Which of these atoms cannot serve as a central atom in a Lewis structure: (a) $\mathrm{O} ;(\mathrm{b}) \mathrm{He} ;(\mathrm{c}) \mathrm{F} ;(\mathrm{d}) \mathrm{H} ;(\mathrm{e}) \mathrm{P}$ ? Explain.

Sidharth A.

Numerade Educator

When is a resonance hybrid needed to adequately depict the bonding in a molecule? Using $\mathrm{NO}_{2}$ as an example, explain how a resonance hybrid is consistent with the actual bond length, bond strength, and bond order.

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In which of these bonding patterns does X obey the octet rule?

Sidharth A.

Numerade Educator

What is required for an atom to expand its valence shell? Which of the following atoms can expand its valence shell: $\mathrm{F}, \mathrm{S}$ , $\mathrm{H}, \mathrm{Al}, \mathrm{Se}, \mathrm{Cl} ?$

Jason D.

Numerade Educator

Draw a Lewis structure for (a) $\mathrm{SiF}_{4} ;(\mathrm{b}) \mathrm{SeCl}_{2} ;(\mathrm{c}) \mathrm{COF}_{2}(\mathrm{Cis}$central).

Sidharth A.

Numerade Educator

Draw a Lewis structure for (a) $\mathrm{PH}_{4}^{+} ;(\mathrm{b}) \mathrm{C}_{2} \mathrm{F}_{4} ;(\mathrm{c}) \mathrm{SbH}_{3}$.

Jason D.

Numerade Educator

Draw a Lewis structure for (a) $\mathrm{PF}_{3} ;(\mathrm{b}) \mathrm{H}_{2} \mathrm{CO}_{3}$ (both $\mathrm{H}$ atoms are attached to $\mathrm{O}$ atoms); (c) $\mathrm{CS}_{2}$ .

Sidharth A.

Numerade Educator

Draw a Lewis structure for $(\mathrm{a}) \mathrm{CH}_{4} \mathrm{S} ;(\mathrm{b}) \mathrm{S}_{2} \mathrm{Cl}_{2} ;(\mathrm{c}) \mathrm{CHCl}_{3}$.

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Draw Lewis structures of all the important resonance forms of (a) $\mathrm{NO}_{2}^{+} ;(\mathrm{b}) \mathrm{NO}_{2} \mathrm{F}(\mathrm{N} \text { is central). }$

Sidharth A.

Numerade Educator

Draw Lewis structures of all the important resonance forms of $(\mathrm{a}) \mathrm{HNO}_{3}\left(\mathrm{HONO}_{2}\right) ;(\mathrm{b}) \mathrm{HAsO}_{4}^{2-}\left(\mathrm{HOAsO}_{3}^{2-}\right)$

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Draw Lewis structures of all the important resonance forms of $(\mathrm{a}) \mathrm{N}_{3}^{-} ;(\mathrm{b}) \mathrm{NO}_{2}^{-}$.

Sidharth A.

Numerade Educator

Draw Lewis structures of all the important resonance forms of (a) $\mathrm{HCO}_{2}^{-}(\mathrm{H} \text { is attached to } \mathrm{C}) ;(\mathrm{b}) \mathrm{HBrO}_{4}\left(\mathrm{HOBrO}_{3}\right)$

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Draw the Lewis structure with lowest formal charges, and determine the charge of each atom in (a) $\mathrm{IF}_{5} ;(\mathrm{b}) \mathrm{AlH}_{4}^{-}$.

Sidharth A.

Numerade Educator

Draw the Lewis structure with lowest formal charges, and determine the charge of each atom in (a) $\mathrm{OCS} ;(\mathrm{b}) \mathrm{NO}$.

Jason D.

Numerade Educator

Draw the Lewis structure with lowest formal charges, and determine the charge of each atom in (a) $\mathrm{CN}^{-} ;(\mathrm{b}) \mathrm{ClO}^{-}$.

Sam L.

Numerade Educator

Draw the Lewis structure with lowest formal charges, and determine the charge of each atom in (a) $\mathrm{BF}_{4}^{-} ;(\mathrm{b}) \mathrm{ClNO}$.

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Draw a Lewis structure for a resonance form of each ion with the lowest possible formal charges, show the charges, and give oxidation numbers of the atoms: (a) $\mathrm{BrO}_{3}^{-} ;(\mathrm{b}) \mathrm{SO}_{3}^{2-}$.

Sam L.

Numerade Educator

Draw a Lewis structure for a resonance form of each ion with the lowest possible formal charges, show the charges, and give oxidation numbers of the atoms: (a) $\mathrm{AsO}_{4}^{3-} ;(\mathrm{b}) \mathrm{ClO}_{2}^{-}$.

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These species do not obey the octet rule. Draw a Lewis structure for each, and state the type of octet rule exception:

(a) $\mathrm{BH}_{3} \quad$ (b) $\mathrm{AsF}_{4}^{-} \quad$ (c) $\mathrm{SeCl}_{4}$

Sidharth A.

Numerade Educator

These species do not obey the octet rule. Draw a Lewis structure for each, and state the type of octet-rule exception:

(a) $\mathrm{PF}_{6}^{-} \quad$ (b) $\mathrm{ClO}_{3} \quad$ (c) $\mathrm{H}_{3} \mathrm{PO}_{3}$ (one $\mathrm{P}-\mathrm{H}$ bond $)$

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These species do not obey the octet rule. Draw a Lewis structure for each, and state the type of octet-rule exception:

(a) $\mathrm{BrF}_{3} \quad$ (b) $\mathrm{ICl}_{2}^{-} \quad$ (c) $\mathrm{BeF}_{2}$

Sidharth A.

Numerade Educator

These species do not obey the octet rule. Draw a Lewis structure for each, and state the type of octet-rule exception:

(a) $\mathrm{O}_{3}^{-} \quad$ (b) $\mathrm{XeF}_{2} \quad$ (c) $\mathrm{SbF}_{4}-$

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Molten beryllium chloride reacts with chloride ion from molten $\mathrm{NaCl}$ to form the $\mathrm{BeCl}_{4}^{2-}$ ion, ion, in which the $\mathrm{Be}$ atom attains an octet. Show the net ionic reaction with Lewis structures.

Sidharth A.

Numerade Educator

Despite many attempts, the perbromate ion $\left(\mathrm{BrO}_{4}^{-}\right)$ was not prepared in the laboratory until about 1970 . In fact, articles were published explaining theoretically why it could never be prepared!) Draw a Lewis structure for $\mathrm{BrO}_{4}^{-}$ in which all atoms have lowest formal charges.

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Cryolite $\left(\mathrm{Na}_{3} \mathrm{AlF}_{6}\right)$ is an indispensable component in the electrochemical production of aluminum. Draw a Lewis structure for the AlF $_{6}^{3-}$ ion.

Sam L.

Numerade Educator

Phosgene is a colorless, highly toxic gas that was employed against troops in World War I and is used today as a key reactant in organic syntheses. From the following resonance structures, select the one with the lowest formal charges:

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If you know the formula of a molecule or ion, what is the first step in predicting its shape?

Sam L.

Numerade Educator

In what situation is the name of the molecular shape the same as the name of the electron-group arrangement?

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Which of the following numbers of electron groups can give rise to a bent (V-shaped) molecule: two, three, four, five, six? Draw an example for each case, showing the shape classification $\left(\mathrm{AX}_{m} \mathrm{E}_{n}\right)$ and the ideal bond angle.

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Name all the molecular shapes that have a tetrahedral electron-group arrangement.

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Consider the following molecular shapes. (a) Which has the most electron pairs (both shared and unshared) around the central atom? (b) Which has the most unshared pairs around the central atom? (c) Do any have only shared pairs around the central atom?

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Use wedge-bond perspective drawings (if necessary) to sketch the atom positions in a general molecule of formula (not shape class) AX $_{n}$ that has each of the following shapes:

(a) V shaped $\quad$ (b) trigonal planar $\quad$ (c) trigonal bipyramidal

(d) T shaped $\quad$ (e) trigonal pyramidal $\quad$ (f) square pyramidal

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What would you expect to be the electron-group arrangement around atom A in each of the following cases? For each arrangement, give the ideal bond angle and the direction of any expected deviation:

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Determine the electron-group arrangement, molecular shape, and ideal bond angle(s) for each of the following:

(a) $\mathrm{O}_{3} \quad$ (b) $\mathrm{H}_{3} \mathrm{O}^{+} \quad$ (c) $\mathrm{NF}_{3}$

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Determine the electron-group arrangement, molecular shape, and ideal bond angle(s) for each of the following:

(a) $\mathrm{SO}_{4}^{2-} \quad$ (b) $\mathrm{NO}_{2}^{-} \quad$ (c) $\mathrm{PH}_{3}$

Sam L.

Numerade Educator

Determine the electron-group arrangement, molecular shape, and ideal bond angle(s) for each of the following:

(a) $\mathrm{CO}_{3}^{2-} \quad$ (b) $\mathrm{SO}_{2} \quad$ (c) $\mathrm{CF}_{4}$

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Determine the electron-group arrangement, molecular shape, and ideal bond angle(s) for each of the following:

(a) $\mathrm{SO}_{3} \quad$ (b) $\mathrm{N}_{2} \mathrm{O}$($\mathrm{N}$ is central) $\quad$ (c) $\mathrm{CH}_{2} \mathrm{Cl}_{2}$

Sam L.

Numerade Educator

Name the shape and give the $\mathrm{AX}_{m} \mathrm{E}_{n}$ classification and ideal bond angle(s) for each of the following general molecules:

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Name the shape and give the $A X_{m}$ E classification and ideal bond angle(s) for each of the following general molecules:

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Determine the shape, ideal bond angle(s), and the direction of any deviation from those angles for each of the following:

(a) $\mathrm{ClO}_{2}^{-} \quad$ (b) $\mathrm{PF}_{5} \quad$ (c) $\mathrm{Se} \mathrm{F}_{4} \quad$ (d) $\mathrm{KrF}_{2}$

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Determine the shape, ideal bond angle(s), and the direction of any deviation from those angles for each of the following:

(a) $\mathrm{ClO}_{3}^{-} \quad$ (b) $\mathrm{IF}_{4}^{-} \quad$ (c) $\mathrm{SeOF}_{2} \quad$ (d) $\mathrm{TeF}_{5}^{-}$

Sam L.

Numerade Educator

Determine the shape around each central atom in each molecule, and explain any deviation from ideal bond angles:

(a) $\mathrm{CH}_{3} \mathrm{OH} \quad$ (b) $\mathrm{N}_{2} \mathrm{O}_{4}\left(\mathrm{O}_{2} \mathrm{NNO}_{2}\right)$

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Determine the shape around each central atom in each molecule, and explain any deviation from ideal bond angles:

(a) $\mathrm{H}_{3} \mathrm{PO}_{4}(\mathrm{no} \mathrm{H}-\mathrm{P} \text { bond }) \quad$ (b) $\mathrm{CH}_{3}-\mathrm{O}-\mathrm{CH}_{2} \mathrm{CH}_{3}$

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Determine the shape around each central atom in each molecule, and explain any deviation from ideal bond angles:

(a) $\mathrm{CH}_{3} \mathrm{COOH} \quad$ (b) $\mathrm{H}_{2} \mathrm{O}_{2}$

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Determine the shape around each central atom in each molecule, and explain any deviation from ideal bond angles:

(a) $\mathrm{H}_{2} \mathrm{SO}_{3}(\mathrm{no} \mathrm{H}-\mathrm{S} \text { bond }) \quad$ (b) $\mathrm{N}_{2} \mathrm{O}_{3}\left(\mathrm{ONNO}_{2}\right)$

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Arrange the following $A F_{n}$ species in order of increasing $\mathrm{F}-\mathrm{A}-\mathrm{F}$ bond angles: $\mathrm{BF}_{3}, \mathrm{BeF}_{2}, \mathrm{CF}_{4}, \mathrm{NF}_{3}, \mathrm{OF}_{2}$

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Arrange the following $\mathrm{ACl}_{n}$ species in order of decreasing $\mathrm{Cl}-\mathrm{A}-\mathrm{Cl}$ bond angles: $\mathrm{SCl}_{2}, \mathrm{OCl}_{2}, \mathrm{PCl}_{3}, \mathrm{SiCl}_{4}, \mathrm{SiCl}_{6}^{2-}$.

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State an ideal value for each of the bond angles in each molecule, and note where you expect deviations:

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Because both tin and carbon are members of Group 4A(14), they form structurally similar compounds. But tin exhibits a greater variety of structures because it forms several ionic species. Predict the shapes and ideal bond angles, including any deviations:

(a) $\mathrm{Sn}\left(\mathrm{CH}_{3}\right)_{2} \quad$ (b) $\mathrm{SnCl}_{3}^{-2} \quad$ (c) $\mathrm{Sn}\left(\mathrm{CH}_{3}\right)_{4}$

(d) $\mathrm{SnF}_{5}^{-} \quad$ (e) $\mathrm{SnF}_{6}^{2-}$

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In the gas phase, phosphorus pentachloride exists as separate molecules. In the solid phase, however, the compound is composed of alternating $\mathrm{PCl}_{4}^{+}$ and $\mathrm{PCl}_{6}-$ ions. What change(s) in molecular shape occur(s) as PCl_ solidifies? How does the $\mathrm{Cl}-\mathrm{P}-\mathrm{Cl}$ angle change?

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For molecules of general formula $\mathrm{AX}_{n}$ (where $n>2 )$ , how do you determine if a molecule is polar?

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How can a molecule with polar covalent bonds not be polar? Give an example.

Sam L.

Numerade Educator

Explain in general why the shape of a biomolecule is important to its function.

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Consider the molecules $\mathrm{SCl}_{2}, \mathrm{F}_{2}, \mathrm{CS}_{2}, \mathrm{CF}_{4},$ and $\mathrm{BrCl}$ .

(a) Which has bonds that are the most polar?

(b) Which molecules have a dipole moment?

Sam L.

Numerade Educator

Consider the molecules $\mathrm{BF}_{3}, \mathrm{PF}_{3}, \mathrm{BrF}_{3}, \mathrm{SF}_{4},$ and $\mathrm{SF}_{6}$

(a) Which has bonds that are the most polar?

(b) Which molecules have a dipole moment?

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Which molecule in each pair has the greater dipole moment? Give the reason for your choice.

(a) $\mathrm{SO}_{2}$ or $\mathrm{SO}_{3} \quad$ (b) $\mathrm{ICl}$ or $\mathrm{IF}$

(c) $\mathrm{SiF}_{4}$ or $\mathrm{SF}_{4} \quad$ (d) $\mathrm{H}_{2} \mathrm{O}$ or $\mathrm{H}_{2} \mathrm{S}$

Sam L.

Numerade Educator

Which molecule in each pair has the greater dipole moment? Give the reason for your choice.

(a) $\mathrm{ClO}_{2}$ or $\mathrm{SO}_{2} \quad$ (b) $\mathrm{HBr}$ or $\mathrm{HCl}$

(c) $\mathrm{BeCl}_{2}$ or $\mathrm{SCl}_{2} \quad$ (d) $\mathrm{AsF}_{3}$ or $\mathrm{AsF}_{5}$

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There are three different dichloroethylenes (molecular formula $C_{2} \mathrm{H}_{2} \mathrm{Cl}_{2},$ which we can designate $\mathrm{X}, \mathrm{Y},$ and $\mathrm{Z}$ . Compound $\mathrm{X}$ has no dipole moment, but compound Z does. Compounds $\mathrm{X}$ and $\mathrm{Z}$ each combine with hydrogen to give the same product:

$$\mathrm{C}_{2} \mathrm{H}_{2} \mathrm{Cl}_{2}(\mathrm{X} \text { or } \mathrm{Z})+\mathrm{H}_{2} \longrightarrow \mathrm{ClCH}_{2}-\mathrm{CH}_{2} \mathrm{Cl}$$

What are the structures of X, Y, and Z? Would you expect compound Y to have a dipole moment?

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Dinitrogen difluoride, $\mathrm{N}_{2} \mathrm{F}_{2},$ is the only stable, simple inorganic molecule with an $\mathrm{N}=\mathrm{N}$ bond. It occurs in $c i s$ and trans forms.

(a) Draw the molecular shapes of the two forms of $\mathrm{N}_{2} \mathrm{F}_{2}$ .

(b) Predict the direction of the polarity, if any, of each form.

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In addition to ammonia, nitrogen forms three other hydrides: hydrazine $\left(\mathrm{N}_{2} \mathrm{H}_{4}\right),$ diazene $\left(\mathrm{N}_{2} \mathrm{H}_{2}\right),$ and tetrazene $\left(\mathrm{N}_{4} \mathrm{H}_{4}\right) .$

(a) Use Lewis structures to compare the strength, length, and order of the nitrogen-nitrogen bonds in hydrazine, diazene, and $\mathrm{N}_{2}$.

(b) Tetrazene (atom sequence $\mathrm{H}_{2} \mathrm{NNN} \mathrm{NH}_{2} )$ decomposes above $0^{\circ} \mathrm{C}$ to hydrazine and nitrogen gas. Draw a Lewis structure for tetrazene, and calculate $\Delta H_{\mathrm{rxn}}^{\circ}$ for this decomposition.

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Draw a Lewis structure for each species: (a) $\mathrm{PF}_{5} ;(\mathrm{b}) \mathrm{CCl}_{4}$; (c) $\mathrm{H}_{3} \mathrm{O}^{+} ;$ (d) $\mathrm{ICl}_{3} ;(\mathrm{e}) \operatorname{BeH}_{2} ;(\mathrm{f}) \mathrm{PH}_{2}^{-} ;(\mathrm{g}) \operatorname{GeBr}_{4} ;(\mathrm{h}) \mathrm{CH}_{3}^{-}$;(i) $\mathrm{BCl}_{3} ;(\mathrm{j}) \mathrm{BrF}_{4}^{+} ;(\mathrm{k}) \mathrm{XeO}_{3} ;(\mathrm{l}) \mathrm{TeF}_{4}$.

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Consider the following reaction of silicon tetrafluoride:

$$\mathrm{SiF}_{4}+\mathrm{F}^{-} \longrightarrow \mathrm{SiF}_{5}^{-}$$

(a) Which depiction below best illustrates the change in molecular shape around $\mathrm{Si} ?$ (b) Give the name and $\mathrm{AX}_{m} \mathrm{E}_{n}$ designation of each shape in the depiction chosen in part (a).

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Both aluminum and iodine form chlorides, Al $_{2} \mathrm{Cl}_{6}$ and $\mathrm{I}_{2} \mathrm{Cl}_{6}$ , with "bridging" Cl atoms. The Lewis structures are

(a) What is the formal charge on each atom? (b) Which of these molecules has a planar shape? Explain.

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The VSEPR model was developed before any xenon compounds had been prepared. Thus, these compounds provided an excellent test of the model's predictive power. What would you have predicted for the shapes of $\mathrm{XeF}_{2}, \mathrm{XeF}_{4},$ and $\mathrm{XeF}_{6} ?$

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When $\mathrm{SO}_{3}$ gains two electrons, SO_{3} ^ { 2 - } forms. (a) Which depiction best illustrates the change in molecular shape around $\mathrm{S}$ ? (b) Does molecular polarity change during this reaction?

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The actual bond angle in NO $_{2}$ is $134.3^{\circ},$ and in $\mathrm{NO}_{2}^{-}$ it is $115.4^{\circ},$ although the ideal bond angle is $120^{\circ}$ for both. Explain.

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"Inert" xenon actually forms several compounds, especially with the highly electronegative elements oxygen and fluorine. The simple fluorides XeF, $\mathrm{XeF}_{4},$ and $\mathrm{XeF}_{6}$ are all formed by direct reaction of the elements. As you might expect from the size of the xenon atom, the $\mathrm{Xe}-\mathrm{F}$ bond is not a strong one. Calculate the $\mathrm{Xe}-\mathrm{F}$ bond energy in $\mathrm{XeF}_{6},$ given that the enthalpy of formation is $-402 \mathrm{kJ} / \mathrm{mol} .$

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Propylene oxide is used to make many products, including plastics such as polyurethane. One method for synthesizing it involves oxidizing propene with hydrogen peroxide:

(a) What is the molecular shape and ideal bond angle around each carbon atom in propylene oxide?

(b) Predict any deviation from the ideal for the actual $\mathrm{C}-\mathrm{C}-\mathrm{C}$

bond angles (assume the three atoms in the ring form an equilateral triangle).

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Chloral, $\mathrm{Cl}_{3} \mathrm{C}-\mathrm{CH}=\mathrm{O}$ , reacts with water to form the sedative and hypnotic agent chloral hydrate, $\mathrm{Cl}_{3} \mathrm{C}-\mathrm{CH} (\mathrm{OH})_{2} .$ Draw Lewis structures for these substances, and describe the change in molecular shape, if any, that occurs around each of the carbon atoms during the reaction.

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Like several other bonds, carbon-oxygen bonds have lengths and strengths that depend on the bond order. Draw Lewis structures for the following species, and arrange them in order of increasing carbon-oxygen bond length and then by increasing carbon-oxygen bond strength: (a) $\mathrm{CO}

(\mathrm{b}) \mathrm{CO}_{3}^{2-} ;(\mathrm{c}) \mathrm{H}_{2} \mathrm{CO} ;$ (d) $\mathrm{CH}_{4} \mathrm{O} ;(\mathrm{e}) \mathrm{HCO}_{3}^{-}(\mathrm{H}$ attached to $\mathrm{O})$ .

Katherine D.

Numerade Educator

In the $1980 s,$ there was an international agreement to destroy all stockpiles of mustard gas, $\mathrm{ClCH}_{2} \mathrm{CH}_{2} \mathrm{SCH}_{2} \mathrm{CH}_{2} \mathrm{Cl}$ . When this substance contacts the moisture in eyes, nasal passages, and skin, the - OH groups of water replace the Cl atoms and create high local concentrations of hydrochloric acid, which cause-severe blistering and tissue destruction. Write a balanced equation for this reaction, and calculate $\Delta H_{\text { rxn. }}^{\circ}$

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The four bonds of carbon tetrachloride (CCl_ are polar, but the molecule is nonpolar because the bond polarity is canceled by the symmetric tetrahedral shape. When other atoms substitute for some of the Cl atoms, the symmetry is broken and the molecule becomes polar. Use Figure 9.21$(\mathrm{p} .381)$ to rank the following molecules from the least polar to the most polar: $\mathrm{CH}_{2} \mathrm{Br}_{2}, \mathrm{CF}_{2} \mathrm{Cl}_{2}$ , $\mathrm{CH}_{2} \mathrm{F}_{2}, \mathrm{CH}_{2} \mathrm{Cl}_{2}, \mathrm{CBr}_{4}, \mathrm{CF}_{2} \mathrm{Br}_{2}$.

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Ethanol $\left(\mathrm{CH}_{3} \mathrm{CH}_{2} \mathrm{OH}\right)$ is being used as a gasoline additive or alternative in many parts of the world.

(a) Use bond energies to find $\Delta H_{\mathrm{rm}}^{\circ}$ for the combustion of gaseous ethanol. (Assume $\mathrm{H}_{2} \mathrm{O}$ forms as a gas.)

(b) In its standard state at $25^{\circ} \mathrm{C}$ , ethanol is a liquid. Its vaporization requires 40.5 $\mathrm{kJ} / \mathrm{mol}$ . Correct the value from part (a) to find the enthalpy of reaction for the combustion of liquid ethanol.

(c) How does the value from part (b) compare with the value you calculate from standard enthalpies of formation (Appendix B)?

(d) "Greener" methods produce ethanol from corn and other plant material, but the main industrial method involves hydrating ethylene from petroleum. Use Lewis structures and bond energies to calculate $\Delta H_{\mathrm{ran}}^{\circ}$ for the formation of gaseous ethanol from ethylene gas with water vapor.

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In the following compounds, the $C$ atoms form a single ring. Draw a Lewis structure for each compound, identify cases for which resonance exists, and determine the carbon-carbon bond order(s): (a) $\mathrm{C}_{3} \mathrm{H}_{4} ;(\mathrm{b}) \mathrm{C}_{3} \mathrm{H}_{6} ;(\mathrm{c}) \mathrm{C}_{4} \mathrm{H}_{6} ;(\mathrm{d}) \mathrm{C}_{4} \mathrm{H}_{4} ;(\mathrm{e}) \mathrm{C}_{6} \mathrm{H}_{6}$.

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An experiment requires 50.0 $\mathrm{mL}$ of 0.040 $\mathrm{M}$ NaOH for the titration of 1.00 $\mathrm{mmol}$ of acid. Mass analysis of the acid shows 2.24$\%$ hydrogen, 26.7$\%$ carbon, and 71.1$\%$ oxygen. Draw the Lewis structure of the acid.

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A gaseous compound has a composition by mass of 24.8$\%$ carbon, 2.08$\%$ hydrogen, and 73.1$\%$ chlorine. At STP, the gas has a density of 4.3 $\mathrm{g} / \mathrm{L}$ . Draw a Lewis structure that fits these facts. Would another structure be equally satisfactory? Explain.

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Perchlorates are powerful oxidizing agents used in fireworks, flares, and the booster rockets of space shuttles. Lewisstructures for the perchlorate ion (ClO $_{4}^{-} )$ can be drawn with all single bonds or with one, two, or three double bonds. Draw each of these possible resonance forms, use formal charges to determine the most important, and calculate its average bond order.

Sam L.

Numerade Educator

Methane burns in oxygen to form carbon dioxide and water vapor. Hydrogen sulfide burns in oxygen to form sulfur dioxide and water vapor. Use bond energies (Table $9.2, \mathrm{p} .371$ ) to determine the enthalpy of each reaction per mole of $\mathrm{O}_{2}$ (assume Lewis structures with zero formal charges; $\mathrm{EE}$ of $\mathrm{S}=\mathrm{O}$ is 552 $\mathrm{kJ} / \mathrm{mol} )$.

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Use Lewis structures to determine which two of the following are unstable: (a) $\mathrm{SF}_{2}

(\mathrm{b}) \mathrm{SF}_{3} ;(\mathrm{c}) \mathrm{SF}_{4} ;(\mathrm{d}) \mathrm{SF}_{5}

(\mathrm{e}) \mathrm{SF}_{6} .$

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A major short-lived, neutral species in flames is OH.

(a) What is unusual about the electronic structure of $\mathrm{OH}$ ?

(b) Use the standard enthalpy of formation of $\mathrm{OH}(g)(39.0 \mathrm{kJ} / \mathrm{mol})$ and bond energies to calculate the $\mathrm{O}-\mathrm{H}$ bond energy in $\mathrm{OH}(g)$.

(c) From the average value for the O- $\mathrm{H}$ bond energy in Table 9.2 $(\mathrm{p} .371)$ and your value for the $\mathrm{O}-\mathrm{H}$ bond energy in $\mathrm{OH}(g),$ find the energy needed to break the first $\mathrm{O}- \mathrm{H}$ bond in water.

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Pure $\mathrm{HN}_{3}$ ( atom sequence HNNN) is explosive. In aqueous solution, it is a weak acid that yields the azide ion, $\mathrm{N}_{3}$ . Draw resonance structures to explain why the nitrogen-nitrogen bond lengths are equal in $\mathrm{N}_{3}-$ but unequal in $\mathrm{HN}_{3}$ .

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Except for nitrogen, the elements of Group 5 $\mathrm{A}(15)$ all form pentafluorides, and most form pentachlorides. The chlorine atoms of $\mathrm{PCl}_{5}$ can be replaced with fluorine atoms one at a time to give, successively, $\mathrm{PCl}_{4} \mathrm{F}, \mathrm{PCl}_{3} \mathrm{F}_{2}, \ldots, \mathrm{PF}_{5} .$ (a) Given the sizes of $\mathrm{F}$ and $\mathrm{Cl},$ would you expect the first two $\mathrm{F}$ substitutions to be at axial or equatorial positions? Explain. (b) Which of the five fluorine-containing molecules have no dipole moment?

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Dinitrogen monoxide $\left(\mathrm{N}_{2} \mathrm{O}\right)$ supports combustion in a manner similar to oxygen, with the nitrogen atoms forming $\mathrm{N}_{2} .$ Draw three resonance structures for $\mathrm{N}_{2} \mathrm{O}$ (one $\mathrm{N}$ is central), and use formal charges to decide the relative importance of each. What correlation can you suggest between the most important structure and the observation that $\mathrm{N}_{2} \mathrm{O}$ supports combustion?

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Oxalic acid $\left(\mathrm{H}_{2} \mathrm{C}_{2} \mathrm{O}_{4}\right)$ is found in toxic concentrations in rhubarb leaves. The acid forms two ions, $\mathrm{HC}_{2} \mathrm{O}_{4}^{-}$ and $\mathrm{C}_{2} \mathrm{O}_{4}^{2-}$ by the sequential loss of $\mathrm{H}^{+}$ ions. Draw Lewis structures for the three species, and comment on the relative lengths and strengths of their carbon-oxygen bonds. The connections among the atoms are shown below with single bonds only.

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The Murchison meteorite that landed in Australia in 1969 contained 92 different amino acids, including 21 found in Earth organisms. A skeleton structure (single bonds only) of one of these extraterrestrial amino acids is shown below.

Draw a Lewis structure, and identify any atoms having a nonzero formal charge.

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Hydrazine $\left(\mathrm{N}_{2} \mathrm{H}_{4}\right)$ is used as a rocket fuel because it reacts very exothermically with oxygen to form nitrogen gas and water vapor. The heat released and the increase in number of moles of gas provide thrust. Calculate the enthalpy of reaction.

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A student isolates a product with the molecular shape shown at right (F is orange). (a) If the species is a neutral compound, can the black sphere represent selenium (Se)? (b) If the species is an anion, can the black sphere represent $\mathrm{N} ?$ (c) If the black sphere represents Br, what is the charge of the species?

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When gaseous sulfur trioxide is dissolved in concentrated sulfuric acid, disulfuric acid forms:

$$\mathrm{SO}_{3}(g)+\mathrm{H}_{2} \mathrm{SO}_{4}(l) \longrightarrow \mathrm{H}_{2} \mathrm{S}_{2} \mathrm{O}_{7}(l)$$

Use bond energies (Table $9.2, \mathrm{p} .371 )$ to determine $\Delta H_{\mathrm{rxn}^{+}}^{\circ} .$ The $\mathrm{S}$ atoms in $\mathrm{H}_{2} \mathrm{S}_{2} \mathrm{O}_{7}$ are bonded through an $\mathrm{O}$ atom. Assume Lewis

structures with zero formal charges; BE of $\mathrm{S}=\mathrm{O}$ is 552 $\mathrm{kJ} / \mathrm{mol}$ .)

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A molecule of formula $A Y_{3}$ is found experimentally to be polar. Which molecular shapes are possible and which are impossible for $A Y_{3} ?$

Sam L.

Numerade Educator

Consider the following molecular shapes:

(a) Match each shape with one of the following species: $\mathrm{XeF}_{3}^{+}$ , $\mathrm{SbBr}_{3}, \mathrm{GaCl}_{3}$ .

(b) Which, if any, is polar?

(c) Which has the most valence electrons around the central atom?

Ava P.

Numerade Educator

Hydrogen cyanide can be catalytically reduced with hydrogen to form methylamine. Use Lewis structures and bond energies to determine $\Delta H_{\text { ren }}^{\circ}$ for

$$\mathrm{HCN}(g)+2 \mathrm{H}_{2}(g) \longrightarrow \mathrm{CH}_{3} \mathrm{NH}_{2}(g)$$

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Ethylene, $\mathrm{C}_{2} \mathrm{H}_{4},$ and tetrafluoroethylene, $\mathrm{C}_{2} \mathrm{F}_{4},$ are used to make the polymers polyethylene and polytetrafluoroethylene (Teflon), respectively.

(a) Draw the Lewis structures for $\mathrm{C}_{2} \mathrm{H}_{4}$ and $\mathrm{C}_{2} \mathrm{F}_{4},$ and give the ideal $\mathrm{H}-\mathrm{C}-\mathrm{H}$ and $\mathrm{F}-\mathrm{C}-\mathrm{F}$ bond angles.

(b) The actual $\mathrm{H}-\mathrm{C}-\mathrm{H}$ and $\mathrm{F}-\mathrm{C}-\mathrm{F}$ bond angles are $117.4^{\circ}$ and $112.4^{\circ},$ respectively. Explain these deviations.

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Using bond lengths in Table 9.2$(\mathrm{p} .371)$ and assuming ideal geometry, calculate each of the following distances:

(a) Between $\mathrm{H}$ atoms in $\mathrm{C}_{2} \mathrm{H}_{2}$

(b) Between $\mathrm{F}$ atoms in $\mathrm{SF}_{6}$ (two answers)

(c) Between equatorial $\mathrm{F}$ atoms in $\mathrm{PF}_{5}$

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Phosphorus pentachloride, a key industrial compound with annual world production of about $2 \times 10^{7} \mathrm{kg},$ is used to make other compounds. It reacts with sulfur dioxide to produce phosphorus oxychloride $\left(\mathrm{POCl}_{3}\right)$ and thionyl chloride $\left(\mathrm{SOCl}_{2}\right)$ Draw a Lewis structure and name the molecular shape of each of these products.

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