Exercises

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

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: P, I, Mg, Cl, In, Cs, O, Pb, Co?

Predict the charge on the monatomic ions formed from the following atoms in binary ionic compounds:

(a) P

(b) Mg

(c) Al

(d) O

(e) Cl

(f) Cs

Write the electron configuration for each of the following ions:

(a) As^3–

(b) I^–

(c) Be²⁺

(d) Cd²⁺

(e) O^2–

(f) Ga³⁺

(g) Li⁺

(h) N^3–

(i) Sn²⁺

(j) Co²⁺

(k) Fe²⁺

(l) As³⁺

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) Al

(b) Br

(c) Sr

(d) Li

(e) As

(f) S

Why is it incorrect to speak of a molecule of solid NaCl?

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) Cl₂CO

(b) MnO

(c) NCl₃

(d) CoBr₂

(e) K₂S

(f) CO

(g) CaF₂

(h) HI

(i) CaO

(j) IBr

(k) CO₂

From its position in the periodic table, determine which atom in each pair is more electronegative:

(a) Br or Cl

(b) N or O

(c) S or O

(d) P or S

(e) Si or N

(f) Ba or P

(g) N or K

From their positions in the periodic table, arrange the atoms in each of the following series in order of increasing electronegativity:

(a) C, F, H, N, O

(b) Br, Cl, F, H, I

(c) F, H, O, P, S

(d) Al, H, Na, O, P

(e) Ba, H, N, O, As

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

Identify the more polar bond in each of the following pairs of bonds:

(a) HF or HCl

(b) NO or CO

(c) SH or OH

(d) PCl or SCl

(e) CH or NH

(f) SO or PO

(g) CN or NN

Name the following compounds:

(a) CsCl

(b) BaO

(c) K₂S

(d) BeCl₂

(e) HBr

(f) AlF₃

Write the formulas of the following compounds:

(a) rubidium bromide

(b) magnesium selenide

(c) sodium oxide

(d) calcium chloride

(e) hydrogen fluoride

(f) gallium phosphide

(g) aluminum bromide

(h) ammonium sulfate

Write the formulas of the following compounds:

(a) chlorine dioxide

(b) dinitrogen tetraoxide

(c) potassium phosphide

(d) silver sulfide

(e) aluminum fluoride trihydrate

(f) silicon dioxide

Each of the following compounds contains a metal that can exhibit more than one ionic charge. Name these compounds:

(a) Cr₂O₃

(b) FeCl₂

(c) CrO₃

(d) TiCl₄

(e) CoCl₂∙6H₂O

(f) MoS₂

The following ionic compounds are found in common household products. Write the formulas for each compound:

(a) potassium phosphate

(b) copper(II) sulfate

(c) calcium chloride

(d) titanium(IV) oxide

(e) ammonium nitrate

(f) sodium bisulfate (the common name for sodium hydrogen sulfate)

What are the IUPAC names of the following compounds?

(a) manganese dioxide

(b) mercurous chloride (Hg₂Cl₂)

(c) ferric nitrate [Fe(NO₃)₃]

(d) titanium tetrachloride

(e) cupric bromide (CuBr₂)

Write the Lewis symbols for each of the following ions:

(a) As^3–

(b) I^–

(c) Be²⁺

(d) O^2–

(e) Ga³⁺

(f) Li⁺

(g) N^3–

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) MgS

(b) Al₂O₃

(c) GaCl₃

(d) K₂O

(e) Li₃N

(f) KF

Write the Lewis structure for the diatomic molecule P₂, an unstable form of phosphorus found in high-temperature phosphorus vapor.

Write Lewis structures for the following:

(a) O₂

(b) H₂CO

(c) AsF₃

(d) ClNO

(e) SiCl₄

(f) H₃O⁺

(g) ${\text{NH}}_4{}^{+}$

(h) ${\text{BF}}_4{}^{\text{−}}$

(i) HCCH

(j) ClCN

(k) ${\text{C}}_2{}^{\text{2+}}$

Write Lewis structures for the following:

(a) SeF₆

(b) XeF₄

(c) ${\text{SeCl}}_3{}^{+}$

(d) Cl₂BBCl₂ (contains a B–B bond)

Correct the following statement: “The bonds in solid PbCl₂ are ionic; the bond in a HCl molecule is covalent. Thus, all of the valence electrons in PbCl₂ are located on the Cl^– ions, and all of the valence electrons in a HCl molecule are shared between the H and Cl atoms.”

Methanol, H₃COH, is used as the fuel in some race cars. Ethanol, C₂H₅OH, is used extensively as motor fuel in Brazil. Both methanol and ethanol produce CO₂ and H₂O when they burn. Write the chemical equations for these combustion reactions using Lewis structures instead of chemical formulas.

Carbon tetrachloride was formerly used in fire extinguishers for electrical fires. It is no longer used for this purpose because of the formation of the toxic gas phosgene, Cl₂CO. Write the Lewis structures for carbon tetrachloride and phosgene.

The arrangement of atoms in several biologically important molecules is given here. Complete the Lewis structures of these molecules by adding multiple bonds and lone pairs. Do not add any more atoms.

(a) the amino acid serine:

(b) urea:

(c) pyruvic acid:

(d) uracil:

(e) carbonic acid:

A compound with a molar mass of about 42 g/mol contains 85.7% carbon and 14.3% hydrogen by mass. Write the Lewis structure for a molecule of the compound.

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

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

(a) sulfur dioxide, SO₂

(b) carbonate ion, ${\text{CO}}_3{}^{\text{2−}}$

(c) hydrogen carbonate ion, ${\text{HCO}}_3{}^{\text{−}}$ (C is bonded to an OH group and two O atoms)

(d) pyridine:

(e) the allyl ion:

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

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

(a) CO₂

(b) CO

Determine the formal charge of each element in the following:

(a) HCl

(b) CF₄

(c) PCl₃

(d) PF₅

Calculate the formal charge of chlorine in the molecules Cl₂, BeCl₂, and ClF₅.

Draw all possible resonance structures for each of these compounds. Determine the formal charge on each atom in each of the resonance structures:

(a) O₃

(b) SO₂

(c) ${\text{NO}}_2{}^{\text{−}}$

(d) ${\text{NO}}_3{}^{\text{−}}$

Based on formal charge considerations, which of the following would likely be the correct arrangement of atoms in hypochlorous acid: HOCl or OClH?

Draw the structure of hydroxylamine, H₃NO, and assign formal charges; look up the structure. Is the actual structure consistent with the formal charges?

Write the Lewis structure and chemical formula of the compound with a molar mass of about 70 g/mol that contains 19.7% nitrogen and 80.3% fluorine by mass, and determine the formal charge of the atoms in this compound.

Sulfuric acid is the industrial chemical produced in greatest quantity worldwide. About 90 billion pounds are produced each year in the United States alone. Write the Lewis structure for sulfuric acid, H₂SO₄, which has two oxygen atoms and two OH groups bonded to the sulfur.

Explain why the HOH molecule is bent, whereas the HBeH molecule is linear.

Explain the difference between electron-pair geometry and molecular structure.

Explain how a molecule that contains polar bonds can be nonpolar.

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

(a) SF₆

(b) PCl₅

(c) BeH₂

(d) ${\text{CH}}_3{}^{+}$

What are the electron-pair geometry and the molecular structure of each of the following molecules or ions?

(a) ClF₅

(b) ${\text{ClO}}_2{}^{\text{−}}$

(c) ${\text{TeCl}}_4{}^{\text{2−}}$

(d) PCl₃

(e) SeF₄

(f) ${\text{PH}}_2{}^{\text{−}}$

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

(a) ClNO (N is the central atom)

(b) CS₂

(c) Cl₂CO (C is the central atom)

(d) Cl₂SO (S is the central atom)

(e) SO₂F₂ (S is the central atom)

(f) XeO₂F₂ (Xe is the central atom)

(g) ${\text{ClOF}}_2{}^{+}$ (Cl is the central atom)

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

(a) ClF₅

(b) ${\text{ClO}}_2{}^{\text{−}}$

(c) ${\text{TeCl}}_4{}^{\text{2−}}$

(d) PCl₃

(e) SeF₄

(f) ${\text{PH}}_2{}^{\text{−}}$

(g) XeF₂

Which of the following molecules have dipole moments?

(a) CS₂

(b) SeS₂

(c) CCl₂F₂

(d) PCl₃ (P is the central atom)

(e) ClNO (N is the central atom)

The molecule XF₃ has a dipole moment. Is X boron or phosphorus?

Is the Cl₂BBCl₂ molecule polar or nonpolar?

Describe the molecular structure around the indicated atom or atoms:

(a) the sulfur atom in sulfuric acid, H₂SO₄ [(HO)₂SO₂]

(b) the chlorine atom in chloric acid, HClO₃ [HOClO₂]

(c) the oxygen atom in hydrogen peroxide, HOOH

(d) the nitrogen atom in nitric acid, HNO₃ [HONO₂]

(e) the oxygen atom in the OH group in nitric acid, HNO₃ [HONO₂]

(f) the central oxygen atom in the ozone molecule, O₃

(g) each of the carbon atoms in propyne, CH₃CCH

(h) the carbon atom in Freon, CCl₂F₂

(i) each of the carbon atoms in allene, H₂CCCH₂

A molecule with the formula AB₂, 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 for each shape.

Draw the Lewis electron dot structures for these molecules, including resonance structures where appropriate:

(a) ${\text{CS}}_3{}^{\text{2−}}$

(b) CS₂

(c) CS

(d) predict the molecular shapes for ${\text{CS}}_3{}^{\text{2−}}$ and CS₂ and explain how you arrived at your predictions

A compound with a molar mass of about 42 g/mol contains 85.7% carbon and 14.3% hydrogen. What is its molecular structure?

Use the simulation to perform the following exercises for a real molecule. You may need to rotate the molecules in three dimensions to see certain dipoles.

(a) Sketch the bond dipoles and molecular dipole (if any) for O_3. Explain your observations.

(b) Look at the bond dipoles for NH₃. Use these dipoles to predict whether N or H is more electronegative.

(c) Predict whether there should be a molecular dipole for NH₃ and, if so, in which direction it will point. Check the molecular dipole box to test your hypothesis.

Use the Molecule Shape simulator to explore real molecules. On the Real Molecules tab, select H₂O. Switch between the “real” and “model” modes. Explain the difference observed.