7 Ionic Compounds: Formulas, Names, and Properties

69



2.7 Ionic Compounds: Formulas, Names, and Properties







rEvIEW & cHEcK FOr SEctIOn 2.6

Cysteine, whose molecular model and structural formula are illustrated here, is an important

amino acid and a constituent of many living things. What is its molecular formula?



+



NH3 H



−



O

C

O



Molecular model



(a)



C3H6O2S



C



C



H



H



S



H



Structural

al fform

formula

ula



(b) C3H7NO2S



(c)



C3H7N2OS



(d) C3H7NO2



2.7 IonicCompounds:Formulas,Names,

andProperties

The compounds you have encountered so far in this chapter are molecular

compounds—that is, compounds that consist of discrete molecules at the particulate level. Ionic compounds constitute another major class of compounds. They

consist of ions, atoms, or groups of atoms that bear a positive or negative electric

charge. Many familiar compounds are composed of ions (Figure 2.16). Table salt, or

sodium chloride (NaCl), and lime (CaO) are just two. To be able to recognize ionic

compounds and to write formulas for these compounds, it is important to know the

formulas and charges of common ions. You also need to know the names of ions

and be able to name the compounds they form.



Modules 2: Predicting Ion Charges

and 3: Names to Formulas of Ionic

Compounds cover concepts in this

section.



Ions

Atoms of many elements can gain or lose electrons in the course of a chemical reaction. To be able to predict the outcome of chemical reactions [▶ Chapter 3], you need

to know whether an element will likely gain or lose electrons and, if so, how many.



Cations

If an atom loses an electron (which is transferred to an atom of another element in

the course of a reaction), the atom now has one fewer negative electrons than it has

positive protons in the nucleus. The result is a positively charged ion called a cation



FIGURE2.16 Some common ionic compounds.



Calcite

Fluorite

Gypsum



Hematite

Orpiment



kotz_48288_02_0050-0109.indd 69



Hematite, Fe2O3

Name



Formula



Ions Involved



Calcium

carbonate

Calcium

fluoride

Calcium

sulfate

dihydrate

Iron(III)

oxide

Arsenic

sulfide



CaCO3



Ca2+, CO32−



CaF2



Ca2+, F−



CaSO4 ∙ 2 H2O



Ca2+, SO42−



Fe2O3



Fe3+, O2−



As2S3



As3+, S2−



© Cengage Learning/Charles D. Winters



Common

Name



Gypsum, CaSO4 ⋅ 2 H2O



Calcite, CaCO3



Fluorite, CaF2



Orpiment, As2S3



11/18/10 2:06 PM



70



c h a p t er 2   Atoms, Molecules, and Ions

e–+



Li atom

(3 protons and 3 electrons)



FIGURE 2.17   Ions.  A lithium-6

atom is electrically neutral because the

number of positive charges (three

protons) and negative charges (three

electrons) are the same. When it loses

one electron, it has one more positive

charge than negative charge, so it has a

net charge of 1+. We symbolize the

resulting lithium cation as Li+. A fluorine atom is also electrically neutral,

having nine protons and nine electrons.

A fluorine atom can acquire an electron

to produce an F− anion. This anion has

one more electron than it has protons,

so it has a net charge of 1−.



3e –



Li + cation

(3 protons and 2 electrons)



e–



2e –

3p

3n



3p

3n



Li



Li +



3p



3p



3n



3n



3e –



2e –



Lithium ion, Li +

Lithium, Li



9e –

9p

10n



10e –



e–



9p

10n



F



F–



9p

10n

9e –



9p

10n

10e –



Fluorine, F

Fluoride ion, F –



(Figure 2.17). (The name is pronounced “cat´-i-on.”) Because it has an excess of

one positive charge, we write the cation’s symbol as, for example, Li+:

Li atom  →  e−  +  Li+ cation

(3 protons and 3 electrons)   (3 protons and 2 electrons)



Anions

Conversely, if an atom gains one or more electrons, there is now one or more negatively charged electrons than protons. The result is an anion (pronounced “ann´-i-on”)

(Figure 2.17).

O atom + 2 e−  →  O2− anion

(8 protons and 8 electrons)   (8 protons and 10 electrons)



Here the O atom has gained two electrons, so we write the anion’s symbol as O2−.

How do you know whether an atom is likely to form a cation or an anion? It

depends on whether the element is a metal or a nonmetal.

•

•



Metals generally lose electrons in their reactions to form cations.

Nonmetals frequently gain one or more electrons to form anions in their

reactions.



Monatomic Ions

Monatomic ions are single atoms that have lost or gained electrons. As indicated in

Figure 2.18, metals typically lose electrons to form monatomic cations, and nonmetals typically gain electrons to form monatomic anions.

How can you predict the number of electrons gained or lost? Like lithium in

Figure 2.18, metals of Groups 1A–3A form positive ions having a charge equal to the group

number of the metal.



kotz_48288_02_0050-0109.indd 70



Group



Metal Atom



Electron Change



1A

2A

3A



Na(11 protons, 11 electrons)

Ca(20 protons, 20 electrons)

Al (13 protons, 13 electrons)



−1

−2

−3



Resulting Metal Cation

→

→

→



Na+ (11 protons, 10 electrons)

Ca2+(20 protons, 18 electrons)

Al3+ (13 protons, 10 electrons)



11/18/10 2:06 PM



2.7 Ionic Compounds: Formulas, Names, and Properties







1A



7A



H+



Metals

Transition metals

Metalloids

Nonmetals



2A



Li+

Na+ Mg2+

K+ Ca2+



3B



4B



5B



Ti4+



3A



4A



5A



6A



N3− O2−



8B

6B 7B

1B 2B

Cr2+ Mn2+ Fe2+ Co2+ 2+ Cu+

Ni

Cr3+

Fe3+ Co3+

Cu2+ Zn2+



Rb+ Sr2+



Ag+ Cd2+



Cs+ Ba2+



Hg22+

Hg2+



Al3+



P3−



8A



H−

F−



S2− Cl−



7B

8B

8B



Metal Atom



Electron Change

−2

−2

−3



Mn (25 protons, 25 electrons)

Fe (26 protons, 26 electrons)

Fe (26 protons, 26 electrons)



FIGURE2.18 Charges on some

common monatomic cations and

anions. Metals usually form cations

and nonmetals usually form anions.

(The boxed areas show ions of identical charge.)



Se2− Br−

Sn2+



Te2− I−



Pb2+ Bi3+



Transition metals (B-group elements) also form cations. Unlike the A-group

metals, however, no easily predictable pattern of behavior occurs for transition

metal cations. In addition, transition metals often form several different ions. An

iron-containing compound, for example, may contain either Fe2+ or Fe3+ ions. Indeed, 2+ and 3+ ions are typical of many transition metals (Figure 2.18).

Group



71



• Writing Ion Formulas When writing

the formula of an ion, the charge on

the ion must be included.



Resulting Metal Cation

→

→

→



Mn2+ (25 protons, 23 electrons)

Fe2+ (26 protons, 24 electrons)

Fe3+ (26 protons, 23 electrons)



Nonmetals often form ions having a negative charge equal to the group number of the element minus 8. For example, nitrogen is in Group 5A, so it forms an ion having a

3− charge because a nitrogen atom can gain three electrons.

Group



Nonmetal Atom



Electron Change



Resulting Nonmetal Anion



5A



N (7 protons, 7 electrons)



+3



→



6A



S (16 protons, 16 electrons)



+2



→



7A



Br (35 protons, 35 electrons)



+1



→



N3− (7 protons, 10 electrons)

Charge = 5 − 8 = −3

S2− (16 protons, 18 electrons)

Charge = 6 − 8 = −2

Br− (35 protons, 36 electrons)

Charge = 7 − 8 = −1



Notice that hydrogen appears at two locations in Figure 2.18. The H atom can

either lose or gain electrons, depending on the other atoms it encounters.

Electron lost:



H (1 proton, 1 electron) →



Electron gained:



H (1 proton, 1 electron) + e− →



H+ (1 proton, 0 electrons) + e−

H− (1 proton, 2 electrons)



Finally, the noble gases very rarely form monatomic cations and never form

monatomic anions in chemical reactions.



Ion Charges and the Periodic Table

The metals of Groups 1A, 2A, and 3A form ions having 1+, 2+, and 3+ charges

(Figure 2.18); that is, their atoms lose one, two, or three electrons, respectively. For

Group 1A and 2A metals and aluminum, the number of electrons remaining on the cation is

the same as the number of electrons in an atom of the noble gas that precedes it in the periodic

table. For example, Mg2+ has 10 electrons, the same number as in an atom of the

noble gas neon (atomic number 10).

An atom of a nonmetal near the right side of the periodic table would have to

lose a great many electrons to achieve the same number as a noble gas atom of lower

atomic number. (For instance, Cl, whose atomic number is 17, would have to lose

7 electrons to have the same number of electrons as Ne.) If a nonmetal atom were

to gain just a few electrons, however, it would have the same number as a noble gas



kotz_48288_02_0050-0109.indd 71



1A



2A



3A



Group 1A, 2A, 3A metals form

Mn+ cations where n = group number.



Main group metals.



11/18/10 2:06 PM



c h a p t er 2   Atoms, Molecules, and Ions



Photos © Cengage Learning/

Charles D. Winters



72



Calcite, CaCO3

Calcium carbonate



CO32–



PO43–



Apatite, Ca5F(PO4)3

Calcium fluorophosphate



SO42–

Celestite, SrSO4

Strontium sulfate



FIGURE 2.19   Common ionic compounds based on polyatomic ions.

atom of higher atomic number. For example, an oxygen atom has eight electrons.

By gaining two electrons per atom it forms O2−, which has 10 electrons, the same

number as neon. Anions having the same number of electrons as the noble gas atom succeeding it in the periodic table are commonly observed in chemical compounds.



Polyatomic Ions

Polyatomic ions are made up of two or more atoms, and the collection has an electric charge (Figure 2.19 and Table 2.4). For example, carbonate ion, CO32−, a common polyatomic anion, consists of one C atom and three O atoms. The ion has two

Table 2.4  Formulas and Names of Some Common Polyatomic Ions

Formula



Name



Formula



Name



Cation: Positive Ion

NH4+



Ammonium ion



Anions: Negative Ions

Based on a Group 4A element



Based on a Group 7A element



CN−



Cyanide ion



ClO−



Hypochlorite ion



CH3CO2−



Acetate ion



ClO2−



Chlorite ion



CO32−



Carbonate ion



ClO3−



Chlorate ion



HCO3−



Hydrogen carbonate ion

(or bicarbonate ion)



ClO4−



Perchlorate ion



C2O42−



Oxalate ion



Based on a Group 5A element



Based on a transition metal



NO2−



Nitrite ion



CrO42−



Chromate ion



NO3−



Nitrate ion



Cr2O72−



Dichromate ion



PO43−



Phosphate ion



MnO4−



Permanganate ion



HPO42−



Hydrogen phosphate ion



H2PO4−



Dihydrogen phosphate ion



Based on a Group 6A element



kotz_48288_02_0050-0109.indd 72



OH−



Hydroxide ion



SO32−



Sulfite ion



SO42−



Sulfate ion



HSO4−



Hydrogen sulfate ion

(or bisulfate ion)



11/18/10 2:06 PM



2.7  Ionic Compounds: Formulas, Names, and Properties







73



units of negative charge because there are two more electrons (a total of 32) in the

ion than there are protons (a total of 30) in the nuclei of one C atom and three

O atoms.

The ammonium ion, NH4+, is a common polyatomic cation. In this case, four

H atoms surround an N atom, and the ion has a 1+ electric charge. This ion has

10 electrons, but there are 11 positively charged protons in the nuclei of the N and

H atoms (seven and one each, respectively).



Compounds are electrically neutral; that is, they have no net electric charge. Thus,

in an ionic compound the numbers of positive and negative ions must be such that

the positive and negative charges balance. In sodium chloride, the sodium ion has

a 1+ charge (Na+) and the chloride ion has a 1− charge (Cl−). These ions must be

present in a 1 : 1 ratio, and so the formula is NaCl.

The gemstone ruby is largely the compound formed from aluminum ions (Al3+)

and oxide ions (O2−) (but the color comes from a trace of Cr3+ ions.) Here the ions

have positive and negative charges that are of different absolute value. To have a

compound with the same number of positive and negative charges, two Al3+ ions

[total charge  =  2  ×  (3+)  =  6+] must combine with three O2− ions [total

charge = 3 × (2−) = 6−] to give a formula of Al2O3.

Calcium is a Group 2A metal, and it forms a cation having a 2+ charge. It can

combine with a variety of anions to form ionic compounds such as those in the following table:

Compound



Ion Combination



Overall Charge on Compound



CaCl2

CaCO3

Ca3(PO4)2



Ca2+ + 2 Cl−

Ca2+ + CO32−

3 Ca2+ + 2 PO43−



(2+) + 2 × (1−) = 0

(2+) + (2−) = 0

3 × (2+) + 2 × (3−) = 0



In writing formulas of ionic compounds, the convention is that the symbol of the cation

is given first, followed by the anion symbol. Also notice the use of parentheses when more

than one polyatomic ion of a given kind is present [as in Ca3(PO4)2]. (None, however, are used when only one polyatomic ion is present, as in CaCO3.)



EXAMPLE 2.4



Ionic Compound Formulas



Problem  For each of the following ionic compounds, write the symbols for the ions present and

give the relative number of each: (a) Li2CO3, and (b) Fe2(SO4)3.



© Cengage Learning/Charles D. Winters



Formulas of Ionic Compounds



Ruby, Al2O3.  Gems called rubies are

largely composed of Al3+ and O2- ions

with a trace of Cr3+ ion. It is the

chromium(III) ions that give the gem

its color.



• Balancing Ion Charges in Formulas 

Aluminum, a metal in Group 3A, loses

three electrons to form the Al3+ cation.

Oxygen, a nonmetal in Group 6A, gains

two electrons to form an O2– anion.

Notice that in the compound formed

from these ions the charge on the cation is the subscript on the anion, and

vice versa.

2 Al3+ + 3 O2– n Al­2O3

This often works well, but be careful.

The subscripts of Ti4+ and O2– are

reduced to the simplest ratio in TiO2

(1 Ti to 2 O, rather than 2 Ti to 4 O).

Ti4+ + 2 O2– n TiO2



What Do You Know?  You know the formulas of the ionic compounds, the predicted charges on

monatomic ions (Figure 2.18), and the formulas and charges of polyatomic ions (Table 2.4).

Strategy  Divide the formula of the compound into the cations and anions. To accomplish this

you will have to recognize, and remember, the composition and charges of common ions.

Solution

(a)  Li2C O3 is composed of two lithium ions, Li+, for each carbonate ion, CO32−.  Li is a Group 1A

element and always has a 1+ charge in its compounds. Because the two 1+ charges balance

the negative charge of the carbonate ion, the latter must be 2−.

(b)  Fe2(SO4)3 contains two iron(III) ions, Fe3+, for every three sulfate ions, SO42−.  The way to

recognize this is to recall that sulfate has a 2− charge. Because three sulfate ions are present

(with a total charge of 6−), the two iron cations must have a total charge of 6+. This is

possible only if each iron cation has a charge of 3+.

Think about Your Answer  The charges predicted are in line with those in Figure 2.18 and Table

2.4. Remember that the formula for an ion must include its composition and its charge. Formulas

for ionic compounds are always written with the cation first and then the anion, but ion charges

are not included.



kotz_48288_02_0050-0109.indd 73



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74



c h a p t er 2   Atoms, Molecules, and Ions



Check Your Understanding

(a) Give the number and identity of the constituent ions in each of the following ionic compounds: NaF, Cu(NO3)2, and NaCH3CO2.

(b) Iron, a transition metal, forms ions having 2+ and 3+ charges. Write the formulas of the

compounds formed between chloride ions and these two different iron cations.



EXAMPLE 2.5



Ionic Compound Formulas



Problem  Write formulas for ionic compounds composed of an aluminum cation and each of the

following anions: (a) fluoride ion, (b) sulfide ion, and (c) nitrate ion.

What Do You Know?  You know the names of the ions involved, the predicted charges on

monatomic ions (Figure 2.18), and the names of polyatomic ions (Table 2.4).

Strategy  First decide on the formula of the Al cation and the formula of each anion. Combine

the Al cation with each type of anion to form electrically neutral compounds.

Solution  An aluminum cation is predicted to have a charge of 3+ because Al is a metal in

Group 3A.

(a) Fluorine is a Group 7A element. The charge of the fluoride ion is predicted to be 1− (from

7 − 8 = 1−). Therefore, we need 3 F− ions to combine with one Al3+. The formula of the

compound is  AlF3. 

(b) Sulfur is a nonmetal in Group 6A, so it forms a 2− anion. Thus, we need to combine two

Al3+ ions [total charge is 6+ = 2 × (3+)] with three S2− ions [total charge is 6− = 3 × (2−)].

The compound has the formula  Al2S3. 

(c) The nitrate ion has the formula NO3− (see Table 2.4). The answer here is therefore similar to

the AlF3 case, and the compound has the formula  Al(NO3)3.  Here we place parentheses

around NO3 to show that three polyatomic NO3− ions are involved.

Think about Your Answer  The most common error students make is not knowing the correct

charge on an ion.

Check Your Understanding

Write the formulas of all neutral ionic compounds that can be formed by combining the cations

Na+ and Ba2+ with the anions S2− and PO43−.



Names of Ions

Naming Positive Ions (Cations)



• “-ous” and “-ic” Endings  An older



naming system for metal ions uses the

ending -ous for the ion of lower charge

and -ic for the ion of higher charge. For

example, there are cobaltous (Co2+) and

cobaltic (Co3+) ions. In addition, this

older system sometimes uses the root of

the Latin name of some elements in the

names of their ions. For example, the

Latin name for iron is ferrum, and this

system calls the iron cations the ferrous

(Fe2+) and ferric (Fe3+) ions. We do not

use this system in this book, but some

chemical manufacturers continue to

use it.



kotz_48288_02_0050-0109.indd 74



With a few exceptions (such as NH4+), the positive ions described in this text are

metal ions. Positive ions are named by the following rules:

1. For a monatomic positive ion (that is, a metal cation) the name is that of the

metal plus the word “cation.” For example, we have already referred to Al3+ as

the aluminum cation.

2. Some cases occur, especially in the transition series, in which a metal can

form more than one type of positive ion. In these cases the charge of the ion

is indicated by a Roman numeral in parentheses immediately following the

ion’s name. For example, Co2+ is the cobalt(II) cation, and Co3+ is the

cobalt(III) cation.

Finally, you will encounter the ammonium cation, NH4+, many times in this

book and in the laboratory. Do not confuse the ammonium cation with the ammonia molecule, NH3, which has no electric charge and one less H atom.



11/18/10 2:06 PM



2.7 Ionic Compounds: Formulas, Names, and Properties







Naming Negative Ions (Anions)



1−



There are two types of negative ions: those having only one atom (monatomic) and

those having several atoms (polyatomic).



2.



A monatomic negative ion is named by adding -ide to the stem of the name of

the nonmetal element from which the ion is derived (Figure 2.20). The anions

of the Group 7A elements, the halogens, are known as the fluoride, chloride,

bromide, and iodide ions and as a group are called halide ions.

Polyatomic negative ions are common, especially those containing oxygen

(called oxoanions). The names of some of the most common oxoanions are

given in Table 2.4. Although most of these names must simply be learned, some

guidelines can help. For example, consider the following pairs of ions:

NO3− is the nitrate ion, whereas NO2− is the nitrite ion.

SO42− is the sulfate ion, whereas SO32− is the sulfite ion.



The oxoanion having the greater number of oxygen atoms is given the suffix

-ate, and the oxoanion having the smaller number of oxygen atoms has the suffix

-ite. For a series of oxoanions having more than two members, the ion with the largest number of oxygen atoms has the prefix per- and the suffix -ate. The ion having

the smallest number of oxygen atoms has the prefix hypo- and the suffix -ite. The

chlorine oxoanions are the most commonly encountered example.

ClO4−



Perchlorate ion



ClO3−



Chlorate ion



ClO2−



Chlorite ion



ClO−



Hypochlorite ion



Oxoanions that contain hydrogen are named by adding the word “hydrogen”

before the name of the oxoanion. If two hydrogens are in the anion, we say “dihydrogen.” Many hydrogen-containing oxoanions have common names that are

used as well. For example, the hydrogen carbonate ion, HCO3–, is called the bicarbonate ion.

Ion



Systematic Name



HPO42−

H2PO4−

HCO3−

HSO4−

HSO3−



Hydrogen phosphate ion

Dihydrogen phosphate ion

Hydrogen carbonate ion

Hydrogen sulfate ion

Hydrogen sulfite ion



Common Name



3−



2−



H−

hydride

ion



F−



N3− O2−

nitride

ion



oxide

ion



fluoride

ion



P3−



S2−



Cl−



phosphide sulfide

ion

ion



chloride

ion



Se2− Br−

selenide bromide

ion

ion



Te2−



I−



telluride

ion



iodide

ion



FIGURE2.20 Names

and charges of some common

monatomic ions.



• Naming Oxoanions

per . . . ate

increasing

oxygen content



1.



75



. . . ate

. . . ite

hypo . . . ite



Bicarbonate ion

Bisulfate ion

Bisulfite ion



Names of Ionic Compounds

The name of an ionic compound is built from the names of the positive and negative ions in the compound. The name of the positive cation is given first, followed

by the name of the negative anion. Examples of ionic compound names are given

below.

Ionic Compound



Ions Involved



Name



CaBr2

NaHSO4

(NH4)2CO3

Mg(OH)2

TiCl2

Co2O3



Ca2+ and 2 Br−

Na+ and HSO4−

2 NH4+ and CO32−

Mg2+ and 2 OH−

Ti2+ and 2 Cl−

2 Co3+ and 3 O2−



Calcium bromide

Sodium hydrogen sulfate

Ammonium carbonate

Magnesium hydroxide

Titanium(II) chloride

Cobalt(III) oxide



kotz_48288_02_0050-0109.indd 75



• Names of Compounds Containing

Transition Metal Cations Be sure to

notice that the charge on a transition

metal cation is indicated by a Roman

numeral and is included in the name.



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76



c h a p t er 2 Atoms, Molecules, and Ions



PrOBLEM SOLvInG tIP 2.1

Writing formulas for ionic compounds takes

practice, and it requires that you know the

formulas and charges of the most common

ions. The charges on monatomic ions are

often evident from the position of the element in the periodic table, but you simply

have to remember the formula and charges

of polyatomic ions, especially the most common ones such as nitrate, sulfate, carbonate,

phosphate, and acetate.



Formulas for Ions and Ionic Compounds

If you cannot remember the formula of

a polyatomic ion or if you encounter an ion

you have not seen before, you may be able

to figure out its formula. For example, suppose you are told that NaCHO2 is sodium

formate. You know that the sodium ion is

Na+, so the formate ion must be the remaining portion of the compound; it must have

a charge of 1− to balance the 1+ charge on



the sodium ion. Thus, the formate ion must

be CHO2−.

Finally, when writing the formulas of

ions, you must include the charge on the ion

(except in the formula of an ionic compound). Writing Na when you mean sodium

ion is incorrect. There is a vast difference in

the properties of the element sodium (Na)

and those of its ion (Na+).



Properties of Ionic Compounds

When a particle having a negative electric charge is brought near another particle having a positive electric charge, there is a force of attraction between them (Figure 2.21).

In contrast, there is a repulsive force when two particles with the same charge—both

positive or both negative—are brought together. These forces are called electrostatic

forces, and the force of attraction (or repulsion) between ions is given by Coulomb’s

law (Equation 2.3):

charge on + and − ions



Force = −k



charge on electron



(n+e)(n−e)

d2



proportionality constant



(2.3)



distance between ions



where, for example, n+ is +3 for Al3+ and n− is −2 for O2−. Based on Coulomb’s law,

the force of attraction between oppositely charged ions increases

as the ion charges (n+ and n−) increase. Thus, the attraction between ions

having charges of 2+ and 2− is greater than that between ions having 1+ and

1− charges (Figure 2.21).

as the distance between the ions becomes smaller (Figure 2.21).



•

•



Ionic compounds do not consist of simple pairs or small groups of positive and

negative ions. The simplest ratio of cations to anions in an ionic compound is represented by its formula, but an ionic solid consists of millions upon millions of ions

arranged in an extended three-dimensional network called a crystal lattice. A



+1

n+ = 1



+



Li+



−

+2



F−



d small



d



n– = –1



+



−1



−2



d large



LiF



(a)



(a) Ions such as Li+ and F– are held together by an electrostatic force. Here a lithium ion is attracted to a fluoride ion,

and the distance between the nuclei of the two ions is d.



As ion charge increases,

force of attraction increases



As distance increases,

force of attraction decreases



(b)



(b) Forces of attraction between ions of opposite charge

increase with increasing ion charge and decrease with

increasing distance (d).



FIGURE2.21 Coulomb’s law and electrostatic forces.



kotz_48288_02_0050-0109.indd 76



11/18/10 2:06 PM



77



2.7 Ionic Compounds: Formulas, Names, and Properties







PrOBLEM SOLvInG tIP 2.2



Is a Compound Ionic?

behavior: All elements to the left of a

diagonal line running from boron to tellurium in the periodic table are metallic.

2. If there is no metal in the formula, it is

likely that the compound is not ionic. The

exceptions here are compounds composed of polyatomic cations based on

nonmetals (e.g., NH4Cl or NH4NO3).

3. Learn to recognize the formulas of polyatomic ions (see Table 2.4). Chemists

write the formula of ammonium nitrate as



1. Most metal-containing compounds are

ionic. So, if a metal atom appears in the

formula of a compound and especially

when it is the element listed first, a good

first guess is that it is ionic. (There are

interesting exceptions, but few come up

in introductory chemistry.) It is helpful in

this regard to recall trends in metallic



NH4NO3 (not as N2H4O3) to alert others

to the fact that it is an ionic compound

composed of the common polyatomic

ions NH4+ and NO3–.

As an example of these guidelines, you

can be sure that MgBr2 (Mg2+ with Br–) and

K2S (K+ with S2–) are ionic compounds. On

the other hand, the compound CCl4, formed

from two nonmetals, C and Cl, is not ionic.



portion of the lattice for NaCl, illustrated in Figure 2.22, represents a common way

of arranging ions for compounds that have a 1∶1 ratio of cations to anions.

Ionic compounds have characteristic properties that can be understood in

terms of the charges of the ions and their arrangement in the lattice. Because

each ion is surrounded by oppositely charged nearest neighbors, it is held tightly

in its allotted location. At room temperature each ion can move just a bit around

its average position, but considerable energy must be added before an ion can

escape the attraction of its neighboring ions. Only if enough energy is added will

the lattice structure collapse and the substance melt. Greater attractive forces

mean that ever more energy—and higher and higher temperatures—is required

to cause melting. Thus, Al2O3, a solid composed of Al3+ and O2− ions, melts at a

much higher temperature (2072 °C) than NaCl (801 °C), a solid composed of

Na+ and Cl− ions.

Most ionic compounds are “hard” solids. That is, the solids are not pliable or

soft. The reason for this characteristic is again related to the lattice of ions. The

nearest neighbors of a cation in a lattice are anions, and the force of attraction

makes the lattice rigid. However, a blow with a hammer can cause the lattice to

break cleanly along a sharp boundary. The hammer blow displaces layers of ions just

enough to cause ions of like charge to become nearest neighbors, and the repulsion

between these like-charged ions forces the lattice apart (Figure 2.23).



© Cengage Learning/Charles D. Winters



Students often ask how to know whether a

compound is ionic. Here are some useful

guidelines.



FIGURE2.22 Sodium chloride.

A crystal of NaCl consists of an

extended lattice of sodium ions and

chloride ions in a 1∶1 ratio.



© Cengage Learning/Charles D. Winters



FIGURE2.23 Ionic solids.



(a)



(a) An ionic solid is normally rigid owing

to the forces of attraction between

oppositely charged ions. When struck

sharply, however, the crystal can cleave

cleanly.



kotz_48288_02_0050-0109.indd 77



(b)



(b) When a crystal is struck, layers of ions move slightly, and ions of like charge become nearest

neighbors. Repulsions between ions of similar charge cause the crystal to cleave.



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78



c h a p t er 2 Atoms, Molecules, and Ions



rEvIEW & cHEcK FOr SEctIOn 2.7

1.



What is the most likely charge on an ion of barium?

(a)



2.



2−



(b) 2+



loses 3 electrons



(b) gains 3 electrons



nitrohydrogen sulfide



(b) ammonium sulfide



Ba(CH3CO2)2



(b) BaCH3CO2



(d) gains 2 electrons



(c)



ammonium sulfur



(d) ammonia sulfide



(c)



BaMnO4



(d) BaCO3



The name of the compound with the formula V2O3 is

(a)



vanadium(III) oxide



(b) vanadium oxide

6.



loses 2 electrons



The formula of barium acetate is

(a)



5.



(c)



(d) 1−



The name of the compound (NH4)2S is

(a)



4.



3+



When gallium forms an ion, it

(a)



3.



(c)



(c)



divanadium trioxide



(d) vanadium trioxide



Which should have the higher melting point, MgO or NaCl?

(a)



MgO



(b) NaCl



2.8 MolecularCompounds:FormulasandNames

Many familiar compounds are not ionic, they are molecular: the water you drink,

the sugar in your coffee or tea, or the aspirin you take for a headache.

Ionic compounds are generally solids, whereas molecular compounds can

range from gases to liquids to solids at ordinary temperatures (Figure 2.24). As size

and molecular complexity increase, compounds generally exist as solids. We will

explore some of the underlying causes of these general observations in Chapter 12.

Some molecular compounds have complicated formulas that you cannot, at this

stage, predict or even decide if they are correct. However, there are many simple

compounds you will encounter often, and you should understand how to name

them and, in many cases, know their formulas.

Let us look first at molecules formed from combinations of two nonmetals.

These “two-element” or binary compounds of nonmetals can be named in a systematic way.

FIGURE2.24 Molecular compounds. Ionic compounds are gener-



© Cengage Learning/Charles D. Winters



ally solids at room temperature. In

contrast, molecular compounds can be

gases, liquids, or solids. The molecular

models are of caffeine (in coffee), water,

and citric acid (in lemons).



kotz_48288_02_0050-0109.indd 78



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79



2.8 Molecular Compounds: Formulas and Names







Hydrogen forms binary compounds with all of the nonmetals except the noble

gases. For compounds of oxygen, sulfur, and the halogens, the H atom is generally

written first in the formula and is named first. The other nonmetal is named as if it

were a negative ion.

Compound



Name



HF

HCl

H 2S



Hydrogen fluoride

Hydrogen chloride

Hydrogen sulfide



• Formulas of Binary Nonmetal

Compounds Containing Hydrogen

Simple hydrocarbons (compounds of

C and H) such as methane (CH4) and

ethane (C2H6) have formulas written

with H following C, and the formulas

of ammonia and hydrazine have H following N. Water and the hydrogen

halides, however, have the H atom

preceding O or the halogen atom.

Tradition is the only explanation for

such irregularities in writing formulas.



Although there are exceptions, most binary molecular compounds are a combination

of nonmetallic elements from Groups 4A–7A with one another or with hydrogen. The formula

is generally written by putting the elements in order of increasing group number.

When naming the compound, the number of atoms of a given type in the compound is designated with a prefix, such as “di-,” “tri-,” “tetra-,” “penta-,” and so on.

Compound



Systematic Name



NF3

NO

NO2

N 2O

N 2O 4

PCl3

PCl5

SF6

S2F10



Nitrogen trifluoride

Nitrogen monoxide

Nitrogen dioxide

Dinitrogen monoxide

Dinitrogen tetraoxide

Phosphorus trichloride

Phosphorus pentachloride

Sulfur hexafluoride

Disulfur decafluoride



• Hydrocarbons Compounds such as

methane, ethane, propane, and butane

belong to a class of hydrocarbons

called alkanes.



Finally, many of the binary compounds of nonmetals were discovered years ago

and have common names.

Compound



Common Name



Compound



Common Name



CH4

C 2H 6

C 3H 8

C4H10

NH3



Methane

Ethane

Propane

Butane

Ammonia



N 2H 4

PH3

NO

N 2O

H 2O



Hydrazine

Phosphine

Nitric oxide

Nitrous oxide (“laughing gas”)

Water



methane, CH4



propane, C3H8



ethane, C2H6



butane, C4H10



rEvIEW & cHEcK FOr SEctIOn 2.8

1.



What is the formula for dioxygen difluoride?

(a)



2.



O2F2



(b) OF



carbon disulfide, CS2



(b) nitrogen pentaoxide, N2O5



(d) OF2



(c)



boron trifluoride, BF3



(d) sulfur tetrafluoride, SF4



The name of the compound with the formula N2F4 is

(a)



nitrogen fluoride



(b) dinitrogen fluoride

4.



O2F



Compound names and formulas are listed below. Which name is incorrect?

(a)



3.



(c)



(c)



dinitrogen tetrafluoride



(d) nitrogen tetrafluoride



The name of the compound with the formula P4O10 is

(a)



phosphorus oxide



(b) tetraphosphorus decaoxide



kotz_48288_02_0050-0109.indd 79



(c)



tetraphosphorus oxide



(d) phosphorus decaoxide



11/18/10 2:06 PM