:ö: A. The Lewis diagram for BrO, is: :ö-Br-0: :0: The electron-pair geometry around the Br atom in BrO4 is There are lone pair(s) around the central atom, so the geometry of BrO4 is B. The Lewis diagram for SeH2 is: H-Se-H The electron-pair geometry around the Se atom in SeH2 is There are lone pair(s) around the central atom, so the geometry of SeH2 is

:ö: A. The Lewis diagram for BrO, is: :ö-Br-0: :0: The electron-pair geometry around the Br atom in BrO4 is There are lone pair(s) around the central atom, so the geometry of BrO4 is B. The Lewis diagram for SeH2 is: H-Se-H The electron-pair geometry around the Se atom in SeH2 is There are lone pair(s) around the central atom, so the geometry of SeH2 is

Living by Chemistry

2nd Edition

ISBN:9781464142314

Author:Angelica M. Stacy

Publisher:Angelica M. Stacy

ChapterU2: Smells: Molecular Structure And Properties

Section: Chapter Questions

Problem 7STP

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Similar questions

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?
Distinguish between the terms electronegativity versus electron affinity, covalent bond versus ionic bond, and pure covalent bond versus polar covalent bond. Characterize the types of bonds in terms of electronegativity difference. Energetically, why do ionic and covalcnt bonds form?
Given the bonds C N, C H, C Br, and S O, (a) which atom in each is the more electronegative? (b) which of these bonds is the most polar?
Atoms form ions so as to achieve electron configurations similar to those of the noble gases. For the following pairs of noble gas configurations, give the formulas of two simple ionic compounds that would have comparable electron configurations. msp;a.[He]and[Ne]c.[He]and[Ar]b.[Ne]and[Ne]d.[Ne]and[Ar]
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.
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.
Which statement(s) is/are true about bond enthalpy? (a) The bond energy for a triple bond between A and B is three times that of a single bond between A and B. (b) H for the breaking of a bond is always a negative number. (c) Energy is required to make a bond. (d) Bond enthalpy is defined only for bonds broken or formed in the gaseous state. (e) The presence of bonds does not influence the geometry of a molecule. However, the presence of bonds affects the value of the bond enthalpy between two atoms.
rite the Lewis structure for each of the following atoms. msp;a.He(Z=2)d.Ne(Z=10)b.Br(Z=35)e.I(Z=53)c.Sr(Z=38)f.Ra(Z=88)
How does the bond energy of HCl(g) differ from the standard enthalpy of formation of HCl(g)?
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?