Chapter 9

The molecule shown here is diffuoromethane \(\left(\mathrm{CH}_{2} \mathrm{F}_{2}\right),\) which is used as a refrigerant called \(\mathrm{R}-32\) . (a) Based on the structure, how many electron domains surround the C atom in this molecule? (b) Would the molecule have a nonzero dipole moment? (c) If the molecule is polar, which of the following describes the direction of the overall dipole moment vector in the molecule: (i) from the carbon atom toward a fluorine atom, (ii) from the carbon atom to a point midway between the fluorine atoms, (iii) from the carbon atom to a point midway between the hydrogen atoms, or (iv) from the carbon atom toward a hydrogen atom? [Sections 9.2 and 9.3\(]\)

(a) An AB \(_{2}\) molecule is linear. How many non bonding electron pairs are around the A atom from this information? (b) How many non bonding electrons surround the Xe in \(\mathrm{XeF}_{2} ?(\mathbf{c})\) Is XeF \(_{2}\) linear?

(a) Methane \(\left(\mathrm{CH}_{4}\right)\) and the perchlorate ion \(\left(\mathrm{ClO}_{4}^{-}\right)\) are both described as tetrahedral. What does this indicate about their bond angles? (b) The \(\mathrm{NH}_{3}\) molecule is trigonal pyramidal, while \(\mathrm{BF}_{3}\) is trigonal planar. Which of these molecules is flat?

How does a trigonal pyramid differ from a tetrahedron so far as molecular geometry is concerned?

Describe the bond angles to be found in each of the following molecular structures: (a) trigonal planar, (b) tetrahedral, (c) octahedral, (d) linear.

(a) An AB \(_{6}\) molecule has no lone pairs of electrons on the A atom. What is its molecular geometry? (b) An AB \(_{4}\) molecule has two lone pairs of electrons on the A atom (in addition to the four B atoms). What is the electron-domain geometry around the A atom? (c) For the AB \(_{4}\) molecule in part (b), predict the molecular geometry.

Would you expect the nonbonding electron-pair domain in \(\mathrm{NH}_{3}\) to be greater or less in size than the corresponding one in \(\mathrm{PH}_{3}\) ?

In which of these molecules or ions does the presence of nonbonding electron pairs produce an effect on molecular shape? (a) \(\operatorname{SiH}_{4,}(\mathbf{b}) \mathrm{PF}_{3},(\mathbf{c}) \mathrm{HBr},(\mathbf{d}) \mathrm{HCN},(\mathbf{e}) \mathrm{SO}_{2}\)

In which of the following molecules can you confidently predict the bond angles about the central atom, and for which would you be a bit uncertain? Explain in each case. \((\mathbf{a}) \mathrm{H}_{2} \mathrm{S},(\mathbf{b}) \mathrm{BCl}_{3},(\mathbf{c}) \mathrm{CH}_{3} \mathrm{I},(\mathbf{d}) \mathrm{CBr}_{4,}(\mathbf{e}) \mathrm{TeBr}_{4}\)

How many nonbonding electron pairs are there in each of the following molecules: \((\mathrm{a})\left(\mathrm{CH}_{3}\right)_{2} \mathrm{S},(\mathbf{b}) \mathrm{HCN},(\mathbf{c}) \mathrm{C}_{2} \mathrm{H}_{2}\) \((\mathbf{d}) \mathrm{CH}_{3} \mathrm{F} ?\)

Describe the characteristic electron-domain geometry of each of the following numbers of electron domains about a central atom: \((\mathbf{a}), \mathbf{( b )} 4,(\mathbf{c}) 5,(\mathbf{d}) 6\)

Give the electron-domain and molecular geometries of a molecule that has the following electron domains on its central atom: (a) four bonding domains and no nonbonding domains, (b) three bonding domains and two nonbonding domains, (c) five bonding domains and one nonbonding domain, (d) four bonding domains and two nonbonding domains.

What are the electron-domain and molecular geometries of a molecule that has the following electron domains on its central atom? (a) Three bonding domains and no nonbonding domains, (b) three bonding domains and one nonbonding domain, (c) two bonding domains and two nonbonding domains.

Give the electron-domain and molecular geometries for the following molecules and ions: (a) \(\mathrm{HCN},(\mathbf{b}) \mathrm{SO}_{3}^{2-},(\mathbf{c}) \mathrm{SF}_{4}\) \((\mathbf{d}) \mathrm{PF}_{6},(\mathbf{e}) \mathrm{NH}_{3} \mathrm{Cl}^{+},(\mathbf{f}) \mathrm{N}_{3}^{-}\)

Draw the Lewis structure for each of the following molecules or ions, and predict their electron-domain and molecular geometries: (a) \(\operatorname{AsF}_{3},(\mathbf{b}) \mathrm{CH}_{3}^{+},(\mathbf{c}) \operatorname{Br} \mathrm{F}_{3},(\mathbf{d}) \mathrm{ClO}_{3},(\mathbf{e}) \mathrm{XeF}_{2}\) \((\mathbf{f}) \mathrm{BrO}_{2}^{-}\)

Ammonia, \(\mathrm{NH}_{3},\) reacts with incredibly strong bases to produce the amide ion, NH \(_{2}\) . Ammonia can also react with acids to produce the ammonium ion, \(\mathrm{NH}_{4}^{+} .\) (a) Which species (amide ion, ammonia, or ammonium ion) has the largest \(\mathrm{H}-\mathrm{N}-\mathrm{H}\) bond angle? (b) Which species has the smallest \(\mathrm{H}-\mathrm{N}-\mathrm{H}\) bond angle?

In which of the following \(\mathrm{AF}_{n}\) molecules or ions is there more than one \(\mathrm{F}-\mathrm{A}-\mathrm{Fbond}\) angle: \(\mathrm{SiF}_{4}, \mathrm{PF}_{5}, \mathrm{SF}_{4}, \mathrm{AsF}_{3} ?\)

Name the proper three-dimensional molecular shapes for each of the following molecules or ions, showing lone pairs as needed: \((\mathbf{a}) \mathrm{ClO}_{2}^{-}(\mathbf{b}) \mathrm{SO}_{4}^{2-}(\mathbf{c}) \mathrm{NF}_{3}(\mathbf{d}) \mathrm{CCl}_{2} \mathrm{Br}_{2}(\mathbf{e}) \mathrm{SF}_{4}^{2+}\)

What is the distinction between a bond dipole and a molecular dipole moment?

Consider a molecule with formula \(\mathrm{AX}_{3}\) . Supposing the \(\mathrm{A}-\mathrm{X}\) bond is polar, how would you expect the dipole moment of the \(\mathrm{AX}_{3}\) molecule to change as the \(\mathrm{X}-\mathrm{A}-\mathrm{X}\) bond angle increases from \(100^{\circ}\) to \(120^{\circ}\)

(a) Does SCl\(_{2}\) have a dipole moment? If so, in which direction does the net dipole point? (b) Does BeCl\(_{2}\) have a dipole moment? If so, in which direction does the net dipole point?

(a) The PH \(_{3}\) molecule is polar. Does this offer experimental proof that the molecule cannot be planar? Explain. (b) It turns out that ozone, \(\mathrm{O}_{3},\) has a small dipole moment. How is this possible, given that all the atoms are the same?

(a) Is the molecule BF \(_{3}\) polar or nonpolar? (b) If you react BF \(_{3}\) to make the ion \(\mathrm{BF}_{3}^{2-}\) , is this ion planar? (c) Does the molecule BF\(_{2}\)Cl have a dipole moment?

Predict whether each of the following molecules is polar or nonpolar: (a) IF, (b) \(\mathrm{CS}_{2},(\mathbf{c}) \mathrm{SO}_{3},(\mathbf{d}) \mathrm{PCl}_{3},(\mathbf{e}) \mathrm{SF}_{6},(\mathbf{f}) \mathrm{IF}_{5}\)

Predict whether each of the following molecules is polar or nonpolar: \((\mathbf{a}){CCl}_{4},(\mathbf{b}) \mathrm{NH}_{3},(\mathbf{c}) \mathrm{SF}_{4},(\mathbf{d}) \mathrm{XeF}_{4},(\mathbf{e}) \mathrm{CH}_{3} \mathrm{Br}\)

Dichloroethylene \(\left(\mathrm{C}_{2} \mathrm{H}_{2} \mathrm{Cl}_{2}\right)\) has three forms (isomers), each of which is a different substance. (a) Draw Lewis structures of the three isomers, all of which have a carbon-carbon double bond. ( b) Which of these isomers has a zero dipole moment? (c) How many isomeric forms can chloroethylene, \(\mathrm{C}_{2} \mathrm{H}_{3} \mathrm{Cl}\) have? Would they be expected to have dipole moments?

For each statement, indicate whether it is true or false. (a) In order to make a covalent bond, the orbitals on each atom in the bond must overlap. (b) A p orbital on one atom cannot make a bond to an s orbital on another atom. (c) Lone pairs of electrons on an atom in a molecule influence the shape of a molecule. (d) The 1 s orbital has a nodal plane. (e) The \(2p\) orbital has a nodal plane.

Draw sketches illustrating the overlap between the following orbitals on two atoms: (a) the 2 s orbital on each atom, (b) the 2\(p_{z}\) orbital on each atom (assume both atoms are on the \(z\) -axis), (c) the 2 s orbital on one atom and the 2\(p_{z}\) orbital on the other atom.

For each statement, indicate whether it is true or false. (a) The greater the orbital overlap in a bond, the weaker the bond. (b) The greater the orbital overlap in a bond, the shorter the bond. (c) To create a hybrid orbital, you could use the sorbital on one atom with a porbital on another atom. (d) Nonbonding electron pairs cannot occupy a hybrid orbital.

How would you expect the extent of overlap of the bonding atomic orbitals to vary in the series IF, ICl, IBr, and \(I_{2} ?\) Explain your answer.

Consider the molecule \(\mathrm{BF}_{3}\). (a) What is the electron configuration of an isolated B atom? (b) What is the electron configuration of an isolated F atom? (c) What hybrid orbitals should be constructed on the B atom to make the B–F bonds in \(\mathrm{B} \mathrm{F}_{3}\)?(d) What valence orbitals, if any, remain unhybridized on the B atom in \(\mathrm{BF}_{3} ?\)

Consider the \(\mathrm{SCl}_{2}\) molecule. (a) What is the electron configuration of an isolated S atom? (b) What is the electron configuration of an isolated Cl atom? (c) What hybrid orbitals should be constructed on the S atom to make the S-Cl bonds in \(\mathrm{SCl}_{2} ?\) (d) What valence orbitals, if any, remain unhybridized on the S atom in \(\mathrm{SCl}_{2} ?\)

Indicate the hybridization of the central atom in \((\mathbf{a}) \mathrm{BCl}_{3}$$(\mathbf{b}) \mathrm{AlCl}_{4}^{-},(\mathbf{c}) \mathrm{CS}_{2},(\mathbf{d}) \mathrm{GeH}_{4}\)

What is the hybridization of the central atom in (a) \(\mathrm{SiCl}_{4}\), \((\mathbf{b}) \mathrm{HCN},(\mathbf{c}) \mathrm{SO}_{3},(\mathbf{d}) \mathrm{TeCl}_{2} ?\)

(a) Which geometry and central atom hybridization would you expect in the series \(\mathrm{BH}_{4}^{-}, \mathrm{CH}_{4}, \mathrm{NH}_{4}^{+} ?\) (b) What would you expect for the magnitude and direction of the bond dipoles in this series? (c) Write the formulas for the analogous species of the elements of period 3; would you expect them to have the same hybridization at the central atom?

(a) If the valence atomic orbitals of an atom are sp hybridized, how many unhybridized \(p\) orbitals remain in the valence shell? How many \(\pi\) bonds can the atom form? (b) Imagine that you could hold two atoms that are bonded together, twist them, and not change the bond length. Would it be easier to twist (rotate) around a single \(\sigma\) bond or around a double \((\sigma\) plust (rotate) around a single \(\sigma\) bond same?

(a) Draw Lewis structures for ethane \(\left(\mathrm{C}_{2} \mathrm{H}_{6}\right),\) ethylene \(\left(\mathrm{C}_{2} \mathrm{H}_{4}\right),\) and acetylene \(\left(\mathrm{C}_{2} \mathrm{H}_{2}\right)\) (b) What is the hybridization of the carbon atoms in each molecule? (c) Predict which molecules, if any, are planar. (d) How many \(\sigma\) and \(\pi\) bonds are there in each molecule?

The nitrogen atoms in \(\mathrm{N}_{2}\) participate in multiple bonding, whereas those in hydrazine, \(\mathrm{N}_{2} \mathrm{H}_{4},\) do not. (a) Draw Lewis structures for both molecules. (b) What is the hybridization of the nitrogen atoms in each molecule? (c) Which molecule has the stronger \(N-N\) bond?

Ethyl acetate, \(\mathrm{C}_{4} \mathrm{H}_{8} \mathrm{O}_{2},\) is a fragrant substance used both as a solvent and as an aroma enhancer. Its Lewis structure is (a) What is the hybridization at each of the carbon atoms of the molecule? (b) What is the total number of valence electrons in ethyl acetate? (c) How many of the valence electrons are used to make \(\sigma\) bonds in the molecule? (d) How many valence electrons are used to make \(\pi\) bonds? (e) How many valence electrons remain in nonbonding pairs in the molecule?

(a) What is the difference between a localized \(\pi\) bond and a delocalized one? (b) How can you determine whether a molecule or ion will exhibit delocalized \(\pi\) bonding? (c) Is the \(\pi\) bond in \(\mathrm{NO}_{2}^{-}\) localized or delocalized?

(a) Write a single Lewis structure for \(S O_{3},\) and determine the hybridization at the S atom. (b) Are there other equivalent Lewis structures for the molecule? (c) Would you expect SO \(_{3}\) to exhibit delocalized \(\pi\) bonding?

In the formate ion, \(\mathrm{HCO}_{2}^{-}\) , the carbon atom is the central atom with the other three atoms attached to it. (a) Draw a Lewis structure for the formate ion. (b) What hybridization is exhibited by the C atom? (c) Are there multiple equivalent resonance structures for the ion? (d) How many electrons are in the \(\pi\) system of the ion?

(a) What is the difference between hybrid orbitals and molecular orbitals? (b) How many electrons can be placed into each MO of a molecule? (c) Can antibonding molecular orbitals have electrons in them?

(a) If you combine two atomic orbitals on two different atoms to make a new orbital, is this a hybrid orbital or a molecular orbital? (b) If you combine two atomic orbitals on one atom to make a new orbital, is this a hybrid orbital or a molecular orbital? (c) Does the Pauli exclusion principle (Section 6.7) apply to MOs? Explain.

Consider the \(\mathrm{H}_{2}^{+}\) ion. (a) Sketch the molecular orbitals of the ion and draw its energy-level diagram. (b) How many electrons are there in the \(\mathrm{H}_{2}+\) ion? (c) Write the electron configuration of the ion in terms of its MOs. (d) What is the bond order in \(\mathrm{H}_{2}^{+} ?\) (e) Suppose that the ion is excited by light so that an electron moves from a lower-energy to a higher-energy MO. Would you expect the excited-state \(\mathrm{H}_{2}^{+}\) ion to be stable or to fall apart? (f) Which of the following statements about part (e) is correct: (i) The light excites an electron from a bonding orbital to an antibonding orbital, (ii) The light excites an electron from an antibonding orbital to a bonding orbital, or (iii) In the excited state there are more bonding electrons than antibonding electrons?

(a) Sketch the molecular orbitals of the \(\mathrm{H}_{2}^{-}\) ion and draw its energy-level diagram.(b) Write the electron configuration of the ion in terms of its MOs. (c) Calculate the bond order in \(\mathrm{H}_{2}^{-} .(\mathbf{d})\) Suppose that the ion is excited by light, so that an electron moves from a lower-energy to a higher-energy molecular orbital. Would you expect the excited-state \(\mathrm{H}_{2}\) -ion to be stable? (e) Which of the following statements about part (d) is correct: (i) The light excites an electron from a bonding orbital to an antibonding orbital, (ii) The light excites an electron from an antibonding orbital to a bonding orbital, or (iii) In the excited state there are more bonding electrons than antibonding electrons?

Draw a picture that shows all three 2\(p\) orbitals on one atom and all three 2\(p\) orbitals on another atom. (a) Imagine the atoms coming close together to bond. How many \(\sigma\) bonds can the two sets of 2\(p\) orbitals make with each other? (b) How many \(\pi\) bonds can the two sets of 2\(p\) orbitals make with each other? (c) How many antibonding orbitals, and of what type, can be made from the two sets of 2\(p\) orbitals?

Indicate whether each statement is true or false. (a) \(s\) orbitals can only make \(\sigma\) or \(\sigma^{*}\) molecular orbitals. (b) The probability is 100\(\%\) for finding an electron at the nucleus in a \(\pi^{*}\) orbital. (c) Antibonding orbitals are higher in energy than bonding orbitals (if all orbitals are created from the same atomic orbitals). (d) Electrons cannot occupy an antibonding orbital.

(a) What does the term diamagnetism mean? (b) How does a diamagnetic substance respond to a magnetic field? (c) Which of the following ions would you expect to be diamagnetic: \(N_{2}^{2-}, \mathrm{O}_{2}^{2-}, \mathrm{Be}_{2}^{2+}, \mathrm{C}_{2}^{-} ?\)

(a) What does the term paramagnetism mean? (b) How can one determine experimentally whether a substance is paramagnetic? (c) Which of the following ions would you expect to be paramagnetic: \(\mathrm{O}_{2}^{+}, \mathrm{N}_{2}^{2-}, \mathrm{Li}_{2}^{+}, \mathrm{O}_{2}^{2-} ?\) For those ions that are paramagnetic, determine the number of unpaired electrons.