Chapter 8

Does the number of valence electrons in a neutral atom ever equal the atomic number?

Does the number of valence electrons in a neutral atom ever equal the group number?

Do all the elements in a group in the periodic table hav the same number of valence electrons?

Distinguish between an atom's valence electrons and its total electron count.

Some of his critics described G. N. Lewis's approach to explaining covalent bonding as an exercise in double counting and therefore invalid. Explain the basis for this criticism.

Does the octet rule mean that a diatomic molecule must have 16 valence electrons?

Why would you not expect to find hydrogen atoms in the bonding arrangement \(\mathrm{X}-\mathrm{H}-\mathrm{X} ?\)

Does each atom in a pair that is covalently bonded always contribute the same number of valence electrons to form the bonds between them?

Draw Lewis symbols of atoms of lithium, magnesium, and aluminum.

Draw Lewis symbols of atoms of nitrogen, oxygen, fluorine, and chlorine.

Draw Lewis symbols for \(\operatorname{In}^{+}, \mathrm{I}^{-}, \mathrm{Ca}^{2+},\) and \(\mathrm{Sn}^{2+} .\) Which ions have a complete valence-shell octet?

Draw Lewis symbols of \(\mathrm{Xe}, \mathrm{Sr}^{2+}, \mathrm{Cl},\) and \(\mathrm{Cl}^{-} .\) How many valence electrons are in each atom or ion?

Draw the Lewis symbol of an ion that has the following: a. \(1+\) charge and 1 valence electron b. \(3+\) charge and 0 valence electrons

Draw the Lewis symbol of an ion that has the following: a. \(1-\) charge and 8 valence electrons b. \(1+\) charge and 5 valence electrons

How many valence electrons does each of the following species contain? (a) \(\mathrm{BN} ;\) (b) \(\mathrm{HF} ;\) (c) \(\mathrm{OH}^{-} ;\) (d) \(\mathrm{CN}^{-}\).

How many valence electrons does each of the following species contain? (a) \(\mathrm{N}_{2}^{+} ;\) (b) \(\mathrm{CS}^{+} ;\) (c) \(\mathrm{CN} ;\) (d) CO.

Draw Lewis structures for the following diatomic molecules and ions: (a) \(\mathrm{CO}_{i}\) (b) \(\mathrm{O}_{2}\); (c) \(\mathrm{ClO}^{-}\); (d) \(\mathrm{CN}^{-}\).

Draw Lewis structures for the following diatomic molecules and ions: (a) \(\mathrm{F}_{2} ;\) (b) \(\mathrm{NO}^{+} ;\) (c) \(\mathrm{SO} ;\) (d) HI.

Which of the groups in the periodic table will carry negative partial charges in diatomic compounds with hydrogen, HX and \(\mathrm{H}_{2} \mathrm{X} ?\)

Greenhouse Gases Chlorofluorocarbons (CFCs) are linked to the depletion of stratospheric ozone. They are also greenhouse gases. Draw Lewis structures for the following CFCs: a. \(\left.\mathrm{CF}_{2} \mathrm{Cl}_{2} \text { (Freon } 12\right)\) b. \(\left.\mathrm{Cl}_{2} \mathrm{FCCF}_{2} \mathrm{Cl} \text { (Freon } 113, \text { containing a } \mathrm{C}-\mathrm{C} \text { bond }\right)\) c. \(\mathrm{C}_{2} \mathrm{ClF}_{3}\) (Freon 1113 , containing a \(\mathrm{C}=\mathrm{C}\) bond).

Skunks and Rotten Eggs Many sulfur-containing organic compounds have characteristically foul odors: butanethiol \(\left(\mathrm{CH}_{3} \mathrm{CH}_{2} \mathrm{CH}_{2} \mathrm{CH}_{2} \mathrm{SH}\right)\) is responsible for the odor of skunks, and rotten eggs smell the way they do because they produce tiny amounts of pungent hydrogen sulfide, \(\mathrm{H}_{2} \mathrm{S} .\) Draw the Lewis structures for \(\mathrm{CH}_{3} \mathrm{CH}_{2} \mathrm{CH}_{2} \mathrm{CH}_{2} \mathrm{SH}\) and \(\mathrm{H}_{2} \mathrm{S}\).

Chlorine Bleach Chlorine combines with oxygen in several proportions. Dichlorine monoxide (Cl \(_{2} \mathrm{O}\) ) is used in the manufacture of bleaching agents. Potassium chlorate \(\left(\mathrm{KClO}_{3}\right)\) is used in oxygen generators aboard aircraft. Draw the Lewis structures for \(\mathrm{Cl}_{2} \mathrm{O}\) and \(\mathrm{ClO}_{3}^{-}\).

How can we use electronegativity to predict whether a bond between two atoms is likely to be covalent or ionic?

How do the electronegativities of the elements change across a period and down a group?

Explain on the basis of atomic structure why trends in electronegativity are related to trends in atomic size.

Is the element with the most valence electrons in a period also the most electronegative? Explain.

What is meant by the term polar covalent bond?

What factor is responsible for the existence of polar covalent bonds?

Describe how atmospheric greenhouse gases act like the panes of glass in a greenhouse.

Understanding the Greenhouse Effect Water vapor in the atmosphere contributes more to the greenhouse effect than carbon dioxide, yet water vapor is not considered an important factor in global warming. Propose a reason why.

Increasing concentrations of nitrous oxide in the atmosphere may be contributing to climate change. Is the ability of \(\mathrm{N}_{2} \mathrm{O}\) to absorb infrared radiation due to nitrogen-nitrogen bond stretching, nitrogen-oxygen bond stretching, or both? Explain your answer.

Is the ability of \(\mathrm{H}_{2} \mathrm{O}\) molecules to absorb photons of infrared radiation due to symmetrical stretching or asymmetrical stretching of its \(\mathrm{O}-\mathrm{H}\) bonds, or both? Explain your answer. (Hint: The angle between the two \(\mathrm{O}-\mathrm{H}\) bonds in \(\mathrm{H}_{2} \mathrm{O}\) is \(104.5^{\circ} .\)

Which of the following bonds are polar: \(\mathrm{C}-\mathrm{Se}, \mathrm{C}-\mathrm{O}\) \(\mathrm{Cl}-\mathrm{Cl}, \mathrm{O}=\mathrm{O}, \mathrm{N}-\mathrm{H}, \mathrm{C}-\mathrm{H} ?\) In the bond or bonds that you selected, which atom has the greater electronegativity?

Which is the least polar bond: \(C-S e, C=O, C 1-B r\) \(\mathrm{O}=\mathrm{O}, \mathrm{N}-\mathrm{H}, \mathrm{C}-\mathrm{H} ?\)

In which of the following binary compounds is the bond expected to have the least ionic character? LiCl; Cs I; \(\mathrm{KBr} ; \mathrm{NaF}\).

In which of the following compounds is the bond between the atoms expected to have the most covalent character? \(\mathrm{AlCl}_{3} ; \mathrm{AlBr}_{3} ; \mathrm{AlI}_{3} ; \mathrm{GaF}_{3}\).

Which substance has the most polar covalent bonds: \(\mathrm{PF}_{3}\) \(\mathrm{S}_{8}, \mathrm{RbCl},\) or \(\mathrm{SF}_{2} ?\)

Atoms of which element are held together by nonpolar covalent bonds: lithium, phosphorus, or xenon?

Explain the concept of resonance.

How does resonance influence the stability of a molecule or an ion?

What factors determine whether a molecule or ion exhibits resonance?

What structural features do all the resonance forms of a molecule or ion have in common?

Explain why \(\mathrm{NO}_{2}\) is more likely to exhibit resonance than \(\mathrm{CO}_{2}\).

Are these two skeletal structures resonance forms: \(\mathrm{X}-\mathrm{X}-\mathrm{O}\) and \(\mathrm{X}-\mathrm{O}-\mathrm{X} ?\) Explain.

Draw Lewis structures for fulminic acid (HCNO), showing all resonance forms.

Draw Lewis structures for hydrazoic acid (HN \(_{3}\) ), showing all resonance forms.

Oxygen and nitrogen combine to form a variety of nitrogen oxides, including the following two unstable compounds, each with two nitrogen atoms per molecule: \(\mathrm{N}_{2} \mathrm{O}_{2}\) and \(\mathrm{N}_{2} \mathrm{O}_{3} .\) Draw Lewis structures for these molecules, showing all resonance forms.

Oxygen and sulfur combine to form a variety of sulfur oxides. Some are stable molecules and some, including \(\mathrm{S}_{2} \mathrm{O}_{2}\) and \(\mathrm{S}_{2} \mathrm{O}_{3},\) decompose when they are heated. Draw Lewis structures for these two compounds, showing all resonance forms.

The oxygen-oxygen distance in \(\mathrm{F}_{2} \mathrm{O}_{2}\) is about \(20 \%\) shorter than in hydrogen peroxide. a. Draw Lewis structures for both \(\mathrm{F}_{2} \mathrm{O}_{2}\) and \(\mathrm{H}_{2} \mathrm{O}_{2}\) b. It was proposed that \(\mathrm{F}_{2} \mathrm{O}_{2}\) has a resonance form of \(\left[\mathrm{FO}_{2}\right]^{+} \mathrm{F}^{-} .\) Draw a Lewis structure for this "ionic" form of \(\mathrm{F}_{2} \mathrm{O}_{2}\) and explain how the structure is consistent with the observed O-O distance.

The nitrogen-oxygen bond distance in \(\mathrm{NOF}_{3}\) is shorter than in \(\mathrm{NO}\left(\mathrm{CH}_{3}\right)_{3}\) a. Draw Lewis structures for \(\mathrm{NOF}_{3}\) and \(\mathrm{NO}\left(\mathrm{CH}_{3}\right)_{3}\) b. It was proposed that \(\mathrm{NOF}_{3}\) has a resonance form of \(\left[\mathrm{NOF}_{2}\right]^{+} \mathrm{F}^{-} .\) Draw a Lewis structure for this "ionic" form of \(\mathrm{NOF}_{3}\) and explain how the structure is consistent with the observed \(\mathrm{N}\) - \(\mathrm{O}\) distance.