Key Terms, Key Equations, Summaries, and Exercises (Chapter 4)

OpenStax

Private: Chapter Four

Key Terms, Key Equations, Summaries, and Exercises (Chapter 4)

Key Terms

axial position location in a trigonal bipyramidal geometry in which there is another atom at a 180° angle and the equatorial positions are at a 90° angle

binary acid compound that contains hydrogen and one other element, bonded in a way that imparts acidic properties to the compound (ability to release H+ ions when dissolved in water)

binary compound compound containing two different elements.

bond angle angle between any two covalent bonds that share a common atom

bond dipole moment separation of charge in a bond that depends on the difference in electronegativity and the bond distance represented by partial charges or a vector

bond distance (also, bond length) distance between the nuclei of two bonded atoms

bond length distance between the nuclei of two bonded atoms at which the lowest potential energy is achieved

covalent bond bond formed when electrons are shared between atoms

dipole moment property of a molecule that describes the separation of charge determined by the sum of the individual bond moments based on the molecular structure

double bond covalent bond in which two pairs of electrons are shared between two atoms

electron-pair geometry arrangement around a central atom of all regions of electron density (bonds, lone pairs, or unpaired electrons)

electronegativity tendency of an atom to attract electrons in a bond to itself

equatorial position one of the three positions in a trigonal bipyramidal geometry with 120° angles between them; the axial positions are located at a 90° angle

formal charge charge that would result on an atom by taking the number of valence electrons on the neutral atom and subtracting the nonbonding electrons and the number of bonds (one-half of the bonding electrons)

free radical molecule that contains an odd number of electrons

hypervalent molecule molecule containing at least one main group element that has more than eight electrons in its valence shell

inert pair effect tendency of heavy atoms to form ions in which their valence s electrons are not lost

ionic bond strong electrostatic force of attraction between cations and anions in an ionic compound

Lewis structure diagram showing lone pairs and bonding pairs of electrons in a molecule or an ion

Lewis symbol symbol for an element or monatomic ion that uses a dot to represent each valence electron in the element or ion

linear shape in which two outside groups are placed on opposite sides of a central atom

lone pair two (a pair of) valence electrons that are not used to form a covalent bond

molecular structure arrangement of atoms in a molecule or ion

molecular structure structure that includes only the placement of the atoms in the molecule

nomenclature system of rules for naming objects of interest

octahedral shape in which six outside groups are placed around a central atom such that a three-dimensional shape is generated with four groups forming a square and the other two forming the apex of two pyramids, one above and one below the square plane

octet rule guideline that states main group atoms will form structures in which eight valence electrons interact with each nucleus, counting bonding electrons as interacting with both atoms connected by the bond

oxyacid compound that contains hydrogen, oxygen, and one other element, bonded in a way that imparts acidic properties to the compound (ability to release H+ ions when dissolved in water)

polar covalent bond covalent bond between atoms of different electronegativities; a covalent bond with a positive end and a negative end

polar molecule (also, dipole) molecule with an overall dipole moment

pure covalent bond (also, nonpolar covalent bond) covalent bond between atoms of identical electronegativities

resonance situation in which one Lewis structure is insufficient to describe the bonding in a molecule and the average of multiple structures is observed

resonance forms two or more Lewis structures that have the same arrangement of atoms but different arrangements of electrons

resonance hybrid average of the resonance forms shown by the individual Lewis structures

single bond bond in which a single pair of electrons is shared between two atoms

tetrahedral shape in which four outside groups are placed around a central atom such that a three-dimensional shape is generated with four corners and 109.5° angles between each pair and the central atom

trigonal bipyramidal shape in which five outside groups are placed around a central atom such that three form a flat triangle with 120° angles between each pair and the central atom, and the other two form the apex of two pyramids, one above and one below the triangular plane

trigonal planar shape in which three outside groups are placed in a flat triangle around a central atom with 120° angles between each pair and the central atom

triple bond bond in which three pairs of electrons are shared between two atoms

valence shell electron-pair repulsion theory (VSEPR) theory used to predict the bond angles in a molecule based on positioning regions of high electron density as far apart as possible to minimize electrostatic repulsion

vector quantity having magnitude and direction

2
Key Equations

formal charge = # valence shell electrons (free atom) − # one pair electrons − 1 # bonding electrons

Summary

Ionic Bonding

Atoms gain or lose electrons to form ions with particularly stable electron configurations. The charges of cations formed by the representative metals may be determined readily because, with few exceptions, the electronic structures

of these ions have either a noble gas configuration or a completely filled electron shell. The charges of anions formed by the nonmetals may also be readily determined because these ions form when nonmetal atoms gain enough electrons to fill their valence shells.

Covalent Bonding

Covalent bonds form when electrons are shared between atoms and are attracted by the nuclei of both atoms. In pure covalent bonds, the electrons are shared equally. In polar covalent bonds, the electrons are shared unequally, as one atom exerts a stronger force of attraction on the electrons than the other. The ability of an atom to attract a pair of electrons in a chemical bond is called its electronegativity. The difference in electronegativity between two atoms determines how polar a bond will be. In a diatomic molecule with two identical atoms, there is no difference in electronegativity, so the bond is nonpolar or pure covalent. When the electronegativity difference is very large, as is the case between metals and nonmetals, the bonding is characterized as ionic.

Chemical Nomenclature

Chemists use nomenclature rules to clearly name compounds. Ionic and molecular compounds are named using somewhat-different methods. Binary ionic compounds typically consist of a metal and a nonmetal. The name of the metal is written first, followed by the name of the nonmetal with its ending changed to –ide. For example, K₂O is called potassium oxide. If the metal can form ions with different charges, a Roman numeral in parentheses follows the name of the metal to specify its charge. Thus, FeCl₂ is iron(II) chloride and FeCl₃ is iron(III) chloride. Some compounds contain polyatomic ions; the names of common polyatomic ions should be memorized. Molecular compounds can form compounds with different ratios of their elements, so prefixes are used to specify the numbers of atoms of each element in a molecule of the compound. Examples include SF₆, sulfur hexafluoride, and N₂O₄, dinitrogen tetroxide. Acids are an important class of compounds containing hydrogen and having special nomenclature rules. Binary acids are named using the prefix hydro-, changing the –ide suffix to –ic, and adding “acid;” HCl is hydrochloric acid. Oxyacids are named by changing the ending of the anion (–ate to –ic and –ite to

–ous), and adding “acid;” H₂CO₃ is carbonic acid.

Lewis Symbols and Structures

Valence electronic structures can be visualized by drawing Lewis symbols (for atoms and monatomic ions) and Lewis structures (for molecules and polyatomic ions). Lone pairs, unpaired electrons, and single, double, or triple bonds are used to indicate where the valence electrons are located around each atom in a Lewis structure. Most structures—especially those containing second row elements—obey the octet rule, in which every atom (except H) is surrounded by eight electrons. Exceptions to the octet rule occur for odd-electron molecules (free radicals), electron- deficient molecules, and hypervalent molecules.

Formal Charges and Resonance

In a Lewis structure, formal charges can be assigned to each atom by treating each bond as if one-half of the electrons are assigned to each atom. These hypothetical formal charges are a guide to determining the most appropriate Lewis structure. A structure in which the formal charges are as close to zero as possible is preferred. Resonance occurs in cases where two or more Lewis structures with identical arrangements of atoms but different distributions of electrons can be written. The actual distribution of electrons (the resonance hybrid) is an average of the distribution indicated by the individual Lewis structures (the resonance forms).

Molecular Structure and Polarity

VSEPR theory predicts the three-dimensional arrangement of atoms in a molecule. It states that valence electrons will assume an electron-pair geometry that minimizes repulsions between areas of high electron density (bonds and/ or lone pairs). Molecular structure, which refers only to the placement of atoms in a molecule and not the electrons, is equivalent to electron-pair geometry only when there are no lone electron pairs around the central atom. A dipole moment measures a separation of charge. For one bond, the bond dipole moment is determined by the difference in electronegativity between the two atoms. For a molecule, the overall dipole moment is determined by both the individual bond moments and how these dipoles are arranged in the molecular structure. Polar molecules (those with

an appreciable dipole moment) interact with electric fields, whereas nonpolar molecules do not.

Exercises

Ionic Bonding

1.

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

2.

Iron(III) sulfate [Fe₂(SO₄)₃] is composed of Fe³⁺ and ${SO}_{4}^{2−}$

ions. Explain why a sample of iron(III) sulfate is uncharged.

3.

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?

4.

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?

5.

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

6.

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

7.

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⁺

8.

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

9.

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

10.

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.)

4.2 Covalent Bonding

11.

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

12.

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

13.

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

14.

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

15.

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

16.

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

17.

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

18.

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

19.

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

20.

Which is the most polar bond?

(a) C–C

(b) C–H

(c) N–H

(d) O–H

(e) Se–H

21.

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

22.

Which of the following molecules or ions contain polar bonds?

(a) O₃

(b) S₈

(c) $O_{2}^{2−}$

(d) ${NO}_{3}^{−}$

(e) CO₂

(f) H₂S

(g) ${BH}_{4}^{−}$

4.3 Chemical Nomenclature

23.

Name the following compounds:

(a) CsCl

(b) BaO

(c) K₂S

(d) BeCl₂

(e) HBr

(f) AlF₃

24.

Name the following compounds:

(a) NaF

(b) Rb₂O

(c) BCl₃

(d) H₂Se

(e) P₄O₆

(f) ICl₃

25.

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

26.

Write the formulas of the following compounds:

(a) lithium carbonate

(b) sodium perchlorate

(c) barium hydroxide

(d) ammonium carbonate

(e) sulfuric acid

(f) calcium acetate

(g) magnesium phosphate

(h) sodium sulfite

27.

Write the formulas of the following compounds:

(a) chlorine dioxide

(b) dinitrogen tetraoxide

(c) potassium phosphide

(d) silver(I) sulfide

(e) aluminum fluoride trihydrate

(f) silicon dioxide

28.

Write the formulas of the following compounds:

(a) barium chloride

(b) magnesium nitride

(c) sulfur dioxide

(d) nitrogen trichloride

(e) dinitrogen trioxide

(f) tin(IV) chloride

29.

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

30.

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

(a) NiCO₃

(b) MoO₃

(c) Co(NO₃)₂

(d) V₂O₅

(e) MnO₂

(f) Fe₂O₃

31.

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)

32.

The following ionic compounds are found in common household products. Name each of the compounds:

(a) Ca(H₂PO₄)₂

(b) FeSO₄

(c) CaCO₃

(d) MgO

(e) NaNO₂

(f) KI

33.

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)

4.4 Lewis Symbols and Structures

34.

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–

35.

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

36.

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

37.

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)

Two Lewis structures are shown side-by-side, each surrounded by brackets. The left structure shows the symbol M with a superscripted two positive sign. The right shows the symbol X surrounded by four lone pairs of electrons with a superscripted two negative sign outside of the brackets.

(b)

Two Lewis structures are shown side-by-side, each surrounded by brackets. The left structure shows the symbol M with a superscripted three positive sign. The right structure shows the symbol X surrounded by four lone pairs of electrons with a superscripted negative sign and a subscripted three both outside of the brackets.

(c)

Two Lewis structures are shown side-by-side, each surrounded by brackets. The left structure shows the symbol M with a superscripted positive sign and a subscripted two outside of the brackets. The right structure shows the symbol X surrounded by four lone pairs of electrons with a superscripted two negative sign outside of the brackets.

(d)

Two Lewis structures are shown side-by-side, each surrounded by brackets. The left structure shows the symbol M with a superscripted three positive sign and a subscripted two outside of the brackets. The right structure shows the symbol X surrounded by four lone pairs of electrons with a superscripted two negative sign and subscripted three both outside of the brackets.

38.

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

39.

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^–

40.

Write Lewis structures for the following:

(a) O₂

(b) H₂CO

(c) AsF₃

(d) ClNO

(e) SiCl₄

(f) H₃O⁺

(g) ${NH}_{4}^{+}$

(h) ${BF}_{4}^{−}$

(i) HCCH

(j) ClCN

(k) $C_{2}^{2+}$

41.

Write Lewis structures for the following:

(a) ClF₃

(b) PCl₅

(c) BF₃

(d) ${PF}_{6}^{−}$

42.

Write Lewis structures for the following:

(a) SeF₆

(b) XeF₄

(c) ${SeCl}_{3}^{+}$

(d) Cl₂BBCl₂ (contains a B–B bond)

43.

Write Lewis structures for:

(a) ${PO}_{4}^{3−}$

(b) ${ICl}_{4}^{−}$

(c) ${SO}_{3}^{2−}$

(d) HONO

44.

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.”

45.

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)

46.

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.

47.

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.

48.

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.

49.

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

50.

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:

A Lewis structure is shown. A nitrogen atom is single bonded to two hydrogen atoms and a carbon atom. The carbon atom is single bonded to an oxygen atom and another nitrogen atom. That nitrogen atom is then single bonded to two hydrogen atoms.

(c) pyruvic acid:

A Lewis structure is shown. A carbon atom is single bonded to three hydrogen atoms and another carbon atom. The second carbon atom is single bonded to an oxygen atom and a third carbon atom. This carbon is then single bonded to two oxygen atoms, one of which is single bonded to a hydrogen atom.

(d) uracil:

(e) carbonic acid:

A Lewis structure is shown. A carbon atom is single bonded to three oxygen atoms. Two of those oxygen atoms are each single bonded to a hydrogen atom.

51.

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.

52.

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.

53.

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.

54.

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

4.5 Formal Charges and Resonance

55.

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

(a) selenium dioxide, OSeO

(b) nitrate ion, ${NO}_{3}^{−}$

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

(d) benzene, C₆H₆:

A Lewis structure shows a hexagonal ring composed of six carbon atoms. They form single bonds to each another and single bonds to one hydrogen atom each.

(e) the formate ion:

A Lewis structure shows a carbon atom single bonded to two oxygen atoms and a hydrogen atom. The structure is surrounded by brackets and there is a superscripted negative sign.

56.

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

(a) sulfur dioxide, SO₂

(b) carbonate ion, ${CO}_{3}^{2−}$

(c) hydrogen carbonate ion, ${HCO}_{3}^{−}$

(C is bonded to an OH group and two O atoms)

(d) pyridine:

A Lewis structure depicts a hexagonal ring composed of five carbon atoms and one nitrogen atom. Each carbon atom is single bonded to a hydrogen atom.

(e) the allyl ion:

A Lewis structure shows a carbon atom single bonded to two hydrogen atoms and a second carbon atom. The second carbon atom is single bonded to a hydrogen atom and a third carbon atom. The third carbon atom is single bonded to two hydrogen atoms. The whole structure is surrounded by brackets, and there is a superscripted negative sign.

57.

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

58.

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

59.

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:

60.

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

61.

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.

62.

Determine the formal charge of each element in the following:

(a) HCl

(b) CF₄

(c) PCl₃

(d) PF₅

63.

Determine the formal charge of each element in the following:

(a) H₃O⁺

(b) ${SO}_{4}^{2−}$

(c) NH₃

(d) $O_{2}^{2−}$

(e) H₂O₂

64.

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

65.

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

(a) F₂CO

(b) NO^–

(c) ${BF}_{4}^{−}$

(d) ${SnCl}_{3}^{−}$

(e) H₂CCH₂

(f) ClF₃

(g) SeF₆

(h) ${PO}_{4}^{3−}$

66.

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) ${NO}_{2}^{−}$

(d) ${NO}_{3}^{−}$

67.

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

68.

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

69.

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

70.

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

71.

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

72.

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.

73.

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

74.

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.

4.6 Molecular Structure and Polarity

75.

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

76.

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?

77.

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

78.

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 ${NH}_{4}^{+}$

identical to the H–C–H bond angle in CH₄?

79.

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

80.

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?

81.

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

(a) SF₆

(b) PCl₅

(c) BeH₂

(d) ${CH}_{3}^{+}$

82.

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

(a) ${IF}_{6}^{+}$

(b) CF₄

(c) BF₃

(d) ${SiF}_{5}^{−}$

(e) BeCl₂

83.

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

(a) ClF₅

(b) ${ClO}_{2}^{−}$

(c) ${TeCl}_{4}^{2−}$

(d) PCl₃

(e) SeF₄

(f) ${PH}_{2}^{−}$

84.

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

(a) H₃O⁺

(b) ${PCl}_{4}^{−}$

(c) ${SnCl}_{3}^{−}$

(d) ${BrCl}_{4}^{−}$

(e) ICl₃

(f) XeF₄

(g) SF₂

85.

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) ${ClOF}_{2}^{+}$

(Cl is the central atom)

86.

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) ${NO}_{2}^{−}$

(g) ${SiO}_{4}^{4−}$

87.

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

(a) ClF₅

(b) ${ClO}_{2}^{−}$

(c) ${TeCl}_{4}^{2−}$

(d) PCl₃

(e) SeF₄

(f) ${PH}_{2}^{−}$

(g) XeF₂

88.

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

(a) H₃O⁺

(b) ${PCl}_{4}^{−}$

(c) ${SnCl}_{3}^{−}$

(d) ${BrCl}_{4}^{−}$

(e) ICl₃

(f) XeF₄

(g) SF₂

89.

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)

90.

Identify the molecules with a dipole moment:

(a) SF₄

(b) CF₄

(c) Cl₂CCBr₂

(d) CH₃Cl

(e) H₂CO

91.

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

92.

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

93.

Is the Cl₂BBCl₂ molecule polar or nonpolar?

94.

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.

95.

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

96.

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

(a) CO₂

(b) ${NO}_{2}^{−}$

(c) SO₃

(d) ${SO}_{3}^{2−}$

97.

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.

98.

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.

99.

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

(a) ${CS}_{3}^{2−}$

(b) CS₂

(c) CS

(d) predict the molecular shapes for ${CS}_{3}^{2−}$

and CS₂ and explain how you arrived at your predictions

100.

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

101.

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

102.

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.

103.

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.

104.

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.

105.

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.

106.

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.

License

Icon for the Creative Commons Attribution 4.0 International License

Key Terms

2 Key Equations

Summary

Exercises

4.2 Covalent Bonding

4.3 Chemical Nomenclature

4.4 Lewis Symbols and Structures

4.5 Formal Charges and Resonance

4.6 Molecular Structure and Polarity

License

Share This Book

2
Key Equations