Chapter 8

(a) What are valence electrons? (b) How many valence electrons does a nitrogen atom possess? (c) An atom has the electron configuration \(1 s^{2} 2 s^{2} 2 p^{6} 3 s^{2} 3 p^{2} .\) How many valence electrons does the atom have?

(a) What is the octet rule? (b) How many electrons must a sulfur atom gain to achieve an octet in its valence shell? (c) If an atom has the electron configuration \(1 s^{2} 2 s^{2} 2 p^{3}\), how many electrons must it gain to achieve an octet?

Write the electron configuration for phosphorus. Identify the valence electrons in this configuration and the nonvalence electrons. From the standpoint of chemical reactivity, what is the important difference between them?

Write the Lewis symbol for atoms of each of the following elements: (a) \(\mathrm{Al}\), (b) \(\mathrm{Br}\), (c) \(\mathrm{Ar}\), (d) \(\mathrm{Sr}\).

What is the Lewis symbol for each of the following atoms or ions: (a) \(\mathrm{Ca}\), (b) \(\mathrm{P}_{,}(\mathrm{c}) \mathrm{Mg}^{2+}\), (d) \(\mathrm{S}^{2-}\) ?

Using Lewis symbols, diagram the reaction between magnesium and oxygen atoms to give the ionic substance \(\mathrm{MgO}\).

Use Lewis symbols to represent the reaction that occurs between \(\mathrm{Ca}\) and \(\mathrm{F}\) atoms.

Predict the chemical formula of the ionic compound formed between the following pairs of elements: (a) \(\mathrm{Al}\) and \(\mathrm{F}\), (b) \(\mathrm{K}\) and \(\mathrm{S}\), (c) \(\mathrm{Y}\) and \(\mathrm{O}\), (d) \(\mathrm{Mg}\) and \(\mathrm{N}\).

Which ionic compound is expected to form from combining the following pairs of elements: (a) barium and fluorine, (b) cesium and chlorine, (c) lithium and nitrogen, (d) aluminum and oxygen?

Write the electron configuration for each of the following ions, and determine which ones possess noble-gas configurations: (a) \(\mathrm{Sr}^{2+}\), (b) \(\mathrm{Ti}^{2+}\), (c) \(\mathrm{Se}^{2-}\), (d) \(\mathrm{Ni}^{2+}\), (e) \(\mathrm{Br}^{-}\), (f) \(\mathrm{Mn}^{3+}\).

Write electron configurations for the following ions, and determine which have noble-gas configurations: (a) \(\mathrm{Zn}^{2+}\), (b) \(\mathrm{Te}^{2-}\) (c) \(\mathrm{Sc}^{3+}\), (d) \(\mathrm{Rh}^{3+}\), (e) \(\mathrm{Tl}^{+}\), (f) \(\mathrm{Bi}^{3+}\).

(a) Define the term lattice energy. (b) Which factors govern the magnitude of the lattice energy of an ionic compound?

(a) Does the lattice energy of an ionic solid increase or decrease (i) as the charges of the ions increase, (ii) as the sizes of the ions increase? (b) Using a periodic table, arrange the following substances according to their expected lattice energies, listing them from lowest lattice energy to the highest: \(\mathrm{ScN}, \mathrm{KBr}, \mathrm{MgO}, \mathrm{NaF}\). Compare your list with the data in Table \(8.2\).

The lattice energies of \(\mathrm{KBr}\) and \(\mathrm{CsCl}\) are nearly equal (Table 8.2). What can you conclude from this observation?

Explain the following trends in lattice energy: (a) \(\mathrm{CaF}_{2}>\mathrm{BaF}_{2} ;\) (b) \(\mathrm{NaCl}>\mathrm{RbBr}>\mathrm{CsBr} ;\) (c) \(\mathrm{BaO}>\mathrm{KF}\)

Energy is required to remove two electrons from Ca to form \(\mathrm{Ca}^{2+}\) and is required to add two electrons to \(\mathrm{O}\) to form \(\mathrm{O}^{2-}\). Why, then, is \(\mathrm{CaO}\) stable relative to the free elements?

(a) Based on the lattice energies of \(\mathrm{MgCl}_{2}\) and \(\mathrm{SrCl}_{2}\) given in Table \(8.2\), what is the range of values that you would expect for the lattice energy of \(\mathrm{CaCl}_{2} ?\) (b) Using data from Appendix \(C\), Figure \(7.12\), and Figure \(7.14\) and the value of the second ionization energy for \(\mathrm{Ca}\), \(1145 \mathrm{~kJ} / \mathrm{mol}\), calculate the lattice energy of \(\mathrm{CaCl}_{2} .\)

(a) What is meant by the term covalent bond? (b) Give three examples of covalent bonding. (c) A substance XY, formed from two different elements, boils at \(-33^{\circ} \mathrm{C}\). Is XY likely to be a covalent or an ionic substance? Explain.

Which of these elements is unlikely to form covalent bonds: \(\mathrm{S}, \mathrm{H}, \mathrm{K}, \mathrm{Ar}, \mathrm{Si}\) ? Explain your choices.

Using Lewis symbols and Lewis structures, diagram the formation of \(\mathrm{SiCl}_{4}\) from \(\mathrm{Si}\) and \(\mathrm{Cl}\) atoms.

Use Lewissymbols and Lewis structures to diagram the formation of \(\mathrm{PF}_{3}\) from \(\mathrm{P}\) and \(\mathrm{F}\) atoms.

(a) Construct a Lewis structure for \(\mathrm{O}_{2}\) in which each atom achieves an octet of electrons. (b) Explain why it is necessary to form a double bond in the Lewis structure. (c) The bond in \(\mathrm{O}_{2}\) is shorter than the \(\mathrm{O}-\mathrm{O}\) bond in compounds that contain an \(\mathrm{O}-\mathrm{O}\) single bond. Explain this observation.

(a) Construct a Lewis structure for hydrogen peroxide, \(\mathrm{H}_{2} \mathrm{O}_{2}\), in which each atom achieves an octet of electrons. (b) Do you expect the \(\mathrm{O}-\mathrm{O}\) bond in \(\mathrm{H}_{2} \mathrm{O}_{2}\) to be longer or shorter than the \(\mathrm{O}-\mathrm{O}\) bond in \(\mathrm{O}_{2}\) ?

(a) What is meant by the term electronegativity? (b) On the Pauling scale what is the range of electronegativity values for the elements? (c) Which element has the greatest electronegativity? (d) Which element has the smallest electronegativity?

(a) What is the trend in electronegativity going from left to right in a row of the periodic table? (b) How do electronegativity values generally vary going down a column in the periodic table? (c) How do periodic trends in electronegativity relate to those for ionization energy and electron affinity?

Using only the periodic table as your guide, select the most electronegative atom in each of the following sets: (a) Se, \(\mathrm{Rb}, \mathrm{O}, \mathrm{In} ;\) (b) \(\mathrm{Al}, \mathrm{Ca}, \mathrm{C}, \mathrm{Si} ;\) (c) \(\mathrm{Ge}, \mathrm{As}, \mathrm{P}, \mathrm{Sn} ;\) (d) \(\mathrm{Li}\), \(\mathrm{Rb}, \mathrm{Be}, \mathrm{Sr}\)

By referring only to the periodic table, select (a) the most electronegative element in group \(6 \mathrm{~A} ;\) (b) the least electronegative element in the group \(\mathrm{Al}, \mathrm{Si}, \mathrm{P} ;\) (c) the most electronegative element in the group Ga, \(\mathrm{P}, \mathrm{Cl}, \mathrm{Na}\); (d) the element in the group \(\mathrm{K}, \mathrm{C}, \mathrm{Zn}, \mathrm{F}\), that is most likely to form an ionic compound with Ba.

Which of the following bonds are polar: (a) \(\mathrm{B}-\mathrm{F}\), (b) \(\mathrm{Cl}-\mathrm{Cl}\), (c) \(\mathrm{Se}-\mathrm{O}\), (d) \(\mathrm{H}-\mathrm{I}\) ? Which is the more electronegative atom in each polar bond?

Arrange the bonds in each of the following sets in order of increasing polarity: (a) \(\mathrm{C}-\mathrm{F}, \mathrm{O}-\mathrm{F}, \mathrm{Be}-\mathrm{F}\) (b) \(\mathrm{O}-\mathrm{Cl}, \mathrm{S}-\mathrm{Br}, \mathrm{C}-\mathrm{P}\) (c) \(\mathrm{C}-\mathrm{S}, \mathrm{B}-\mathrm{F}, \mathrm{N}-\mathrm{O}\)

The dipole moment and bond distance measured for the highly reactive gas phase OH molecule are \(1.78 \mathrm{D}\) and \(0.98 \AA\), respectively. (a) Given these values calculate the effective charges on the \(\mathrm{H}\) and \(\mathrm{O}\) atoms of the OH molecule in units of the electronic charges \(e\). (b) Is this bond more or less polar than the \(\mathrm{H}-\mathrm{Cl}\) bond in an \(\mathrm{HCl}\) molecule? (c) Is that what you would have expected based on electronegativities?

The iodine monobromide molecule, IBr, has a bond length of \(2.49 \AA\) and a dipole moment of \(1.21 \mathrm{D}\). (a) Which atom of the molecule is expected to have a negative charge? Explain. (b) Calculate the effective charges on the I and Br atoms in IBr, in units of the electronic charge \(e\).

In the following pairs of binary compounds determine which one is a molecular substance and which one is an ionic substance. Use the appropriate naming convention (for ionic or molecular substances) to assign a name to each compound: (a) \(\mathrm{SiF}_{4}\) and \(\mathrm{LaF}_{3}\), (b) \(\mathrm{FeCl}_{2}\) and \(\operatorname{ReCl}_{6}\), (c) \(\mathrm{PbCl}_{4}\) and \(\mathrm{RbCl}\).

Draw Lewis structures for the following: (a) \(\mathrm{SiH}_{4}\), (b) \(\mathrm{CO}\), (c) \(\mathrm{SF}_{2}\), (d) \(\mathrm{H}_{2} \mathrm{SO}_{4}(\mathrm{H}\) is bonded to \(\mathrm{O})\), (e) \(\mathrm{ClO}_{2}^{-}\) (f) \(\mathrm{NH}_{2} \mathrm{OH}\).

Write Lewis structures for the following: (a) \(\mathrm{H}_{2} \mathrm{CO}\) (both \(\mathrm{H}\) atoms are bonded to \(\mathrm{C}\) ), (b) \(\mathrm{H}_{2} \mathrm{O}_{2}\), (c) \(\mathrm{C}_{2} \mathrm{~F}_{6}\) (contains a \(\mathrm{C}-\mathrm{C}\) bond \(),\) (d) \(\mathrm{AsO}_{3}{\underline{\phantom{xx}}}^{3-}\), (e) \(\mathrm{H}_{2} \mathrm{SO}_{3}(\mathrm{H}\) is bonded to \(\mathrm{O})\), (f) \(\mathrm{C}_{2} \mathrm{H}_{2}\)

(a) When talking about atoms in a Lewis structure, what is meant by the term formal charge? (b) Does the formal charge of an atom represent the actual charge on that atom? Explain. (c) How does the formal charge of an atom in a Lewis structure differ from the oxidation number of the atom?

(a) Write a Lewis structure for the phosphorus trifluoride molecule, \(\mathrm{PF}_{3}\). Is the octet rule satisfied for all the atoms in your structure? (b) Determine the oxidation numbers of the \(\mathrm{P}\) and \(\mathrm{F}\) atoms. (c) Determine the formal charges of the \(\mathrm{P}\) and \(\mathrm{F}\) atoms. (d) Is the oxidation number for the \(\mathrm{P}\) atom the same as its formal charge? Explain why or why not.

Write Lewis structures that obey the octet rule for each of the following, and assign oxidation numbers and formal charges to each atom: (a) \(\mathrm{NO}^{+}\), (b) \(\mathrm{POCl}_{3}\) (P is bonded to the three \(\mathrm{Cl}\) atoms and to the \(\mathrm{O}\), (c) \(\mathrm{ClO}_{4}^{-}\), (d) \(\mathrm{HClO}_{3}(\mathrm{H}\) is bonded to \(\mathrm{O})\).

(a) Write one or more appropriate Lewis structures for the nitrite ion, \(\mathrm{NO}_{2}^{-} .\) (b) With what allotrope of oxygen is it isoelectronic? (c) What would you predict for the lengths of the bonds in \(\mathrm{NO}_{2}^{-}\) relative to \(\mathrm{N}-\mathrm{O}\) single bonds?

Consider the nitryl cation, \(\mathrm{NO}_{2}^{+} .\) (a) Write one or more appropriate Lewis structures for this ion. (b) Are resonance structures needed to describe the structure? (c) With what familiar molecule is it isoelectronic?

Predict the ordering of the \(\mathrm{C}-\mathrm{O}\) bond lengths in \(\mathrm{CO}\), \(\mathrm{CO}_{2}\), and \(\mathrm{CO}_{3}^{2-}\)

Based on Lewis structures, predict the ordering of \(\mathrm{N}-\mathrm{O}\) bond lengths in \(\mathrm{NO}^{+}, \mathrm{NO}_{2}^{-}\), and \(\mathrm{NO}_{3}^{-} .\)

(a) Use the concept of resonance to explain why all six \(\mathrm{C}-\mathrm{C}\) bonds in benzene are equal in length. (b) Are the \(\mathrm{C}-\mathrm{C}\) bond lengths in benzene shorter than \(\mathrm{C}-\mathrm{C} \mathrm{sin}-\) gle bonds? Are they shorter than \(\mathrm{C}=\mathrm{C}\) double bonds?

(a) State the octet rule. (b) Does the octet rule apply to ionic as well as to covalent compounds? Explain, using examples as appropriate.

Considering the nonmetals, what is the relationship between the group number for an element (carbon, for example, belongs to group \(4 \mathrm{~A}\); see the periodic table on the inside front cover) and the number of single covalent bonds that element needs to form to conform to the octet rule?

What is the most common exception to the octet rule? Give two examples.

For elements in the third row of the periodic table and beyond, the octet rule is often not obeyed. What factors are usually cited to explain this fact?

Draw the Lewis structures for each of the following ions or molecules. Identify those that do not obey the octet rule, and explain why they do not. (a) \(\mathrm{SO}_{3}{\underline{\phantom{xx}}}^{2-}\), (b) \(\mathrm{AlH}_{3}\), (c) \(\mathrm{N}_{3}^{-}\), (d) \(\mathrm{CH}_{2} \mathrm{Cl}_{2}\), (e) \(\mathrm{SbF}_{5}\).

Draw the Lewis structures for each of the following molecules or ions. Which do not obey the octet rule? (b) \(\mathrm{SCN}^{-}\), (a) \(\mathrm{NH}_{4}^{+}\), (c) \(\mathrm{PCl}_{3}\), (d) \(\mathrm{TeF}_{4}\), (e) \(\mathrm{XeF}_{2}\).

In the vapor phase, \(\mathrm{BeCl}_{2}\) exists as a discrete molecule. (a) Draw the Lewis structure of this molecule, using only single bonds. Does this Lewis structure satisfy the octet rule? (b) What other resonance forms are possible that satisfy the octet rule? (c) Using formal charges, select the resonance form from among all the Lewis structures that is most important in describing \(\mathrm{BeCl}_{2}\).

(a) Describe the molecule chlorine dioxide, \(\mathrm{ClO}_{2}\), using three possible resonance structures. (b) Do any of these resonance structures satisfy the octet rule for every atom in the molecule? Why or why not? (c) Using formal charges, select the resonance structure(s) that is (are) most important.