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Chapter 10: Problem 1

Write Lewis symbols for the following atoms. (a) \(\mathrm{Kr}\); (b) Ge; (c) \(\mathrm{N} ;\) (d) Ga; (e) As; (f) Rb.

Short Answer

Expert verified
(a) \(\mathrm{Kr}: . . . . . . . .\) (eight dots evenly spread and paired around Kr symbol). (b) \(\mathrm{Ge}: . . . .\) (four dots around Ge symbol forming a square). (c) \(\mathrm{N}: . . . . .\) (three paired and two unpaired dots around N). (d) \(\mathrm{Ga}: . . .\) (three unpaired dots around Ga). (e) \(\mathrm{As}: . . . . . \) (three paired and two unpaired dots around As). (f) \(\mathrm{Rb}: .\) (one dot next to Rb).

Step by step solution

01

Understand Lewis Symbols

Lewis symbols or Lewis dot diagrams are diagrams that show the bonding between atoms of a molecule, and the lone pairs of electrons that may exist in the molecule. Each dot represents a valence electron, each line represents a bonded pair of electrons and each pair of dots or line represents a single electron pair.
02

Identify atomic numbers

Identify the atomic numbers respective to given atoms. (a) Kr (krypton) has an atomic number of 36. (b) Ge (germanium) has an atomic number of 32. (c) N (nitrogen) has an atomic number of 7. (d) Ga (gallium) has an atomic number of 31. (e) As (arsenic) has an atomic number of 33. (f) Rb (rubidium) has an atomic number of 37.
03

Identify number of valence electrons

Based on atomic number determine the number of valence electrons. Kr has 8, Ge has 4, N has 5, Ga has 3, As has 5, and Rb has 1 valence electron.
04

Draw Lewis symbols

Draw Lewis symbols. Since each dot represents a valence electron, draw corresponding dots around the element's symbol to represent its valence electrons. Arrange the first four dots around the symbol to denote one electron in each of the four sides (up, right, down, left), then start pairing until the number of dots equal the number of valence electrons for that atom.

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Most popular questions from this chapter

Explain the important distinctions between (a) ionic and covalent bonds; (b) lone-pair and bond-pair electrons; (c) molecular geometry and electron-group geometry; (d) bond dipole and resultant dipole moment; (e) polar molecule and nonpolar molecule.

Which of the following molecules would you expect to be polar: (a) \(\mathrm{HCN} ;\) (b) \(\mathrm{SO}_{3} ;\) (c) \(\mathrm{CS}_{2} ;\) (d) OCS; (e) \(\operatorname{SOCl}_{2} ;\) (f) \(\operatorname{SiF}_{4} ;\) (g) \(\operatorname{POF}_{3}\) ? Give reasons for your conclusions.

What is the formal charge of the indicated atom in each of the following structures? (a) the central \(\mathrm{O}\) atom in \(\mathrm{O}_{3}\) (b) \(\mathrm{Al}\) in \(\mathrm{AlH}_{4}^{-}\) (c) \(\mathrm{Cl}\) in \(\mathrm{ClO}_{3}^{-}\) (d) \(\sin \sin ^{2} \theta^{-}\) (e) \(\mathrm{Cl}\) in \(\mathrm{Cl} \mathrm{F}_{3}\)

Assign formal charges to the atoms in the following species, and then select the more likely skeletal structure. (a) \(\mathrm{H}_{2} \mathrm{NOH}\) or \(\mathrm{H}_{2} \mathrm{ONH}\) (b) SCS or CSS (c) NFO or FNO (d) \(\mathrm{SOCl}_{2}\) or \(\mathrm{OSCl}_{2}\) or \(\mathrm{OCl}_{2} \mathrm{S}\) (e) \(\mathrm{F}_{3} \mathrm{SN}\) and \(\mathrm{F}_{3} \mathrm{NS}\)

Which of the following molecules would you expect to have a resultant dipole moment \((\mu) ;\) (a) \(\mathrm{F}_{2}\), (b) \(\mathrm{NO}_{2}\) (c) \(\mathrm{BF}_{3},\) (d) \(\mathrm{HBr},\) (e) \(\mathrm{H}_{2} \mathrm{CCl}_{2}\), (f) \(\mathrm{SiF}_{4},\) (g) OCS? Explain.
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