Exercises

Iron(III) sulfate [Fe₂(SO₄)₃] is composed of Fe³⁺ and ${\text{SO}}_4{}^{\text{2−}}$ ions. Explain why a sample of iron(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: 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) I

(b) Sr

(c) K

(d) N

(e) S

(f) In

Write the electron configuration for the monatomic ions formed from the following elements (which form the greatest concentration of monatomic ions in seawater):

(a) Cl

(b) Na

(c) Mg

(d) Ca

(e) K

(f) Br

(g) Sr

(h) F

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?

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) N or P

(b) N or Ge

(c) S or F

(d) Cl or S

(e) H or C

(f) Se or P

(g) 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) As, H, N, P, Sb

(b) Cl, H, P, S, Si

(c) Br, Cl, Ge, H, Sr

(d) Ca, H, K, N, Si

(e) Cl, Cs, Ge, H, Sr

Which is the most polar bond?

(a) C–C

(b) C–H

(c) N–H

(d) O–H

(e) Se–H

Which of the following molecules or ions contain polar bonds?

(a) O₃

(b) S₈

(c) ${\text{O}}_2{}^{\text{2−}}$

(d) ${\text{NO}}_3{}^{\text{−}}$

(e) CO₂

(f) H₂S

(g) ${\text{BH}}_4{}^{\text{−}}$

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) Cl

(b) Na

(c) Mg

(d) Ca

(e) K

(f) Br

(g) Sr

(h) F

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:

(a)

(b)

(c)

(d)

Write Lewis structures for the following:

(a) H₂

(b) HBr

(c) PCl₃

(d) SF₂

(e) H₂CCH₂

(f) HNNH

(g) H₂CNH

(h) NO^–

(i) N₂

(j) CO

(k) CN^–

Write Lewis structures for the following:

(a) ClF₃

(b) PCl₅

(c) BF₃

(d) ${\text{PF}}_6{}^{\text{−}}$

Write Lewis structures for:

(a) ${\text{PO}}_4{}^{\text{3−}}$

(b) ${\text{ICl}}_4{}^{\text{−}}$

(c) ${\text{SO}}_3{}^{\text{2−}}$

(d) HONO

Write Lewis structures for the following molecules or ions:

(a) SbH₃

(b) XeF₂

(c) Se₈ (a cyclic molecule with a ring of eight Se atoms)

Many planets in our solar system contain organic chemicals including methane (CH₄) and traces of ethylene (C₂H₄), ethane (C₂H₆), propyne (H₃CCCH), and diacetylene (HCCCCH). Write the Lewis structures for each of these molecules.

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) 1s²2s²2p⁵

(b) 1s²2s²2p⁶3s²

(c) 1s²2s²2p⁶3s²3p⁶4s²3d¹⁰

(d) 1s²2s²2p⁶3s²3p⁶4s²3d¹⁰4p⁴

(e) 1s²2s²2p⁶3s²3p⁶4s²3d¹⁰4p¹

A compound with a molar mass of about 28 g/mol contains 85.7% carbon and 14.3% 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 45 g/mol that contains 52.2% C, 13.1% H, and 34.7% O by mass. Write the Lewis structures for the two molecules.

Write resonance forms that describe the distribution of electrons in each of these molecules or ions.

(a) selenium dioxide, OSeO

(b) nitrate ion, ${\text{NO}}_3{}^{\text{−}}$

(c) nitric acid, HNO₃ (N is bonded to an OH group and two O atoms)

(d) benzene, C₆H₆:

(e) the formate ion:

Write the resonance forms of ozone, O₃, the component of the upper atmosphere that protects the Earth from ultraviolet radiation.

In terms of the bonds present, explain why acetic acid, CH₃CO₂H, 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:

Toothpastes containing sodium hydrogen carbonate (sodium bicarbonate) and hydrogen peroxide are widely used. Write Lewis structures for the hydrogen carbonate ion and hydrogen peroxide molecule, with resonance forms where appropriate.

Determine the formal charge of each element in the following:

(a) H₃O⁺

(b) ${\text{SO}}_4{}^{\text{2−}}$

(c) NH₃

(d) ${\text{O}}_2{}^{\text{2−}}$

(e) H₂O₂

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

(a) F₂CO

(b) NO^–

(c) ${\text{BF}}_4{}^{\text{−}}$

(d) ${\text{SnCl}}_3{}^{\text{−}}$

(e) H₂CCH₂

(f) ClF₃

(g) SeF₆

(h) ${\text{PO}}_4{}^{\text{3−}}$

Based on formal charge considerations, which of the following would likely be the correct arrangement of atoms in nitrosyl chloride: ClNO or ClON?

Based on formal charge considerations, which of the following would likely be the correct arrangement of atoms in sulfur dioxide: OSO or SOO?

Iodine forms a series of fluorides (listed here). Write Lewis structures for each of the four compounds and determine the formal charge of the iodine atom in each molecule:

(a) IF

(b) IF₃

(c) IF₅

(d) IF₇

Which of the following structures would we expect for nitrous acid? Determine the formal charges:

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

(a) C–C or $\text{C}=\text{C}$

(b) C–N or $\text{C}\equiv \text{N}$

(c)$\text{C}\equiv \text{O}$ or $\text{C}=\text{O}$

(d) H–F or H–Cl

(e) C–H or O–H

(f) C–N or C–O

Using the bond energies in Table 7.2, determine the approximate enthalpy change for each of the following reactions:

(a) ${\text{Cl}}_2(g)+3{\text{F}}_2(g)⟶2{\text{ClF}}_3(g)$

(b) ${\text{H}}_2\text{C}={\text{CH}}_2(g)+{\text{H}}_2(g)⟶{\text{H}}_3{\text{CCH}}_3(g)$

(c) $2{\text{C}}_2{\text{H}}_6(g)+7{\text{O}}_2(g)⟶4{\text{CO}}_2(g)+6{\text{H}}_2\text{O}\left(g\right)$

How does the bond energy of HCl(g) differ from the standard enthalpy of formation of HCl(g)?

Using the standard enthalpy of formation data in Appendix G, calculate the bond energy of the carbon-sulfur double bond in CS₂.

Using the standard enthalpy of formation data in Appendix G, determine which bond is stronger: the P–Cl bond in PCl₃(g) or in PCl₅(g)?

Use the bond energy to calculate an approximate value of ΔH for the following reaction. Which is the more stable form of FNO₂?

The lattice energy of LiF is 1023 kJ/mol, and the Li–F distance is 200.8 pm. NaF crystallizes in the same structure as LiF but with a Na–F distance of 231 pm. Which of the following values most closely approximates the lattice energy of NaF: 510, 890, 1023, 1175, or 4090 kJ/mol? Explain your choice.

The reaction of a metal, M, with a halogen, X₂, proceeds by an exothermic reaction as indicated by this equation: $\text{M}\left(s\right)+{\text{X}}_2(g)⟶{\text{MX}}_2(s).$ For each of the following, indicate which option will make the reaction more exothermic. Explain your answers.

(a) a large radius vs. a small radius for M⁺²

(b) a high ionization energy vs. a low ionization energy for M

(c) an increasing bond energy for the halogen

(d) a decreasing electron affinity for the halogen

(e) an increasing size of the anion formed by the halogen

Which compound in each of the following pairs has the larger lattice energy? Note: Mg²⁺ and Li⁺ have similar radii; O^2– and F^– have similar radii. Explain your choices.

(a) MgO or MgSe

(b) LiF or MgO

(c) Li₂O or LiCl

(d) Li₂Se or MgO

Which of the following compounds requires the most energy to convert one mole of the solid into separate ions?

(a) MgO

(b) SrO

(c) KF

(d) CsF

(e) MgF₂

The lattice energy of KF is 794 kJ/mol, and the interionic distance is 269 pm. The Na–F

distance in NaF, which has the same structure as KF, is 231 pm. Which of the following values is the closest approximation of the lattice energy of NaF: 682 kJ/mol, 794 kJ/mol, 924 kJ/mol, 1588 kJ/mol, or 3175 kJ/mol? Explain your answer.

What feature of a Lewis structure can be used to tell if a molecule’s (or ion’s) electron-pair geometry and molecular structure will be identical?

Why is the H–N–H angle in NH₃ smaller than the H–C–H bond angle in CH₄? Why is the H–N–H angle in ${\text{NH}}_4{}^{+}$ identical to the H–C–H bond angle in CH₄?

As a general rule, MX_n molecules (where M represents a central atom and X represents terminal atoms; n = 2 – 5) are polar if there is one or more lone pairs of electrons on M. NH₃ (M = N, X = H, n = 3) is an example. There are two molecular structures with lone pairs that are exceptions to this rule. What are they?

Identify the electron pair geometry and the molecular structure of each of the following molecules or ions:

(a) ${\text{IF}}_6{}^{+}$

(b) CF₄

(c) BF₃

(d) ${\text{SiF}}_5{}^{\text{−}}$

(e) BeCl₂

Predict the electron pair geometry and the molecular structure of each of the following ions:

(a) H₃O⁺

(b) ${\text{PCl}}_4{}^{\text{−}}$

(c) ${\text{SnCl}}_3{}^{\text{−}}$

(d) ${\text{BrCl}}_4{}^{\text{−}}$

(e) ICl₃

(f) XeF₄

(g) SF₂

Predict the electron pair geometry and the molecular structure of each of the following:

(a) IOF₅ (I is the central atom)

(b) POCl₃ (P is the central atom)

(c) Cl₂SeO (Se is the central atom)

(d) ClSO⁺ (S is the central atom)

(e) F₂SO (S is the central atom)

(f) ${\text{NO}}_2{}^{\text{−}}$

(g) ${\text{SiO}}_4{}^{\text{4−}}$

Which of these molecules and ions contain polar bonds? Which of these molecules and ions have dipole moments?

(a) H₃O⁺

(b) ${\text{PCl}}_4{}^{\text{−}}$

(c) ${\text{SnCl}}_3{}^{\text{−}}$

(d) ${\text{BrCl}}_4{}^{\text{−}}$

(e) ICl₃

(f) XeF₄

(g) SF₂

Identify the molecules with a dipole moment:

(a) SF₄

(b) CF₄

(c) Cl₂CCBr₂

(d) CH₃Cl

(e) H₂CO

The molecule XCl₂ has a dipole moment. Is X beryllium or sulfur?

There are three possible structures for PCl₂F₃ with phosphorus as the central atom. Draw them and discuss how measurements of dipole moments could help distinguish among them.

Draw the Lewis structures and predict the shape of each compound or ion:

(a) CO₂

(b) ${\text{NO}}_2{}^{\text{−}}$

(c) SO₃

(d) ${\text{SO}}_3{}^{\text{2−}}$

A molecule with the formula AB₃, in which A and B represent different atoms, could have one of three different shapes. Sketch and name the three different shapes that this molecule might have. Give an example of a molecule or ion that has each shape.

What is the molecular structure of the stable form of FNO₂? (N is the central atom.)

Use the simulation to perform the following exercises for a two-atom molecule:

(a) Adjust the electronegativity value so the bond dipole is pointing toward B. Then determine what the electronegativity values must be to switch the dipole so that it points toward A.

(b) With a partial positive charge on A, turn on the electric field and describe what happens.

(c) With a small partial negative charge on A, turn on the electric field and describe what happens.

(d) Reset all, and then with a large partial negative charge on A, turn on the electric field and describe what happens.

Use the Molecule Shape simulator to build a molecule. Starting with the central atom, click on the double bond to add one double bond. Then add one single bond and one lone pair. Rotate the molecule to observe the complete geometry. Name the electron group geometry and molecular structure and predict the bond angle. Then click the check boxes at the bottom and right of the simulator to check your answers.

Use the Molecule Shape simulator to explore real molecules. On the Real Molecules tab, select “model” mode and S₂O. What is the model bond angle? Explain whether the “real” bond angle should be larger or smaller than the ideal model angle.