelectronsae =are in the molecule.in the space to the rightconnect all of the atomsto the central atom andthen make each atomfollow the octet ruleTrial Structure:(duet rule for hydrogen).How many electrons arenecessary in the trialne =structure?Circle the correctne = aene < aene>aerelationship between neand ae.Draw the correctedLewis Structure to theright.Add Later:e- geometry:molecular geOmHybridization:

The given molecule is CF4 named as carbon tetrafluoride, the number of available electrons and the…

1. Molecule: CF4 indicate the number of Jelectrons available electrons that are in the molecule. in the space to the right connect all of the atoms to the central atom and then make each atom follow the octet rule ae = Trial Structure: (duet rule for hydrogen). How many electrons are necessary in the trial structure? he = Circle the correct ne = ae ne < ae ne>ae relationship between ne and ae. Draw the corrected Lewis Structure to the right. Add Later: e- geometry: molecular gesom Hybridization:

1. Molecule: CF4 indicate the number of Jelectrons available electrons that are in the molecule. in the space to the right connect all of the atoms to the central atom and then make each atom follow the octet rule ae = Trial Structure: (duet rule for hydrogen). How many electrons are necessary in the trial structure? he = Circle the correct ne = ae ne < ae ne>ae relationship between ne and ae. Draw the corrected Lewis Structure to the right. Add Later: e- geometry: molecular gesom Hybridization:

Principles of Modern Chemistry

8th Edition

ISBN:9781305079113

Author:David W. Oxtoby, H. Pat Gillis, Laurie J. Butler

Publisher:David W. Oxtoby, H. Pat Gillis, Laurie J. Butler

Chapter3: Atomic Shells And Classical Models Of Chemical Bonding

Section: Chapter Questions

Problem 67P

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A stable triatomic molecule can be formed that contains one atom each of nitrogen, sulfur, and fluorine. Three bonding structures are possible, depending on which is the central atom: NSF, SNF, and SFN. (a) Write a Lewis diagram for each of these molecules, indicating the formal charge on each atom. (b) Often, the structure with the least separation of formal charge is the most stable. Is this statement consistent with the observed structure for this molecule—namely, NSF, which has a central sulfur atom? (c) Does consideration of the electronegativities of N, S, and F from Figure 3.18 help rationalize this observed structure? Explain.
The percent ionic character of a bond can be approximated by the formula 16+3.52 , where is the magnitude of the difference in the electronegativities of the atoms (see Fig. 3.18). Calculate the percent ionic character of HF, HCl, HBr, HI, and CsF, and compare the results with those in Table 3.7.
What is meant by a chemical bond? Why do atoms form bonds with each other? Why do some elements exist as molecules in nature instead of as free atoms?
The structure of pyridine is When a proton becomes bonded to the nitrogen atom by way of its unshared electron pair, the result is _____________________________
7.30 The bond in HF is said to be polar, with the hydrogen carrying a partial positive charge. For this to be true, the hydrogen atom must have less than one electron around it. Yet the Lewis dot structure of HF attributes two electrons to hydrogen. Draw a picture of the electron density distribution for HF and use it to describe how the hydrogen atom can carry a partial positive charge. How can these two models of the HF bond (the electron density and the Lewis structure) seem so different and yet describe the same thing?
It is impossible to draw a legitimate Lewis structure of a neutral NH4 molecule. Hypothetically,how many valence electrons would such a neutral NH4 molecule have ifit could exist? a. The +1 cation, NH4+ , does exist. How many valence electrons does one NH4+ ion have? b. Draw the Lewis structure for NH4+
For each of the following molecules, complete the Lewis structure and use the VSEPR model to determine the bond angles around each central atom. Note that the drawings are only skeleton structures and may depict the angles incorrectly.
A compound is being tested for use as a rocket propellant. Analysis shows that it contains 18.54% F, 34.61% Cl, and 46.85% O. (a) Determine the empirical formula for this compound. (b) Assuming that the molecular formula is the same as the empirical formula, draw a Lewis diagram for this molecule. Review examples elsewhere in this chapter to decide which atom is most likely to lie at the center. (c) Use the VSEPR theory to predict the structure of the molecule from part (b).
Fill in the blanks in each line of the following table that involves characteristics of various bonds between nonmetals. The first line is already completed as an example.
7.84 Which of the following molecules is least likely to actually exist OF4,SF4,SeF4 , or TeF4 ? Why?
Use the data in Table 3.1 to plot the logarithm of ionization energy versus the number of electrons removed for Be. Describe the electronic structure of the Be atom.
Define formal charge and explain how to calculate it. What is the purpose of the formal charge? Organic compounds are composed mostly of carbon and hydrogen but also may have oxygen, nitrogen, and/or halogens in the formula. Formal charge arguments work very well for organic compounds when drawing the best Lewis structure. How do C, H, N, O, and Cl satisfy the octet rule in organic compounds so as to have a formula charge of zero?