Chapter 10

Why are bonding theories important? Provide some examples of what bonding theories can predict.

Why do chemical bonds form? What basic forces are involved in bonding?

What are the three basic types of chemical bonds? What happens to electrons in the bonding atoms in each type?

Describe the octet rule in the Lewis model.

According to the Lewis model, what is a chemical bond?

Why is the formation of solid sodium chloride from solid sodium and gaseous chlorine exothermic, even though it takes more energy to form the \(\mathrm{Na}^{+}\) ion than the amount of energy released upon formation of \(\mathrm{Cl}^{-} ?\)

How does lattice energy relate to ionic radii? To ion charge?

How does the ionic bonding model explain the relatively high melting points of ionic compounds?

How does the ionic bonding model explain the nonconductivity of ionic solids, and at the same time the conductivity of ionic solutions?

In a covalent Lewis structure, what is the difference between lone pair and bonding pair electrons?

In what ways are double and triple covalent bonds different from single covalent bonds?

How does the Lewis model for covalent bonding account for the relatively low melting and boiling points of molecular compounds (compared to ionic compounds)?

What is electronegativity? What are the periodic trends in electronegativity?

How do a pure covalent bond, a polar covalent bond, and an ionic bond differ?

Explain percent ionic character of a bond. Do any bonds have \(100 \%\) ionic character?

What is a dipole moment?

What is the magnitude of the dipole moment formed by separating a proton and an electron by \(100 \mathrm{pm} ? 200 \mathrm{pm} ?\)

What is the basic procedure for writing a covalent Lewis structure?

How do you determine the number of electrons in the Lewis structure of a molecule? A polyatomic ion?

What are resonance structures? What is a resonance hybrid?

Do resonance structures always contribute equally to the overall structure of a molecule? Explain.

What is formal charge? How is formal charge calculated? How is it helpful?

Why does the octet rule have exceptions? List the three major categories of exceptions and an example of each.

What is bond energy? How can you use average bond energies to calculate enthalpies of reaction?

Explain the difference between endothermic reactions and exothermic reactions with respect to the bond energies of the bonds broken and formed.

How does the electron sea model explain the conductivity of metals? The malleability and ductility of metals?

Write the electron configuration for \(\mathrm{N}\). Then write the Lewis symbol for \(\mathrm{N}\) and show which electrons from the electron configuration are included in the Lewis symbol.

Write the electron configuration for Ne. Then write the Lewis symbol for Ne and show which electrons from the electron configuration are included in the Lewis symbol.

Write the Lewis symbol for each atom or ion. a. \(A l\) b. \(\mathrm{Na}^{+}\) c. Cl d.\(\mathrm{Cl}^{-}\)

Write the Lewis symbol for each atom or ion. a. \(s^{2-}\) b. \(\mathrm{Mg}\) c. \(\mathrm{Mg}^{2+}\) d. \(P\)

Write the Lewis symbols for the lons in each ionic compound. a. NaF b. CaO c. \(\mathrm{SrBr}_{2}\) d. \(\mathrm{K}_{2} \mathrm{O}\)

Explain the trend in the lattice energies of the alkaline earth metal oxides. MISSED THIS? Read Section 10.4 $$ \begin{array}{lc} \text { Metal Oxide } & \text { Lattice Energy (k } \mathrm{J} / \mathrm{mol} \text { ) } \\ \hline \mathrm{MgO} & -3795 \\ \hline \mathrm{CaO} & -3414 \\ \hline \mathrm{SrO} & -3217 \\ \hline \mathrm{BaO} & -3029 \\ \hline \end{array} $$

Rubidium iodide has a lattice energy of \(-617 \mathrm{kj} / \mathrm{mol},\) while potassium bromide has a lattice energy of \(-671 \mathrm{~kJ} / \mathrm{mol}\). Why is the lattice energy of potassium bromide more exothermic than the lattice energy of rubidium iodide?

Use covalent Lewis structures to explain why the compound that forms between nitrogen and hydrogen has the formula \(\mathrm{NH}_{3}\). Show why \(\mathrm{NH}_{2}\) and \(\mathrm{NH}_{4}\) are not stable.

Write the Lewis structure for each molecule. MiSSED THIS? Read Section 10.7; Watch KCV 10.7, ME 10.4 a. \(P H_{3}\) b. \(\mathrm{SCl}_{2}\) c. HI d. \(\mathrm{CH}_{4}\)

Write the Lewis structure for each molecule. a. \(\mathrm{NF}_{3}\) b. HBr c. SBr \(_{2}\) d. \(\mathrm{CCl}_{4}\)

Write the Levis structure for each molecule. MISSED THIST Read Section 10.7; Watch KCV 10.7, IWE 10.4 a. \(\mathrm{SF}_{2}\) b. \(\mathrm{SiH}_{4}\) c. HCOOH (both O bonded to C) d. \(\mathrm{CH}_{3} \mathrm{SH}\) (C and \(\mathrm{S}\) central)

Write the Lewis structure for each molecule. a. \(\mathrm{CH}_{2} \mathrm{O}\) b. \(\mathrm{C}_{2} \mathrm{Cl}_{4}\) c. \(\mathrm{CH}_{3} \mathrm{NH}_{2}\) d. CFCl \(_{3}\) (C central)

Determine if a bond between each pair of atoms would be pure covalent, polar covalent, or ionic. a. Cand N b. N and S c. Kand F d. N and \(N\)

Write the Lewis structure for each molecule or ion. a. \(\mathrm{N}_{2} \mathrm{H}_{2}\) b. \(\mathrm{N}_{2} \mathrm{H}_{4}\) c. \(\mathrm{C}_{2} \mathrm{H}_{2}\) d. \(\mathrm{C}_{2} \mathrm{H}_{4}\)

Write the Lewis structure for each molecule or ion. a. \(\mathrm{H}_{3} \mathrm{COCH}_{3}\) b. CN c. \(\mathrm{NO}_{2}^{-}\) d. \(\mathrm{ClO}^{-}\)

Draw the Lewis structure (including resonance structures) for the acetate lon \(\left(\mathrm{CH}_{3} \mathrm{COO}^{-}\right)\). For each resonance structure, assign formal charges to all atoms that have formal charge.

Write the Lewis structure for each ion. Include resonance structures if necessary and assign formal charges to all atoms. If necessary, expand the octet on the central atom to lower formal charge. MISSED THIS? Road Soction 10.9 ; Watch KCV \(10.9,\) HE 10.10 a. \(\mathrm{PO}_{4}^{3-}\) b. CN" c. \(\$ O_{3}^{2-}\) d. \(\mathrm{ClO}_{2}\)

Write Lewis structures for each molecule or ion. Use expanded octets as necessary. a. PIs b. \(I_{3}^{-}\) c. \(\mathrm{SF}_{4}\) d. GeF \(_{4}\)

Write Lewis structures for each molecule or ion. Use expanded octets as necessary. a. ClFs b. \(A \mathrm{sF}_{6}\) c. \(\mathrm{Cl}_{3} \mathrm{PO}\) d. IF \(_{5}\)

Ethane burns in alr to form carbon dixode and water vapor. MISSED THIS? Read Section 10.10 ; Watch ME 10.11 $$ 2 \mathrm{H}_{3} \mathrm{C}-\mathrm{CH}_{3}(\mathrm{x})+7 \mathrm{O}_{2}(\mathrm{~s}) \longrightarrow 4 \mathrm{CO}_{2}(g)+6 \mathrm{H}_{2} \mathrm{O}(\mathrm{s}) $$

Amino acids are the building blocks of proteins. The simplest amino acid is glycine \(\left(\mathrm{H}_{2} \mathrm{NCH}_{2} \mathrm{COOH}\right)\). Draw a Lewis structure for glycine. (Hint: The central atoms in the skeletal structure are nitrogen and the two carbon atoms. Each oxygen atom is bonded directly to the right-most carbon atom.)

The reaction of \(\mathrm{Fe}_{2} \mathrm{O}_{3}(\mathrm{~s})\) with \(\mathrm{Al}(\mathrm{s})\) to form \(\mathrm{Al}_{2} \mathrm{O}_{3}(\mathrm{~s})\) and \(\mathrm{Fe}(\mathrm{s})\) is called the thermite reaction and is highly exothermic. What role does lattice energy play in the exothermicity of the reaction?

Draw the Lewis structure for nitric acid (the hydrogen atom is attached to one of the oxygen atoms). Include all three resonance structures by alternating the double bond among the three oxygen atoms. Use formal charge to determine which of the resonance structures is most important to the structure of nitric acid.

Draw the Lewis structure for each organic compound from its condensed structural formula. a. \(\mathrm{C}_{3} \mathrm{H}_{8}\) b. \(\mathrm{CH}_{3} \mathrm{OCH}_{3}\) c. \(\mathrm{CH}_{3} \mathrm{COCH}_{3}\) d. \(\mathrm{CH}_{3} \mathrm{COOH}\) e. \(\mathrm{CH}_{3} \mathrm{CHO}\)