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: a \(2+\) charge for \(\mathrm{A}\) and a \(1+\) charge for \(\mathrm{B}\). Ions shown in blue carry negative charges: a \(1-\) charge for \(X\) and a \(2-\) charge for \(Y\). (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? [Section 8.2]

Incomplete Lewis structures for the nitrous acid molecule, \(\mathrm{HNO}_{2}\), and the nitrite ion, \(\mathrm{NO}_{2}^{-}\), are shown here. (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 \(\mathrm{NO}_{2}^{-}\) ? [Sections 8.5 and 8.6 ] $$ \mathrm{H}-\mathrm{O}-\mathrm{N}=\mathrm{O} \quad \mathrm{O}-\mathrm{N}=\mathrm{O} $$

(a) True or false: An element's number of valence electrons is the same as its atomic number. (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) True or false: The hydrogen atom is most stable when it has a full octet of electrons. (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?

Consider the element silicon, Si. (a) Write its electron configuration. (b) How many valence electrons does a silicon atom have? (c) Which subshells hold the valence electrons?

(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 . 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: \((\mathbf{a}) \mathrm{Te},(\mathbf{b}) \mathrm{Si},(\mathbf{c}) \mathrm{Kr},(\mathbf{d}) \mathrm{P}\).

What is the Lewis symbol for each of the following atoms or ions? \((\mathbf{a}) \mathrm{Be},(\mathbf{b}) \mathrm{Rb},(\mathbf{c}) \mathrm{I}^{-},(\mathbf{d}) \mathrm{Se}^{2-} .\)

(a) Using Lewis symbols, make a sketch of the reaction between potassium and bromine atoms to give the ionic substance KBr. (b) How many electrons are transferred? (c) Which atom loses electrons in the reaction?

(a) Use Lewis symbols to represent the reaction that occurs between Li and O atoms. (b) What is the chemical formula of the most likely product? (c) How many electrons are transferred? (d) Which atom loses electrons in the reaction?

Predict the chemical formula of the ionic compound formed between the following pairs of elements: (a) Al and Cl, (d) Li and \(O\). (b) \(\mathrm{Mg}\) and \(\mathrm{O},(\mathbf{c}) \mathrm{Zn}\) and \(\mathrm{Cl},\)

Which ionic compound is expected to form from combining the following pairs of elements? (a) calcium and nitrogen, (b) cesium and bromine, (c) strontium and sulfur, (d) aluminum and selenium.

Write the electron configuration for each of the following ions. and determine which ones possess noble-gas configurations: (a) \(\mathrm{Be}^{2+}\) (b) \(\mathrm{Mn}^{2+},(\mathbf{c}) \mathrm{Cd}^{2+}\) (d) \(\mathrm{Fe}^{3+}\), (e) \(\mathrm{Tl}^{+}\), (f) \(\mathrm{At}^{-}\).

Write the electron configurations for the following ions, and determine which have noble-gas configurations: (a) \(\mathrm{Fe}^{2+}\), (b) \(\mathrm{V}^{3+},(\mathbf{c}) \mathrm{Ni}^{2+},(\mathbf{d}) \mathrm{Pt}^{2+}\) (e) \(\mathrm{Ge}^{2-},(\mathbf{f}) \mathrm{Ba}^{2+}\).

(a) Is lattice energy usually endothermic or exothermic? (b) Write the chemical equation that represents the process of lattice energy for the case of NaCl. (c) Would you expect salts like NaCl, which have singly charged ions, to have larger or smaller lattice energies compared to salts like \(\mathrm{CaO}\) which are composed of doubly-charged ions?

(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 8.1 according to their expected lattice energies, listing them from lowest lattice energy to the highest: \(\mathrm{MgS},\) KI, GaN, LiBr.

Which of the following trends in lattice energy is due to differences in ionic radii? (a) \(\mathrm{LiF}>\mathrm{NaF}>\mathrm{CsF},(\mathbf{b}) \mathrm{CaO}>\mathrm{KCl}\), (c) \(\mathrm{PbS}>\mathrm{Li}_{2} \mathrm{O}\).

Energy is required to remove two electrons from Ca to form \(\mathrm{Ca}^{2+},\) and energy is required to add two electrons to \(\mathrm{O}\) to form \(\mathrm{O}^{2-}\). Yet \(\mathrm{CaO}\) is stable relative to the free elements. Which statement is the best explanation? (a) The lattice energy of \(\mathrm{CaO}\) is large enough to overcome these processes. (b) \(\mathrm{CaO}\) is a covalent compound, and these processes are irrelevant. (c) CaO has a higher molar mass than either Ca or O. (d) The enthalpy of formation of \(\mathrm{CaO}\) is small. \((\mathbf{e}) \mathrm{CaO}\) is stable to atmospheric conditions.

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

Which of these elements are unlikely to form ionic bonds? \(\mathrm{Mg}, \mathrm{Al}, \mathrm{Si}, \mathrm{Br}, \mathrm{I}\).

Using Lewis symbols and Lewis structures, make a sketch of the formation of \(\mathrm{NCl}_{3}\) from \(\mathrm{N}\) and \(\mathrm{Cl}\) atoms, showing valence- shell electrons. (a) How many valence electrons does N have initially? (b) How many bonds Cl has to make in order to achieve an octet? (c) How many valence electrons surround the \(\mathrm{N}\) in the \(\mathrm{NCl}_{3}\) molecule? (d) How many valence electrons surround each Cl in the \(\mathrm{NCl}_{3}\) molecule? (e) How many lone pairs of electrons are in the \(\mathrm{NCl}_{3}\) molecule?

Using Lewis symbols and Lewis structures, diagram the formation of \(\mathrm{BF}_{3}\) from \(\mathrm{B}\) and \(\mathrm{F}\) atoms, showing valence- shell electrons. (a) How many valence electrons does B have initially? (b) How many bonds F has to make in order to achieve an octet? (c) How many valence electrons surround the B in the \(\mathrm{BF}_{3}\) molecule? (d) How many valence electrons surround each F in the \(\mathrm{BF}_{3}\) molecule? (e) Does \(\mathrm{BF}_{3}\) obey the octet rule?

(a) Construct a Lewis structure for \(\mathrm{O}_{2}\) in which each atom achieves an octet of electrons. (b) How many bonding electrons are in the structure? (c) Would you expect the \(\mathrm{O}-\mathrm{O}\) bond in \(\mathrm{O}_{2}\) to be shorter or longer than the \(\mathrm{O}-\mathrm{O}\) bond in compounds that contain an \(\mathrm{O}-\mathrm{O}\) single bond? Explain.

(a) Construct a Lewis structure for hydrogen peroxide, \(\mathrm{H}_{2} \mathrm{O}_{2}\), in which each atom achieves an octet of electrons. (b) How many bonding electrons are between the two oxygen atoms? (c) 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 \(\mathrm{O}_{2}\) ? Explain.

Which of the following statements about electronegativity is false? (a) Electronegativity is the ability of an atom in a molecule to attract electron density toward itself. (b) Electronegativity is the same thing as electron affinity. (c) The numerical values for electronegativity have no units. (d) Fluorine is the most electronegative element. (e) Cesium is the least electronegative element.

(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) True or false: The most easily ionizable elements are the most electronegative.

Using only the periodic table as your guide, select the most electronegative atom in each of the following sets: (a) Se, Te, \(\mathrm{Br}, \mathrm{I} ;(\mathbf{b}) \mathrm{Be}, \mathrm{Mg}, \mathrm{C}, \mathrm{Si} ;(\mathbf{c}) \mathrm{Al}, \mathrm{Si}, \mathrm{P}, \mathrm{S} ;(\mathbf{d}) \mathrm{O}, \mathrm{P}, \mathrm{Ge}, \mathrm{In}\).

By referring only to the periodic table, select \((\mathbf{a})\) the most electronegative element in group \(13 ;(\mathbf{b})\) the least electronegative element in the group As, Se and Br; (c) the most electronegative element in the group K, Mg, Al and In; (d) the element in the group \(\mathrm{Na}\), Be, \(\mathrm{Si}\), Ar, that is most likely to form an ionic compound with \(B \mathrm{r}\).

Which of the following bonds are polar? (a) \(\mathrm{C}-\mathrm{O}\) (b) \(\mathrm{Sl}-\mathrm{F},(\mathbf{c}) \mathrm{N}-\mathrm{Cl},(\mathbf{d}) \mathrm{C}-\mathrm{Cl}\). Which is the more electronegative atom in each polar bond?

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

The iodine monobromide molecule, IBr, has a bond length of \(249 \mathrm{pm}\) and a dipole moment of \(1.21 \mathrm{D}\). (a) Which atom of the molecule is expected to have a negative charge? (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: \((\mathbf{a}) \mathrm{SiF}_{4}\) and \(\mathrm{LaF}_{3},(\mathbf{b}) \mathrm{FeCl}_{2}\) and \(\mathrm{ReCl}_{6},(\mathbf{c}) \mathrm{PbCl}_{4}\) and 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: \((\mathbf{a}) \mathrm{TiCl}_{4}\) and \(\mathrm{CaF}_{2},(\mathbf{b}) \mathrm{ClF}_{3}\) and \(\mathrm{VF}_{3},(\mathbf{c}) \mathrm{SbCl}_{5}\) and \(\mathrm{AlF}_{3} .\)

Draw Lewis structures for the following: \((\mathbf{a}) \mathrm{CH}_{2} \mathrm{Cl}_{2},(\mathbf{b}) \mathrm{ClCN},\) (c) AsF \(_{5},\) (d) \(\mathrm{CH}_{2} \mathrm{O}\) (C the central atom), \((\mathbf{e}) \mathrm{OF}_{2},(\mathbf{f}) \mathrm{NO}_{3}^{-}\).

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

Which one of these statements about formal charge is true? (a) Formal charge is the same as oxidation number. (b) To draw the best Lewis structure, you should minimize formal charge. (c) Formal charge takes into account the different electronegativities of the atoms in a molecule. (d) Formal charge is most useful for ionic compounds. (e) Formal charge is used in calculating the dipole moment of a diatomic molecule.

(a) Draw the dominant Lewis structure for the phosphorus trifluoride molecule, \(\mathrm{PF}_{3}\). (b) Determine the oxidation numbers of the \(\mathrm{P}\) and \(\mathrm{F}\) atoms. \((\mathbf{c})\) Determine the formal charges of the \(\mathrm{P}\) and \(\mathrm{F}\) atoms.

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

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: (a) \(\mathrm{SO}_{2}\), (b) \(\mathrm{SO}_{3}\) (c) \(\mathrm{SO}_{3}^{2-}\). (d) Arrange these molecules/ions in order of increasing \(\mathrm{S}-\mathrm{O}\) bond length.

(a) Draw the best Lewis structure(s) 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 C atom. (a) Draw the best Lewis structure(s) for this ion. (b) Are resonance structures needed to describe the structure? (c) Would you predict that the \(\mathrm{C}-\mathrm{O}\) bond lengths in the formate ion would be longer or shorter relative to those in \(\mathrm{CO}_{2}\) ?

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

(a) Which of these compounds is an exception to the octet rule: carbon dioxide, water, ammonia, phosphorus trifluoride, or arsenic pentafluoride? (b) Which of these compounds or ions is an exception to the octet rule: borohydride \(\left(\mathrm{BH}_{4}^{-}\right)\), borazine \(\left(\mathrm{B}_{3} \mathrm{~N}_{3} \mathrm{H}_{6},\right.\) which is analogous to benzene with alternating \(\mathrm{B}\) and \(\mathrm{N}\) in the ring), or boron trichloride?

Fill in the blank with the appropriate numbers for both electrons and bonds (considering that single bonds are counted as one, double bonds as two, and triple bonds as three). (a) Iodine has valence electrons and makes ___________ bond(s) in compounds. (b) Silicon has ____________ valence electrons and makes _________ bond(s) in compounds. (c) Phosphorus has ___________ valence electrons and makes ________ bond(s) in compounds. (d) Sulphur has __________ valence electrons and makes __________ bond(s) in compounds.

Draw the dominant Lewis structures for these chlorineoxygen molecules/ions: \(\mathrm{ClO}, \mathrm{ClO}^{-}, \mathrm{ClO}_{2}^{-}, \mathrm{ClO}_{3}^{-}, \mathrm{ClO}_{4}^{-}\). Which of these do not obey the octet rule?

For Group 13-17 elements in the third row of the periodic table and beyond, the octet rule is often not obeyed. A friend of yours says this is because these heavier elements are more likely to make double or triple bonds. Another friend of yours says that this is because the heavier elements are larger and can make bonds to more than four atoms at a time. Which friend is more correct?

Draw the Lewis structures for each of the following ions or molecules. Identify those in which the octet rule is not obeyed; state which atom in each compound does not follow the octet rule; and state, for those atoms, how many electrons surround them: \((\mathbf{a}) \mathrm{HCl},(\mathbf{b}) \mathrm{ICl}_{5},(\mathbf{c}) \mathrm{NO},\) (d) \(\mathrm{CF}_{2} \mathrm{Cl}_{2},(\mathbf{e}) \mathrm{I}_{3}^{-}\).

Draw the Lewis structures for each of the following molecules or ions. Identify instances where the octet rule is not obeyed; state which atom in each compound does not follow the octet rule; and state how many electrons surround these atoms: \((\mathbf{a}) \mathrm{PF}_{6}^{-},(\mathbf{b}) \mathrm{BeCl}_{2},(\mathbf{c}) \mathrm{NH}_{3},(\mathbf{d}) \mathrm{XeF}_{2} \mathrm{O}\) (the Xe is the central atom), \((\mathbf{e}) \mathrm{SO}_{4}^{2-} .\)

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 structures are possible that satisfy the octet rule? (c) On the basis of the formal charges, which Lewis structure is expected to be dominant for \(\mathrm{BeCl}_{2} ?\)

(a) Describe the molecule xenon trioxide, \(\mathrm{XeO}_{3},\) using four possible Lewis structures, one each with zero, one, two, or three \(\mathrm{Xe}-\mathrm{O}\) double bonds. (b) Do any of these resonance structures satisfy the octet rule for every atom in the molecule? (c) Do any of the four Lewis structures have multiple resonance structures? If so, how many resonance structures do you find? (d) Which of the Lewis structures in part (a) yields the most favorable formal charges for the molecule?