Chemical Bonding and Molecular Geometry

Which of the following atoms would be expected to form negative ions in binary ionic compounds and which would be expected to form positive ions:\({\rm{P, I, Mg, Cl, In, Cs, O, Pb, Co}}\)?

Does a cation gain protons to form a positive charge or does it lose electrons?

Iron(\({\rm{III}}\)) sulfate (\({\rm{F}}{{\rm{e}}_{\rm{2}}}{{\rm{(S}}{{\rm{O}}_{\rm{4}}}{\rm{)}}_{\rm{3}}}\)) is composed of \({\rm{F}}{{\rm{e}}^{{\rm{3 + }}}}\) and \({\rm{S}}{{\rm{O}}_{\rm{4}}}^{{\rm{2 - }}}\) ions. Explain why a sample of iron(\({\rm{III}}\)) sulfate is uncharged.

Which of the following atoms would be expected to form negative ions in binary ionic compounds and which would be expected to form positive ions:\({\rm{Br, Ca, Na, N, F, Al, Sn, S, Cd}}\)?

Predict the charge on the monatomic ions formed from the following atoms in binary ionic compounds: (a) \({\rm{P}}\) (b) \({\rm{Mg}}\) (c) \({\rm{Al}}\) (d) \({\rm{O}}\) (e) \({\rm{Cl}}\) (f) \({\rm{Cs}}\).

Predict the charge on the monatomic ions formed from the following atoms in binary ionic compounds: (a) \({\rm{I}}\) (b) \({\rm{Sr}}\) (c) \({\rm{K}}\) (d) \({\rm{N}}\) (e) \({\rm{S}}\) (f) \({\rm{In}}\).

Write the electron configuration for each of the following ions: (a) \({\rm{A}}{{\rm{s}}^{{\rm{3 - }}}}\) (b) \({{\rm{I}}^{\rm{ - }}}\) (c) \({\rm{B}}{{\rm{e}}^{{\rm{2 + }}}}\) (d) \({\rm{C}}{{\rm{d}}^{{\rm{2 + }}}}\)(e) \({{\rm{O}}^{{\rm{2 - }}}}\) (f) \({\rm{G}}{{\rm{a}}^{{\rm{3 + }}}}\) (g) \({\rm{L}}{{\rm{i}}^{\rm{ + }}}\) (h) \({{\rm{N}}^{{\rm{3 - }}}}\) (i) \({\rm{S}}{{\rm{n}}^{{\rm{2 + }}}}\) (j) \({\rm{C}}{{\rm{o}}^{{\rm{2 + }}}}\) (k) \({\rm{F}}{{\rm{e}}^{{\rm{2 + }}}}\) (l) \({\rm{A}}{{\rm{s}}^{{\rm{3 + }}}}\) .

Write the electron configuration for the monatomic ions formed from the following elements (which form the greatest concentration of monatomic ions in seawater): (a) \({\rm{Cl}}\) (b) \({\rm{Na}}\) (c) \({\rm{Mg}}\)(d) \({\rm{Ca}}\) (e) \({\rm{K}}\) (f) \({\rm{Br}}\) (g) \({\rm{Sr}}\) (h) \({\rm{F}}\).

Write out the full electron configuration for each of the following atoms and for the monatomic ion found in binary ionic compounds containing the element: (a) \({\rm{Al}}\) (b) \({\rm{Br}}\) (c) \({\rm{Sr}}\) (d) \({\rm{Li}}\) (e) \({\rm{As}}\) (f) \({\rm{S}}\) .

From the labels of several commercial products, prepare a list of six ionic compounds in the products. For each compound, write the formula. (You may need to look up some formulas in a suitable reference.)

What information can you use to predict whether a bond between two atoms is covalent or ionic?

Predict which of the following compounds are ionic and which are covalent, based on the location of their constituent atoms in the periodic table: 

(a) \({\rm{C}}{{\rm{l}}_{\rm{2}}}{\rm{CO}}\) (b) \({\rm{MnO}}\) (c) \({\rm{NC}}{{\rm{l}}_{\rm{3}}}\) (d) \({\rm{CoB}}{{\rm{r}}_{\rm{2}}}\) (e) \({{\rm{K}}_{\rm{2}}}{\rm{S}}\) (f) \({\rm{CO}}\) (g) \({\rm{Ca}}{{\rm{F}}_{\rm{2}}}\) (h) \({\rm{HI}}\) (i) \({\rm{CaO}}\) (j) \({\rm{IBr}}\) (k) \({\rm{C}}{{\rm{O}}_{\rm{2}}}\) .

Explain the difference between a nonpolar covalent bond, a polar covalent bond, and an ionic bond.

From its position in the periodic table, determine which atom in each pair is more electronegative: (a) \({\rm{Br or Cl}}\) (b) \({\rm{N or O}}\) (c) \({\rm{S or O}}\) (d) \({\rm{P or S}}\) (e) \({\rm{Si or N}}\) (f) \({\rm{Ba or P}}\) (g) \({\rm{N or K}}\) .

From its position in the periodic table, determine which atom in each pair is more electronegative: (a) \({\rm{N or P}}\) (b) \({\rm{N or Ge}}\) (c) \({\rm{S or F}}\) (d) \({\rm{Cl or S}}\) (e) \({\rm{H or C}}\) (f) \({\rm{Se or P}}\) (g) \({\rm{C or Si}}\) .

From their positions in the periodic table, arrange the atoms in each of the following series in order of increasing electronegativity: (a) \({\rm{C, F, H, N, O}}\) (b) \({\rm{Br, Cl, F, H, I }}\) (c) \({\rm{F, H, O, P, S }}\) (d) \({\rm{Al, H, Na, O, P}}\) (e) \({\rm{Ba, H, N, O, As}}\) .

From their positions in the periodic table, arrange the atoms in each of the following series in order of increasing electronegativity: (a) \({\rm{As, H, N, P, Sb}}\) (b) \({\rm{Cl, H, P, S, Si}}\) (c) \({\rm{Br, Cl, Ge, H, Sr}}\) (d) \({\rm{Ca, H, K, N, Si}}\) (e) \({\rm{Cl, Cs, Ge, H, Sr}}\) .

Which atoms can bond to sulfur so as to produce a positive partial charge on the sulfur atom?

Which is the most polar bond? (a) \({\rm{C - C}}\) (b) \({\rm{C - H}}\) (c) \({\rm{N - H}}\) (d) \({\rm{O - H}}\) (e) \({\rm{Se - H}}\) .

Identify the more polar bond in each of the following pairs of bonds: (a) \({\rm{HF or HCl}}\) (b) \({\rm{NO or CO}}\) (c) \({\rm{SH or OH}}\) (d) \({\rm{PCl or HCl}}\) (e) \({\rm{CH or NH}}\) (f) \({\rm{SO or PO}}\) (g) \({\rm{CN or NN}}\) .

Which of the following molecules or ions contain polar bonds? (a) \({{\rm{S}}_{\rm{8}}}\) (b) \({{\rm{S}}_{\rm{8}}}\) (c) \({{\rm{O}}_{\rm{2}}}^{{\rm{2 - }}}\) (d) \({\rm{N}}{{\rm{O}}_{\rm{3}}}^{\rm{ - }}\) (e) \({\rm{C}}{{\rm{O}}_{\rm{2}}}\) (f) \({{\rm{H}}_{\rm{2}}}{\rm{S}}\) (g) \({\rm{B}}{{\rm{H}}_{\rm{4}}}^{\rm{ - }}\) .

Many monatomic ions are found in seawater, including the ions formed from the following list of elements. Write the Lewis symbols for the monatomic ions formed from the following elements: (a) \({\rm{CI}}\) (b) \({\rm{Na}}\) (c) \({\rm{Mg}}\) (d) \({\rm{Ca}}\) (e) \({\rm{K}}\) (f) \({\rm{Br}}\) (g) \({\rm{Sr}}\) (h) \({\rm{F}}\).

 Write the Lewis symbols of the ions in each of the following ionic compounds and the Lewis symbols of the atom from which they are formed: (a) \({\rm{MgS}}\) (b) \({\rm{A}}{{\rm{l}}_{\rm{2}}}{{\rm{O}}_{\rm{3}}}\) (c) \({\rm{GaC}}{{\rm{l}}_{\rm{3}}}\) (d) \({{\rm{K}}_{\rm{2}}}{\rm{O}}\) (e) \({\rm{L}}{{\rm{i}}_{\rm{3}}}{\rm{N}}\) (f) \({\rm{KF}}\) .

In the Lewis structures listed here, M and X represent various elements in the third period of the periodic table. Write the formula of each compound using the chemical symbols of each element:

Write the Lewis structure for the diatomic molecule \({{\rm{P}}_{\rm{2}}}\), an unstable form of phosphorus found in high temperature phosphorus vapor.

Write Lewis structures for the following: (a) \({\rm{Se}}{{\rm{F}}_{\rm{6}}}\) (b) \({\rm{Xe}}{{\rm{F}}_{\rm{4}}}\) (c) \({\rm{SeC}}{{\rm{l}}_{\rm{3}}}^{\rm{ + }}\) (d) \({\rm{C}}{{\rm{l}}_{\rm{2}}}{\rm{BBC}}{{\rm{l}}_{\rm{2}}}\) (contains a \({\rm{B - B}}\) bond).

Write Lewis structures for: (a) \({\rm{P}}{{\rm{O}}_{\rm{4}}}^{{\rm{3 - }}}\) (b) \({\rm{IC}}{{\rm{I}}_{\rm{4}}}^{\rm{ - }}\) (c) \({\rm{S}}{{\rm{O}}_{\rm{3}}}^{{\rm{2 - }}}\) (d) \({\rm{HONO}}\) .

Correct the following statement: “The bonds in solid \({\text{PbC}}{{\text{I}}_{\text{2}}}\) are ionic; the bond in a \({\text{HCl}}\) molecule is covalent. Thus, all of the valence electrons in \({\text{PbC}}{{\text{I}}_{\text{2}}}\) are located on the \({\text{C}}{{\text{I}}^{\text{ - }}}\) ions, and all of the valence electrons in a \({\text{HCl}}\) molecule are shared between the \({\text{H}}\) and \({\text{CI}}\) atoms.”

Methanol, \({{\text{H}}_{\text{3}}}{\text{COH}}\), is used as the fuel in some race cars. Ethanol, \({{\text{C}}_{\text{2}}}{{\text{H}}_{\text{5}}}{\text{OH}}\), is used extensively as motor fuel in Brazil. Both methanol and ethanol produce \({\text{C}}{{\text{O}}_{\text{2}}}\) and \({{\text{H}}_{\text{2}}}{\text{O}}\) when they burn. Write the chemical equations for these combustion reactions using Lewis structures instead of chemical formulas.

Identify the atoms that correspond to each of the following electron configurations. Then, write the Lewis symbol for the common ion formed from each atom: (a) \({\text{1}}{{\text{s}}^{\text{2}}}{\text{2}}{{\text{s}}^{\text{2}}}{\text{2}}{{\text{p}}^{\text{5}}}\) (b) \({\text{1}}{{\text{s}}^{\text{2}}}{\text{2}}{{\text{s}}^{\text{2}}}{\text{2}}{{\text{p}}^{\text{6}}}{\text{3}}{{\text{s}}^{\text{2}}}\) (c) \({\text{1}}{{\text{s}}^{\text{2}}}{\text{2}}{{\text{s}}^{\text{2}}}{\text{2}}{{\text{p}}^{\text{6}}}{\text{3}}{{\text{s}}^{\text{2}}}{\text{3}}{{\text{p}}^{\text{6}}}{\text{4}}{{\text{s}}^{\text{2}}}{\text{3}}{{\text{d}}^{{\text{10}}}}\) (d) \({\text{1}}{{\text{s}}^{\text{2}}}{\text{2}}{{\text{s}}^{\text{2}}}{\text{2}}{{\text{p}}^{\text{6}}}{\text{3}}{{\text{s}}^{\text{2}}}{\text{3}}{{\text{p}}^{\text{6}}}{\text{4}}{{\text{s}}^{\text{2}}}{\text{3}}{{\text{d}}^{{\text{10}}}}{\text{4}}{{\text{p}}^{\text{4}}}\) (e) \({\text{1}}{{\text{s}}^{\text{2}}}{\text{2}}{{\text{s}}^{\text{2}}}{\text{2}}{{\text{p}}^{\text{6}}}{\text{3}}{{\text{s}}^{\text{2}}}{\text{3}}{{\text{p}}^{\text{6}}}{\text{4}}{{\text{s}}^{\text{2}}}{\text{3}}{{\text{d}}^{{\text{10}}}}{\text{4}}{{\text{p}}^{\text{1}}}\) .

A compound with a molar mass of about \({\rm{42 g/mol}}\) contains \({\rm{85}}{\rm{.7 \% }}\) carbon and \({\rm{14}}{\rm{.3 \% }}\) hydrogen by mass. Write the Lewis structure for a molecule of the compound.

Two arrangements of atoms are possible for a compound with a molar mass of about \({\rm{45 g/mol}}\) that contains \({\rm{52}}{\rm{.2 \%  C, 13}}{\rm{.1 \%  H}}\), and \({\rm{34}}{\rm{.7 \%  O}}\) by mass. Write the Lewis structures for the two molecules.

How are single, double, and triple bonds similar? How do they differ?

Write resonance forms describing the distribution of electrons in each molecule or ion. 

a) selenium dioxide, \({\rm{OSeO}}\) 

(b) nitrate ion, \({\rm{NO}}_{\rm{3}}^{\rm{ - }}\) 

(c) nitric acid, \({\rm{HN}}{{\rm{O}}_{\rm{3}}}\) (\({\rm{N}}\) is bonded to an \({\rm{OH}}\) group and two \({\rm{O}}\) atoms)

(d) benzene, \({{\rm{C}}_{\rm{6}}}{{\rm{H}}_{\rm{6}}}\):

(e) the formate ion:

(a) sulphur dioxide, \({\rm{S}}{{\rm{O}}_{\rm{2}}}\) 

(b) carbonate ion, \({\rm{CO}}_{\rm{3}}^{{\rm{2 - }}}\)

(c) hydrogen carbonate ion, \({\rm{HCO}}_{\rm{3}}^{\rm{ - }}\) (\({\rm{C}}\) is bonded to an \({\rm{OH}}\) group and two \({\rm{O}}\) atoms) 

(d) pyridine: 

(e) the allyl ion:

Write the resonance forms of ozone, \({{\rm{O}}_{\rm{3}}}\), the component of the upper atmosphere that protects the Earth from ultraviolet radiation.

Sodium nitrite, which has been used to preserve bacon and other meats, is an ionic compound. Write the resonance forms of the nitrite ion, \({\rm{NO}}_{\rm{2}}^{\rm{ - }}\).

In terms of the bonds present, explain why acetic acid, \(C{H_3}C{O_2}H\), contains two distinct types of carbon-oxygen bonds, whereas the acetate ion, formed by loss of a hydrogen ion from acetic acid, only contains one type of carbon-oxygen bond. The skeleton structures of these species are shown:

Write the Lewis structures for the following, and include resonance structures where appropriate. Indicate which has the strongest carbon-oxygen bond. 

(a) \(C{O_2}\) 

(b) \(CO\)

Determine the formal charge of each element in the following: 

(a) \({\rm{HCl}}\) 

(b) \({\rm{C}}{{\rm{F}}_{\rm{4}}}\) 

(c) \({\rm{PC}}{{\rm{l}}_{\rm{3}}}\) 

(d) \({\rm{P}}{{\rm{F}}_{\rm{5}}}\)

Determine the formal charge of each element in the following: 

(a) \({{\rm{H}}_{\rm{3}}}{{\rm{O}}^{\rm{ + }}}\) 

(b) \({\rm{SO}}_{\rm{4}}^{{\rm{2 - }}}\) 

(c) \({\rm{N}}{{\rm{H}}_{\rm{3}}}\) 

(d) \({\rm{O}}_{\rm{2}}^{{\rm{2 - }}}\) 

(e) \({{\rm{H}}_{\rm{2}}}{{\rm{O}}_{\rm{2}}}\)

Calculate the formal charge of chlorine in the molecules \({\rm{C}}{{\rm{l}}_{\rm{2}}}\), \({\rm{BeC}}{{\rm{l}}_{\rm{2}}}\), and \({\rm{Cl}}{{\rm{F}}_{\rm{5}}}\).

Question: Calculate the formal charge of each element in the following compounds and ions: 

(a) \({{\rm{F}}_{\rm{2}}}{\rm{CO}}\) 

(b) \({\rm{N}}{{\rm{O}}^{\rm{ - }}}\) 

(c) \({\rm{BF}}_{\rm{4}}^{\rm{ - }}\) 

(d) \({\rm{SnCl}}_{\rm{3}}^{\rm{ - }}\) 

(e) \({{\rm{H}}_{\rm{2}}}{\rm{CC}}{{\rm{H}}_{\rm{2}}}\) 

(f) \({\rm{Cl}}{{\rm{F}}_{\rm{3}}}\) 

(g) \({\rm{Se}}{{\rm{F}}_{\rm{6}}}\) 

(h) \({\rm{PO}}_{\rm{4}}^{{\rm{3 - }}}\)

Question: Which bond in each of the following pairs of bonds is the strongest? 

(a) \({\rm{C - C}}\) or \({\rm{C = C}}\) 

(b) \({\rm{C - N}}\) or \({\rm{C}} \equiv {\rm{N}}\) 

(c) \({\rm{C}} \equiv {\rm{O}}\) or \({\rm{C = O}}\) 

(d) \({\rm{H - F}}\) or \({\rm{H - Cl}}\) 

(e) \({\rm{C - H}}\) or \({\rm{O - H}}\) 

(f) \({\rm{C - N}}\) or \({\rm{C - O}}\)

Question: Using the bond energies in Table \({\rm{7}}{\rm{.2}}\), determine the approximate enthalpy change for each of the following reactions: 

(a) \({{\rm{H}}_{\rm{2}}}{\rm{(g) + B}}{{\rm{r}}_{\rm{2}}}{\rm{(g)}} \to {\rm{2HBr(g)}}\) 

(b) \({\rm{C}}{{\rm{H}}_{\rm{4}}}{\rm{(g) + }}{{\rm{I}}_{\rm{2}}}{\rm{(g)}} \to {\rm{C}}{{\rm{H}}_{\rm{3}}}{\rm{I(g) + HI(g)}}\) 

(c) \({{\rm{C}}_{\rm{2}}}{{\rm{H}}_4}{\rm{(g) + 3}}{{\rm{O}}_{\rm{2}}}{\rm{(g)}} \to {\rm{2C}}{{\rm{O}}_{\rm{2}}}{\rm{(g) + 2}}{{\rm{H}}_{\rm{2}}}{\rm{O(g)}}\)

Question: Using the bond energies in Table \({\rm{7}}{\rm{.2}}\), determine the approximate enthalpy change for each of the following reactions: 

(a) \({\rm{C}}{{\rm{l}}_{\rm{2}}}{\rm{(g) + 3}}{{\rm{F}}_{\rm{2}}}{\rm{(g)}} \to {\rm{2Cl}}{{\rm{F}}_{\rm{3}}}{\rm{(g)}}\)

(b) \({{\rm{H}}_{\rm{2}}}{\rm{C = C}}{{\rm{H}}_{\rm{2}}}{\rm{(g) + }}{{\rm{H}}_{\rm{2}}}{\rm{(g)}} \to {{\rm{H}}_{\rm{3}}}{\rm{CC}}{{\rm{H}}_{\rm{3}}}{\rm{(g)}}\) 

(c) \({\rm{2}}{{\rm{C}}_{\rm{2}}}{{\rm{H}}_{\rm{6}}}{\rm{(g) + 7}}{{\rm{O}}_{\rm{2}}}{\rm{(g)}} \to {\rm{4C}}{{\rm{O}}_{\rm{2}}}{\rm{(g) + 6}}{{\rm{H}}_{\rm{2}}}{\rm{O(g)}}\)

Question: When a molecule can form two different structures, the structure with the stronger bonds is usually the more stable form. Use bond energies to predict the correct structure of the hydroxylamine molecule:

Question: How does the bond energy of \({\rm{HCl(g)}}\) differ from the standard enthalpy of formation of \({\rm{HCl(g)}}\)?

Question: Using the standard enthalpy of formation data in Appendix G, show how the standard enthalpy of formation of \({\rm{HCl(g)}}\) can be used to determine the bond energy.

Question: Using the standard enthalpy of formation data in Appendix G, calculate the bond energy of the carbon-sulphur double bond in \({\rm{C}}{{\rm{S}}_{\rm{2}}}\).