Chapter 8

For each of these Lewis symbols, indicate the group in the periodic table in which the element \(X\) belongs: [Section 8.1] (a) \(\dot{X}\) (b) \(\mathrm{X}\) (c) \(\dot{X} \cdot\)

Illustrated are four ions - \(A, B, X\), and \(Y-\) showing their relative ionic radii. The ions shown in red carry positive charges: a \(2+\) charge for A and a \(1+\) charge for 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]

(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 4B. 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) \(\mathrm{Al}\), (b) \(\mathrm{Br}\), (c) \(\mathrm{Ar}\), (d) \(\mathrm{Sr}\).

What is the Lewis symbol for each of the following atoms or ions? (a) \(\mathrm{K}\), (b) \(\mathrm{As}\), (c) \(\mathrm{Sn}^{2+}\), (d) \(\mathrm{N}^{3-}\).

(a) Using Lewis symbols, diagram the reaction between magnesium and oxygen atoms to give the ionic substance \(\mathrm{MgO}\). (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 Ca and F 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) \(\mathrm{Al}\) 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) \(\mathrm{Cd}^{2+}\), (b) \(\mathrm{P}^{3-}\), (c) \(\mathrm{Zr}^{4+}\), (d) \(\mathrm{Ru}^{3+}\), (e) \(\mathrm{As}^{3-}\), (f) \(\mathrm{Ag}^{+}\).

(a) Is lattice energy usually endothermic or exothermic? (b) Write the chemical equation that represents the process of lattice energy for the case of \(\mathrm{NaCl}\). (c) Would you expect salts like \(\mathrm{NaCl}\), 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 \(8.2\) 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) \(\mathrm{NaCl}>\mathrm{RbBr}>\mathrm{CsBr}\), (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 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 \(\mathrm{O}\). (d) The enthalpy of formation of \(\mathrm{CaO}\) is small. (e) \(\mathrm{CaO}\) is stable to atmospheric conditions.

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

(a) State whether the bonding in each compound is likely to be covalent or not: (i) iron, (ii) sodium chloride, (iii) water, (iv) oxygen, (v) argon. (b) 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?

Using Lewis symbols and Lewis structures, diagram the formation of \(\mathrm{SiCl}_{4}\) from \(\mathrm{Si}\) and \(\mathrm{Cl}\) atoms, showing valence-shell electrons. (a) How many valence electrons does \(\mathrm{Si}\) have initially? (b) How many valence electrons does each \(\mathrm{Cl}\) have initially? (c) How many valence electrons surround the Si in the \(\mathrm{SiCl}_{4}\) molecule? (d) How many valence electrons surround each \(\mathrm{Cl}\) in the \(\mathrm{SiCl}_{4}\) molecule? (e) How many bonding pairs of electrons are in the \(\mathrm{SiCl}_{4}\) molecule?

Use Lewis symbols and Lewis structures to diagram the formation of \(\mathrm{PF}_{3}\) from \(\mathrm{P}\) and \(\mathrm{F}\) atoms, showing valence- shell electrons. (a) How many valence electrons does \(P\) have initially? (b) How many valence electrons does each \(F\) have initially? (c) How many valence electrons surround the \(\mathrm{P}\) in the \(\mathrm{PF}_{3}\) molecule? (d) How many valence electrons surround each \(\mathrm{F}\) in the \(\mathrm{PF}_{3}\) molecule? (e) How many bonding pairs of electrons are in the \(\mathrm{PF}_{3}\) molecule?

(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) \(\mathrm{Na}, \mathrm{Mg}\), K, Ca; (b) P, S, As, Se; (c) Be, B, C, Si; (d) Zn, Ge, Ga, As.

By referring only to the periodic table, select (a) the most electronegative element in group \(6 A ;(b)\) 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) B-F, (b) \(\mathrm{Cl}-\mathrm{Cl}\), (c) Se-O, (d) H-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 iodine monobromide molecule, IBr, has a bond length of \(2.49 \AA\) and a dipole moment of \(1.21\) 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: (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) \(\mathrm{CO}\), (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) \(\mathrm{NH}_{2} \mathrm{Cl}\).

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, \(\mathrm{PF}_{\mathfrak{3} .}\) (b) Determine the oxidation numbers of the \(P\) and \(F\) atoms. (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: (a) \(\mathrm{OCS}\), (b) \(\mathrm{SOCl}_{2}\) ( \(\mathrm{S}\) is the central atom), (c) \(\mathrm{BrO}_{3}^{-}\), (d) \(\mathrm{HClO}_{2}\) (H is bonded to \(\left.\mathrm{O}\right)\).

(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 \(\mathrm{C}\) 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 \(\mathrm{CO}_{2}\) ?

Predict the ordering, from shortest to longest, of the bond lengths in \(\mathrm{CO}, \mathrm{CO}_{2}\), and \(\mathrm{CO}_{3}{\underline{\phantom{xx}}}^{2-}\).

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) Do the \(\mathrm{C}-\mathrm{C}\) bond lengths in benzene alternate shortlong-short-long around the ring? Why or why not? (b) Are \(\mathrm{C}-\mathrm{C}\) bond lengths in benzene shorter than \(\mathrm{C}-\mathrm{C}\) single bonds? (c) Are \(\mathrm{C}-\mathrm{C}\) bond lengths in benzene shorter than \(\mathrm{C}=\mathrm{C}\) double bonds?

Indicate whether each statement is true or false: (a) The octet rule is based on the fact that filling in all \(s\) and \(p\) valence electrons in a shell gives eight electrons. (b) The Si in \(\mathrm{SiH}_{4}\) does not follow the octet rule because hydrogen is in an unusual oxidation state. (c) Boron compounds are frequent exceptions to the octet rule because they have too few electrons surrounding the boron. (d) Compounds in which nitrogen is the central atom are frequent exceptions to the octet rule because they have too many electrons surrounding the nitrogen.

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) Fluorine has valence electrons and makes bond(s) in compounds. (b) Oxygen has valence electrons and makes bond(s) in compounds. (c) Nitrogen has valence electrons and makes bond(s) in compounds. (d) Carbon has valence electrons and makes bond \((s)\) in compounds.

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

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 these atoms: (a) \(\mathrm{PH}_{3}\), (b) \(\mathrm{AlH}_{3}\), (c) \(\mathrm{N}_{3}^{-}\), (d) \(\mathrm{CH}_{2} \mathrm{Cl}_{2}\), (e) \(\mathrm{SnF}_{6-}\)

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 Xe-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 (a) yields the most favorable formal charges for the molecule?

Consider the following statement: "For some molecules and ions, a Lewis structure that satisfies the octet rule does not lead to the lowest formal charges, and a Lewis structure that leads to the lowest formal charges does not satisfy the octet rule." Illustrate this statement using the hydrogen sulfite ion, \(\mathrm{HSO}_{3}{\underline{\phantom{xx}}}^{-}\), as an example (the \(\mathrm{H}\) atom is bonded to one of the O atoms).

Use Table \(8.4\) to estimate the enthalpy change for each of the following reactions: (a) \(\mathrm{H}_{2} \mathrm{C}=\mathrm{O}(\mathrm{g})+\mathrm{HCl}(\mathrm{g}) \longrightarrow \mathrm{H}_{3} \mathrm{C}-\mathrm{O}-\mathrm{Cl}(g)\) (b) \(\mathrm{H}_{2} \mathrm{O}_{2}(g)+2 \mathrm{CO}(g) \longrightarrow \mathrm{H}_{2}(g)+2 \mathrm{CO}_{2}(g)\) (c) \(3 \mathrm{H}_{2} \mathrm{C}=\mathrm{CH}_{2}(g) \longrightarrow \mathrm{C}_{6} \mathrm{H}_{12}(g)\) (the six carbon atoms form a six-membered ring with two \(\mathrm{H}\) atoms on each C atom)