91影视

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

1. ${\mathbf{SO}}_2\mathbf{or}{\mathbf{SO}}_3$
   
2. $\mathbf{ICl}\mathbf{}\mathbf{or}\mathbf{}\mathbf{IF}$
3. ${\mathbf{SiF}}_4\mathbf{or}{\mathbf{SF}}_4$
   
4. ${\mathbf{H}}_2\mathbf{O}\mathbf{or}{\mathbf{H}}_2\mathbf{S}$
   

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

1. ${\mathbf{ClO}}_2\mathbf{or}{\mathbf{SO}}_2$
   
2. HBr or HCl
3. ${\mathbf{BeCl}}_2\mathbf{or}{\mathbf{SCl}}_2$
   
4. ${\mathbf{AsF}}_3\mathbf{or}{\mathbf{AsF}}_5$
   

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

1. ${\mathbf{ClO}}_2\mathbf{or}{\mathbf{SO}}_2$
   
2. HBr or HCl
3. ${\mathbf{BeCl}}_2\mathbf{or}{\mathbf{SCl}}_2$
   
4. ${\mathbf{AsF}}_3\mathbf{or}{\mathbf{AsF}}_5$
   

Draw a Lewis structure for (a) ${\mathbf{SiF}}_4$; (b)${\mathbf{SeCl}}_2$ ; (c)${\mathbf{COF}}_2$(C is central).

Dinitrogen difluoride, ${\mathbf{N}}_2{\mathbf{F}}_2$, is the only stable, simple inorganicmolecule with an N=N bond. The compound occurs incis and trans forms.(a) Draw the molecular shapes of the two forms of${\mathbf{N}}_2{\mathbf{F}}_2$.(b) Predict the polarity, if any, of each form.

In addition to ammonia, nitrogen forms three other hydrides: hydrazine $\left({\mathbf{N}}_2{\mathbf{H}}_4\right)$, diazene$\left({\mathbf{N}}_2{\mathbf{H}}_2\right)$, and tetrazene$\left({\mathbf{N}}_4{\mathbf{H}}_4\right)$.

1. Use Lewis structures to compare the strength, length, and order of nitrogen-nitrogen bonds in hydrazine, diazene, and${\mathbf{N}}_{\mathbf{2}}$.
2. Tetrazene (atom sequence${\mathbf{H}}_2{\mathbf{NNNNH}}_2$) decomposes above${\mathbf{0}}^{\mathbf{o}}$ to hydrazine and nitrogen gas.Draw a Lewis structure for tetrazene, and calculate${\mathbf{螖贬}}_{\mathrm{rxn}}^{\mathbf{o}}$for this decomposition.

Draw a Lewis structure for each species:

$\begin{array}{l}\left(a\right){\mathbf{PF}}_5;\left(b\right){\mathbf{CCl}}_4;\text{鈥夆赌夆赌�}\left(c\right){\mathbf{H}}_3{\mathbf{O}}^{+};\left(d\right){\mathbf{ICl}}_3;\left(e\right){\mathbf{BeH}}_2;\left(f\right){\mathbf{PH}}_2^{鈭�};\\ \left(g\right){\mathbf{GeBr}}_4;\left(h\right){\mathbf{CH}}_3^{鈭�};\text{鈥夆赌�}\left(i\right){\mathbf{BCl}}_3;\left(j\right){\mathbf{BrF}}_4^{鈭�};\left(k\right){\mathbf{XeO}}_3;\left(l\right){\mathbf{TeF}}_4.\end{array}$

Give the molecular shape of each species in Problem 10.62.

$\begin{array}{l}\left(a\right){\mathbf{PF}}_5;\left(b\right){\mathbf{CCl}}_4;\text{鈥夆赌夆赌�}\left(c\right){\mathbf{H}}_3{\mathbf{O}}^{鈭�};\left(d\right){\mathbf{ICl}}_3;\left(e\right){\mathbf{BeH}}_2;\left(f\right){\mathbf{PH}}_2^{鈭�};\\ \left(g\right){\mathbf{GeBr}}_4;\left(h\right){\mathbf{CH}}_3^{鈭�};\text{鈥夆赌�}\left(i\right){\mathbf{BCl}}_3;\left(j\right){\mathbf{BrF}}_4^{鈭�};\left(k\right){\mathbf{XeO}}_3;\left(l\right){\mathbf{TeF}}_4.\end{array}$

Consider the following reaction of silicon tetrafluoride: ${\mathbf{SiF}}_4\text{+}{\mathbf{F}}^{鈭�}\text{鈥�}\stackrel{}{鈫�}{\mathbf{SiF}}_5^{鈭�}$

1. Which depiction below best illustrates the change in molecular shape around Si?
2. Give the name and ${\mathbf{AX}}_m{\mathbf{E}}_n$ designation of each shape in the depiction chosen in part (a).

Both aluminium and iodine form chlorides, ${\mathbf{Al}}_2{\mathbf{Cl}}_6$and ${\mathbf{l}}_2{\mathbf{Cl}}_6$, with 鈥渂ridging鈥� Cl atoms. The Lewis structures are

1. What is the formal charge on each atom?
2. Which of these molecules has a planar shape? Explain.