Chapter 8

Illustrated are four ions \(-\mathrm{A}, \mathrm{B}, \mathrm{X},\) and \(\mathrm{Y}-\) showing their relative ionic radii. The ions shown in red carry positive charges: \(2+\) charge for \(A\) and a \(1+\) charge for \(B\). Ions shown in blue carry negative charges: a 1 - charge for \(\mathrm{X}\) and a 2 - charge for \(\mathrm{Y}\). (a) Which combinations of these ions produce ionic compounds where there is a 1:1 ratio of cations and anions? (b) Among the combinations in part (a), which leads to the ionic compound having the largest lattice energy? (c) Which combination of ions leads to the ionic compound having the smallest lattice energy? [Section 8.2\(]\)

Incomplete Lewis structures for the nitrous acid molecule, \(\mathrm{HNO}_{2},\) and the nitrite ion, \(\mathrm{NO}_{2}^{-},\) are shown below. (a) Complete each Lewis structure by adding electron pairs as needed. (b) Is the formal charge on \(\mathrm{N}\) the same or different in these two species? (c) Would either \(\mathrm{HNO}_{2}\) or \(\mathrm{NO}_{2}^{-}\) be expected to exhibit resonance? (d) Would you expect the \(\mathrm{N}=\mathrm{O}\) bond in \(\mathrm{HNO}_{2}\) to be longer, shorter, or the same length as the \(\mathrm{N}-\mathrm{O}\) bonds in \(\mathrm{NO}_{2}^{-}\) ? Explain. [Sections 8.5 and \(\left.8.6\right]\)

(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 silicon. Identify the valence electrons in this configuration and the nonvalence electrons. From the standpoint of chemical reactivity, what is the important difference between them?

(a) Write the electron configuration for the element titanium, Ti. How many valence electrons does this atom possess? (b) Hafnium, Hf, is also found in group \(4 \mathrm{~B}\). Write the electron configuration for Hf. (c) Ti and Hf behave as though they possess the same number of valence electrons. Which of the subshells in the electron configuration of Hf behave as valence orbitals? Which behave as core orbitals?

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

What is the Lewis symbol for each of the following atoms or ions: (a) \(K,(b) A s,(c) S n^{2+},\left(\right.\) d) \(N^{3-} ?\)

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 Ca and 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 \(\mathrm{S},(\mathrm{c}) \mathrm{Y}\) 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 configura- tions: (a) \(\mathrm{Sr}^{2+},(\mathbf{b}) \mathrm{Ti}^{2+},(\mathrm{c}) \mathrm{Se}^{2-},(\mathrm{d}) \mathrm{Ni}^{2+},(\mathrm{e}) \mathrm{Br}^{-},\) (f) \(\mathrm{Mn}^{3+}\)

Write electron configurations for the following ions, and determine which have noble-gas configurations: (a) \(\mathrm{Cd}^{2+}\) (b) \(\mathrm{P}^{3-}\) (c) \(Z r^{4+}\) (d) \(\mathrm{Ru}^{3+}\), (e) \(\mathrm{As}^{3-},(\mathrm{f}) \mathrm{Ag}^{+}\)

(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) Arrange the following substances not listed in Table 8.2 according to their expected lattice energies, listing them from lowest lattice energy to the highest: \(\mathrm{MgS}, \mathrm{KI}\), \(\mathrm{GaN}, \mathrm{L} \mathrm{iBr}\)

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

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?

List the individual steps used in constructing a Born-Haber cycle for the formation of \(\mathrm{Bal}_{2}\) from the elements. Which of the steps would you expect to be exothermic?

(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 are unlikely to form covalent bonds: \(\mathrm{S}, \mathrm{H}, \mathrm{K}, \mathrm{Ar},\) 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 Lewis symbols 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?

(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?

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 \(\mathrm{Al}, \mathrm{Si}, \mathrm{P} ;(\mathbf{c})\) the most electronegative element in the group \(\mathrm{Ga}, \mathrm{P}, \mathrm{Cl}, \mathrm{Na} ;\) (d) the element in the group \(\mathrm{K}, \mathrm{C}, \mathrm{Zn}, \mathrm{F}\) that is most likely to form an ionic compound with \(\mathrm{Ba}\).

Which of the following bonds are polar: (a) \(\mathrm{B}-\mathrm{F},\) (b) \(\mathrm{Cl}-\mathrm{Cl}\), $$ \text { (c) } \mathrm{Se}-\mathrm{O} $$ (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) \(\mathrm{O}-\mathrm{Cl}\), \(\mathrm{S}-\mathrm{Br}, \mathrm{C}-\mathrm{P} ;(\mathrm{c}) \mathrm{C}-\mathrm{S}, \mathrm{B}-\mathrm{F}, \mathrm{N}-\mathrm{O}\)

The iodine monobromide molecule, IBr, has a bond length of \(2.49 \mathrm{~A}\) 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 \(\mathrm{I}\) and \(\mathrm{Br}\) 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 \(\mathrm{ReCl}_{6}\) (c) \(\mathrm{PbCl}_{4}\) and \(\mathrm{RbCl}\).

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) \(\mathrm{TiCl}_{4}\) and \(\mathrm{CaF}_{2}\), (b) \(\mathrm{ClF}_{3}\) and \(\mathrm{VF}_{3}\), (c) \(\mathrm{SbCl}_{5}\) and \(\mathrm{AlF}_{3}\).

Draw Lewis structures for the following: (a) \(\mathrm{SiH}_{4},\) (b) 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), \((\mathrm{d}) \mathrm{AsO}_{3}^{3-},(\mathrm{e}) \mathrm{H}_{2} \mathrm{SO}_{3}(\mathrm{H}\) is bonded to \(\mathrm{O}),(\mathrm{f}) \mathrm{C}_{2} \mathrm{H}_{2}\)

(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.

Write Lewis structures that obey the octet rule for each of the following, and assign oxidation numbers and formal charges to each atom: (a) \(\mathrm{OCS},(\mathrm{b}) \mathrm{SOCl}_{2}(\mathrm{~S}\) is bonded to the two \(\mathrm{Cl}\) (d) \(\mathrm{HClO}_{2}(\mathrm{H}\) is bonded to \(\mathrm{O})\). atoms and to the \(\mathrm{O}\) ), (c) \(\mathrm{BrO}_{3}^{-}\),

For each of the following molecules or ions of sulfur and oxygen, write a single Lewis structure that obeys the octet rule, and calculate the oxidation numbers and formal charges on all the atoms: \((\mathrm{a}) \mathrm{SO}_{2},(\mathrm{~b}) \mathrm{SO}_{3},(\mathrm{c}) \mathrm{SO}_{3}{\underline{\phantom{xx}}}^{2-},\) (d) Arrange these mol- ecules/ions in order of increasing \(\mathrm{S}-\mathrm{O}\) bond distance.

(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 and double bonds?

Consider the formate ion, \(\mathrm{HCO}_{2}^{-},\) which is the anion formed when formic acid loses an \(\mathrm{H}^{+}\) ion. The \(\mathrm{H}\) and the two \(\mathrm{O}\) atoms are bonded to the central \(\mathrm{C}\) atom. (a) Write one or more appropriate Lewis structures for this ion. (b) Are resonance structures needed to describe the structure? (c) What would you predict for the \(\mathrm{C}-\mathrm{O}\) bond lengths in the formate ion relative to those in \(\mathrm{CO}_{2} ?\)

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 \(\mathrm{C}-\mathrm{C}\) single 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?

The chlorine oxides, in which a chlorine atom is bonded to one or more oxygen atoms, are important molecules in the chemistry of the atmosphere. Will any of the chlorine oxides obey the octet rule? Why or why not?

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) \(\mathrm{SO}_{3}^{2-}\), (b) \(\mathrm{AlH}_{3}\) (c) \(\mathrm{N}_{3}^{-}\) (d) \(\mathrm{CH}_{2} \mathrm{Cl}_{2}\) (e) \(\mathrm{SbF}_{5}\)