Models of Chemical Bonding

(a) List four physical characteristics of a solid metal. 

(b) List two chemical characteristics of a metallic element.

carbon- carbon bonds form the “backbone” of nearly every organic and biological molecule. The average bond energy of the C-C bond is 347kJ/mol. Calculate the frequency and wavelength of the least energetic photon that can break an average C-C bond. In what region of the electromagnetic spectrum is this radiation?

Acetylene gas (ethyne; HC≡CH) burns in an oxyacetylene torch to produce carbon dioxide and water vapour. The heat of reaction for the combustion of acetylene is 1259 kJ/mol. 

(a) Calculate the C≡C bond energy, and compare your value with that in Table 9.2, p. 353. 

(b) When 500.0 g of acetylene burns, how many kilojoules of heat are given off? 

(c) How many grams of ${\mathbf{CO}}_2$form? 

(d) How many litres of  ${\mathbf{O}}_{\mathbf{2}}$ at 298 K and 18.0 atm are consumed?

Use Lewis electron-dot symbols to represent the formation of 

(a) ${\mathbf{BrF}}_3$ from bromine and fluorine atoms; 

(b) ${\mathbf{AlF}}_3$ from aluminium and fluorine atoms. 

How does electronegativity differ from electron affinity?   

How is the partial ionic character of a bond in a diatomic molecule related to $\mathbf{\text{ΔEN}}$  for the bonded atoms? Why?

Is the H-O bond in water nonpolar covalent, polar-covalent, or ionic? Define each term, and explain your choice. 

Using the periodic table only, arrange the elements in each set-in order of increasing EN: 

(a) S, O, Si; 

(b) Mg, P, As.

Using the periodic table only, arrange the elements in each set-in order of decreasing EN: 

(a) N, P, Si; 

(b) Ca, Ga, As.

Use Figure 9.20, p. 364, to indicate the polarity of each bond with partial charges

(a) Br-Cl

(b) F-Cl

(c) H-O

(d) Se-H

(e) As-H

(f) S-N.

Which is the more polar bond in each of the following pairs from Problem 9.61: (a) or (b); (c) or (d); (e) or (f)?

Even though so much energy is required to form a metal cation with a 2+ charge, the alkaline earth metals form halides with general formula 

 ${\mathbf{MX}}_2$, rather than MX. 

(a) Use the following data to calculate the  of MgCl: 

Mg(s)                      Mg(g)  = 148 kJ 

${\mathbf{Cl}}_2$  (g)          2Cl(g)  -243 kJ 

Mg(g)            (g) + e^-  =738 kJ 

Cl(g) +        (g)  = -349 kJ

   of MgCl = 783.5 kJ/mol.

(b) Is MgCl favoured energetically relative to Mg and  ? Explain. 

(c) Use Hess’s law to calculate ∆H° for the conversion of MgCl to   and Mg ( of   = -641.6 kJ/mol). 

(d) Is MgCl favoured energetically relative to  ? Explain.

How much energy is released in the formation of one molecule of HCl by the following reaction? ${\mathbf{H}}^{+}\mathbf{\left(}\mathbf{g}\mathbf{\right)}\mathbf{+}\mathbf{C}{\mathbf{l}}^{-}\mathbf{\left(}\mathbf{g}\mathbf{\right)}\to \mathbf{HCl}\mathbf{\left(}\mathbf{g}\mathbf{\right)}$  The bond energy of HCl is 431 kJ mol^-1. Additional data can be found in tables in Chapter 8?

By using photons of specific wavelengths, chemists can dissociate gaseous Hi to produce H atoms at certain speeds. When HI dissociates, the H atoms move away rapidly. Whereas the heavier I atoms move away more slowly.

1. What is the longest wavelength (in nm) that can dissociate a molecule of HI?
2. If a photon of 254 nm is used, what is the excess energy (in J) over that needed for dissociation?
3. If the excess energy is carried away by the H atom as kinetic energy, what is the speed (in m/s)?

Linear, triatomic ${\mathbf{CO}}_2$ vibrates by symmetric stretch, bend, and asymmetric stretch with frequencies of $\mathbf{4}\mathbf{.}\mathbf{02}\mathbf{\times }{\mathbf{10}}^{13}{\mathbf{s}}^{-1}$,   $\mathbf{2}\mathbf{.}\mathbf{00}\mathbf{\times }{\mathbf{10}}^{13}{\mathbf{s}}^{-1}$ and $\mathbf{7}\mathbf{.}\mathbf{05}\mathbf{\times }{\mathbf{10}}^{13}{\mathbf{s}}^{-1}$ respectively.

1. In what region of the electromagnetic spectrum are these frequencies?
2. calculate the energy (in J) of each vibration. Which takes the least energy?

What is the general relationship between   and EN for the elements? Why?             

Using the periodic table only, arrange the elements in each set in order of increasing EN: (a) I, Br, N; (b) Ca, H, F

Using the periodic table only, arrange the elements in each set in order of decreasing EN: (a) Br, Cl, P; (b) I, F, O

Use Figure 9.20, p. 364, to indicate the polarity of each bond with a polar arrow:

(a) N-B

(b) N-O

(c) C-S

(d) S-O

(e) N-H

(f) Cl-O

Which is the more polar bond in each of the following pairs from Problem 9.62: (a) or (b); (c) or (d); (e) or (f)?

Are the bonds in each of the following substances ionic, nonpolar covalent, or polar covalent? Arrange the substances with polar covalent bonds in order of increasing bond polarity: 

(a) S₈  

(b) RbCl 

(c) PF₃

(d) SCI₂

(e) F₂

(f) SF₂

Are the bonds in each of the following substances ionic, nonpolar covalent, or polar covalent? Arrange the substances with polar covalent bonds in order of increasing bond polarity: 

(a) KCl 

(b) P₄

(c) BE₃

(d) SO₂

(e) Br₂

(f) NO₂

Rank the members of each set of compounds in order of increasing the ionic character of their bonds. Use polar arrows to indicate the bond polarity of each:

(a) HBr, HCl, HI

(b) H₂O, CH₄ , HF

(c) SCI₂ , PCI₃ , SiCI₄ 

Rank the members of each set of compounds in order of increasing the ionic character of their bonds. Use polar arrows to indicate the bond polarity of each:

(a) PCI₃ , PBr₃ , PI₃

(b) BF₃ , NF₃ , CF₄

(c) SeF₄ , TeF₄ , BrF₃

The energy of the C-C bond is 347 kJ/mol, and that of the Cl-Cl bond is 243 kJ/mol. Which of the following values might you expect for the C-Cl bond energy? Explain. 

(a) 590 kJ/mol (sum of the values given) 

(b) 104 kJ/mol (difference of the values given)

(c) 295 kJ/mol (average of the values given) 

(d) 339 kJ/mol (greater than the average of the values given)

Briefly account for the following relative values: 

(a) The melting points of Na and K are 89^oC and 63^oC , respectively. 

(b) The melting points of Li and Be are 180^oC and 1287^oC , respectively. 

(c) Li boils more than 1100^oC higher than it melts.

Magnesium metal is easily deformed by an applied force, whereas magnesium fluoride is shattered. Why do these two solids behave so differently?

Geologists have a rule of thumb: when molten rock cools and solidifies, crystals of compounds with the smallest lattice energies appear at the bottom of the mass. Suggest a reason for this.

Heats of reaction calculated from bond energies and from heats of formation are often, but not always, close to each other.

a) Industrial ethanol (${\mathbf{CH}}_3{\mathbf{CH}}_2\mathbf{OH}$ ) is produce by a catalytic reaction of ethylene (${\mathbf{H}}_2\mathbf{C}\mathbf{=}{\mathbf{CH}}_2$ ) with water at high pressure and temperatures. Calculate ${{\mathbf{\Delta H}}^o}_{\mathrm{rx}}$  for this gas-phase hydration of ethylene to ethanol, using bond energies and then using heats of formation.

b) ethylene glycol is produced by the catalytic oxidation of ethylene to ethylene oxide, which then reacts with water to form ethylene glycol:

The ${{\mathbf{\Delta H}}^o}_{\mathrm{rx}}$  for this hydrolysis step, based on heat of formation, is -97kJ/mol. Calculate ${{\mathbf{\Delta H}}^o}_{\mathrm{rx}}$  for the hydrolysis using bond energies.

c) why are two values relatively close for the hydration in part (a) but not close for the hydrolysis in part(b).

Which of the bonds in Problem 9.84 is the most polar?

Which of the bonds in the Problem 9.85 is the least polar?

What would be the formula for the simplest compound formed from (a) phosphorus and chlorine, (b) carbon and fluorine, and (c) iodine and chlorine? 

Use the octet rule to predict the formula of the simplest compound formed from hydrogen and (a) selenium, (b) arsenic, and (c) silicon. (Remember that the valence shell of hydrogen can hold only two electrons.)

In the developing concept of electronegativity, Pauling used the term excess bond energy for the difference between the actual bond energy X-Y and the average bond energies of X-X and Y-Y (see text discussion for the case of HF). Based on the values in figure 9.20, p. 364, which of the following substances contains bonds with no excess bond energy?

(a) ${\mathbf{PH}}_3$        (b) ${\mathbf{CS}}_2$         (c) $\mathbf{BrCl}$          (d)  ${\mathbf{BH}}_3$        (e)  ${\mathbf{Se}}_8$

Used condensed electron configuration to predict the relative hardness and melting points of Rubidium (Z=37), Vanadium (Z=23) and Cadmium (Z= 48).

The HF bond length is 92pm, 16% shorter than the sum of the covalent radii of H (37pm) and F(72pm). Suggest a reason for this difference. Similar calculations show that the difference becomes smaller down to the group from HF and HI. Explain

There are two main types of covalent bond breakage. In homolytic breakage, each atom in the bond gets one of the shared electrons. In some cases, the electronegativity of adjacent atoms affects the bond energy. In heterolytic breakage, one atom gets both electrons and the other gets none; thus, a cation and an anion form.

a) Why is C-C bond in ${\mathbf{H}}_3\mathbf{C}\mathbf{-}{\mathbf{CF}}_3$  (423kJ/mol) stronger than that in  ${\mathbf{H}}_3\mathbf{C}\mathbf{-}{\mathbf{CH}}_3$ 

    (376kJ/mol)?

b) use bond energy and any other data to calculate the heat of reaction for the heterolytic cleavage of ${\mathbf{O}}_2$ . 

Lattice energies can also be calculated for covalent solids using a Born-Haber cycle, and the network solid silicon dioxide has one of the highest  ${{\mathbf{\Delta H}}^o}_{\mathrm{Lattice}}$ values. Silicon oxide is found in pure crystalline form as transparent rock quartz. Much harder than glass, this material was one prized for making lenses for optical devices and expensive spectacles. Use appendix B and the following data to calculate ${{\mathbf{\Delta H}}^o}_{\mathrm{Lattice}}$  of  ${\mathbf{SiO}}_2$ :

  _{$\begin{array}{l}\mathbf{Si}\mathbf{\left(}\mathbf{s}\mathbf{\right)}\to \mathbf{Si}\mathbf{\left(}\mathbf{g}\mathbf{\right)}{\mathbf{\Delta H}}^o\mathbf{=}\mathbf{454}\mathbf{kJ}\\ \mathbf{Si}\mathbf{\left(}\mathbf{g}\mathbf{\right)}\to {\mathbf{Si}}^{4+}\mathbf{\left(}\mathbf{g}\mathbf{\right)}\mathbf{+}\mathbf{4}{\mathbf{e}}^{-}{\mathbf{\Delta H}}^o\mathbf{=}\mathbf{9949}\mathbf{kJ}\\ {\mathbf{O}}_2\mathbf{\left(}\mathbf{g}\mathbf{\right)}\to \mathbf{2}\mathbf{O}\mathbf{\left(}\mathbf{g}\mathbf{\right)}{\mathbf{\Delta H}}^o\mathbf{=}\mathbf{498}\mathbf{kJ}\\ \mathbf{O}\mathbf{\left(}\mathbf{g}\mathbf{\right)}\mathbf{+}\mathbf{2}{\mathbf{e}}^{-}\to {\mathbf{O}}^{2-}\mathbf{\left(}\mathbf{g}\mathbf{\right)}\text{\hspace{0.33em}\hspace{0.33em}\hspace{0.33em}\hspace{0.33em}\hspace{0.33em}\hspace{0.33em}\hspace{0.33em}\hspace{0.33em}}{\mathbf{\Delta H}}^o\mathbf{=}\mathbf{737}\mathbf{kJ}\end{array}$}

The average C-H bond energy in ${\mathbf{CH}}_4$   is 415kJ/mol. Use table 9.2(p. 353) and the following to calculate the average C-H bond energy in ethane (${\mathbf{C}}_2{\mathbf{H}}_6$ ;C-C bond), in ethene (${\mathbf{C}}_2{\mathbf{H}}_4$ ;C=C bond), and ethyne (${\mathbf{C}}_2{\mathbf{H}}_2\mathbf{;}\mathbf{C}\equiv \mathbf{C}$ ): 

$\begin{array}{l}{\mathbf{C}}_2{\mathbf{H}}_6\mathbf{\left(}\mathbf{g}\mathbf{\right)}\mathbf{+}{\mathbf{H}}_2\mathbf{\left(}\mathbf{g}\mathbf{\right)}\to \mathbf{2}{\mathbf{CH}}_4\mathbf{\left(}\mathbf{g}\mathbf{\right)}{\mathbf{\Delta H}}^o\mathbf{=}\mathbf{-}\mathbf{65}\mathbf{.}\mathbf{07}\mathbf{kJ}\mathbf{/}\mathbf{mol}\\ {\mathbf{C}}_2{\mathbf{H}}_4\mathbf{\left(}\mathbf{g}\mathbf{\right)}\mathbf{+}\mathbf{2}{\mathbf{H}}_2\mathbf{\left(}\mathbf{g}\mathbf{\right)}\to \mathbf{2}{\mathbf{CH}}_4\mathbf{\left(}\mathbf{g}\mathbf{\right)}{\mathbf{\Delta H}}^o\mathbf{=}\mathbf{-}\mathbf{202}\mathbf{kJ}\mathbf{/}\mathbf{mol}\\ {\mathbf{C}}_2{\mathbf{H}}_2\mathbf{\left(}\mathbf{g}\mathbf{\right)}\mathbf{+}\mathbf{3}{\mathbf{H}}_2\mathbf{\left(}\mathbf{g}\mathbf{\right)}\to \mathbf{2}{\mathbf{CH}}_4\mathbf{\left(}\mathbf{g}\mathbf{\right)}{\mathbf{\Delta H}}^o\mathbf{=}\mathbf{-}\mathbf{376}\mathbf{.}\mathbf{74}\mathbf{kJ}\mathbf{/}\mathbf{mol}\end{array}$ 

Dimethyl ether (${\mathbf{CH}}_3{\mathbf{OCH}}_3$ ) and ethanol (${\mathbf{CH}}_3{\mathbf{CH}}_2\mathbf{OH}$ ) are constitutional isomers.

a) Calculate ${{\mathbf{\Delta H}}^o}_{\mathrm{rx}}$  for the formation of each compound as a gas from methane and oxygen; water vapor forms. 

b) state which reaction is more exothermic.

c) calculate ${{\mathbf{\Delta H}}^o}_{\mathrm{rx}}$  for the conversion of ethanol to dimethyl ether

In a future hydrogen-fuel economy, the cheapest source of ${\mathbf{H}}_2$will certainly be water. It takes467kJ to produce one mol of H atoms from water. What is the frequency, wavelength and minimum energy of a photon that can free an H atom from water?