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C hapter 4
C hemiCal Compounds
4.1 Classification of Matter
4.2 Compounds and Chemical Bonds
4.3 Molecular Compounds
4.4 Naming Binary Covalent Compounds
4.5 Ionic Compounds
Review Skills
The presentation of information in this chapter assumes that you can already perform the tasks listed below. You can test your readiness to proceed by answering the Review Questions at the end of the chapter. This might also be a good time to read the Chapter Objectives, which precede the Review Questions.
Describe the particle nature of solids, liquids, and gases. (Section 3.1) Convert between the names and symbols for the common elements. (Table 3.1) Given a periodic table, write the number of the group to which each element belongs. (Figure 3.3) Given a periodic table, identify the alkali metals, alkaline earth metals, halogens, and noble gases. (Section 3.3)
Using a periodic table, classify elements as metals, nonmetals, or metalloids. (Section 3.3) Describe the nuclear model of the atom. (Section 3.4) Define the terms ion, cation, and anion. (Section 3.4) Define the terms covalent bond, molecule, and diatomic. (Section 3.5) Describe the covalent bond in a hydrogen molecule, H2. (Section 3.5)
ook around you. Do you think you see anything composed of just one element… any objects consisting only of carbon, or of gold, or of hydrogen? The correct answer is almost certainly no. If you are lucky enough to have a diamond ring, you have a piece of carbon that is almost pure (although a gemologist would tell you that diamonds contain slight impurities that give each stone its unique character). If you have a “gold” ring, you have a mixture of gold with other metals, added to give the ring greater strength. Even though a few elements, such as carbon and gold, are sometimes found in elemental form in nature, most of the substances we see around us consist of two or more elements that have combined chemically to form more complex substances called compounds. For example, in nature, the element hydrogen is combined with other elements, such as oxygen and carbon, in compounds such as the water and sugar used to make a soft drink. (Perhaps you are sipping one while you read.) In this chapter, you will learn to (1) define the terms mixture and compound more precisely, (2) distinguish between elements, compounds, and mixtures, (3) describe how elements combine to form compounds, (4) construct systematic names for some chemical compounds, and (5) describe the characteristics of certain kinds of chemical compounds. The chapter will also expand your ability to visualize the basic structures of matter.
The flecks of gold in this pan are the only pure elements visible in this scene.
112 Chapter 4 Chemical Compounds
4.1 Classification of Matter
Before getting started on your chemistry homework, you go into the kitchen to make some pasta for your six‑year‑old nephew. You run water into a pan, adding a few shakes of salt, and while you’re waiting for it to boil, you pour a cup of coffee. When the water begins to boil, you pour in the pasta. Then you add some sugar to your coffee. Pure water, the sucrose in white sugar, and the sodium chloride in table salt are all examples of chemical compounds. A compound is a substance that contains two or more elements, with the atoms of those elements always combining in the same whole‑number ratio (Figure 4.1). There are relatively few chemical elements, but there are millions of chemical compounds. Compounds in our food fuel our bodies, and the compounds in gasoline fuel our cars. They can alter our moods and cure our diseases. Water is composed of molecules that contain two atoms of hydrogen and one atom of oxygen. We describe the composition of water with the chemical formula H2O. White sugar is a highly purified form of sucrose, whose chemical formula is C12H22O11. Its molecules are composed of 12 carbon atoms, 22 hydrogen atoms, and 11 oxygen atoms. Sodium and chlorine atoms combine in a one‑to‑one ratio to form sodium chloride, NaCl, which is the primary ingredient in table salt.
Figure 4. Elements Versus Compounds
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ELEMENTS
COMPOUNDS
Silver exists as an assembly of silver atoms.
Neon is composed of independent atoms.
Hydrogen is composed of molecules with 2 hydrogen atoms.
H 2 molecule Neon atom Silver atom
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Water is composed of molecules that contain one oxygen atom and two hydrogen atoms.
Sodium chloride exists as an assembly of sodium and chloride ions, always in a one-to-one ratio.
Sodium ion
Chloride ion
Water molecule, H 2 O
The following sample study sheet and Figure 4.3 show the questions you can ask to discover whether a sample of matter is an element, a compound, or a mixture.
114 Chapter 4 Chemical Compounds
Figure 4. Classification of Matter
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Sample Study
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Classification
of Matter
Tip-off You are asked to classify a sample of matter as a pure substance or a mixture; or you are asked to classify a pure substance as an element or a compound. General STepS The following general procedure is summarized in Figure 4.3. To classify a sample of matter as a pure substance or a mixture, ask one or both of the following questions: Does it have a constant composition? If it does, it is a pure substance. If it has variable composition, it is a mixture. Can the sample as a whole be described with a chemical formula? If it can, it is a pure substance. If it cannot, it is a mixture. To classify a pure substance as an element or a compound, ask the following question: Can it be described with a single symbol? If it can, it is an element. If its chemical formula contains two or more different element symbols, it is a compound. example See Example 4.1.
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example 4.1 - Classification of Matter
Many of us have a bottle in our medicine cabinet containing a mild disinfectant consisting of hydrogen peroxide and water. The liquid is about 3% hydrogen peroxide, H2O2, and about 97% water. Classify each of the following as a pure substance or a mixture. If it is a pure substance, is it an element or a compound?
a. the liquid disinfectant b. the hydrogen peroxide, H2O (^) 2, used to make the disinfectant c. the hydrogen used to make hydrogen peroxide
Solution
a. We know that the liquid disinfectant is a mixture for two reasons. It is composed of two pure substances (H2O 2 and H2O), and it has variable composition. b. Because hydrogen peroxide can be described with a formula, H2O2, it must be a pure substance. Because the formula contains symbols for two elements, it represents a compound. c. Hydrogen can be described with a single symbol, H or H2, so it is a pure substance and an element.
exerCise 4.1 - Classification of Matter
The label on a container of double‑acting baking powder tells us that it contains cornstarch, bicarbonate of soda (also called sodium hydrogen carbonate, NaHCO3), sodium aluminum sulfate, and acid phosphate of calcium (which chemists call calcium dihydrogen phosphate, Ca(H2PO4)2). Classify each of the following as a pure substance or a mixture. If it is a pure substance, is it an element or a compound?
a. calcium b. calcium dihydrogen phosphate c. double‑acting baking powder
4.2 Compounds and Chemical Bonds
The percentage of H (^) 2O 2 in the mixture of hydrogen peroxide and water that is used as a disinfectant can vary, but the percentage of hydrogen in the compound water is always the same. Why? One of the key reasons that the components of a given compound are always the same, and present in the same proportions, is that the atoms in a compound are joined together by special kinds of attractions called chemical bonds. Because of the nature of these attractions, the atoms combine in specific ratios that give compounds their constant composition. This section will introduce you to the different types of chemical bonds and provide you with the skills necessary to predict the types of chemical bonds between atoms of different elements.
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4.2 Compounds and Chemical Bonds 117
Transfer of Electrons
Sometimes one atom in a bond attracts electrons so much more strongly than the
other that one or more electrons are fully transferred from one atom to another.
This commonly happens when metallic atoms combine with nonmetallic atoms. A
nonmetallic atom usually attracts electrons so much more strongly than a metallic
atom that one or more electrons shift from the metallic atom to the nonmetallic atom.
For example, when the element sodium combines with the element chlorine to form
sodium chloride, NaCl, the chlorine atoms attract electrons so much more strongly
than the sodium atoms that one electron is transferred from each sodium atom to a
chlorine atom.
When an electron is transferred completely from one uncharged atom to another,
the atom that loses the electron is left with one more proton than electron and acquires
a +1 charge overall. It therefore becomes a cation (Section 3.4). For example, when
an uncharged sodium atom with 11 protons and 11 electrons loses an electron, it is
left with 11 protons (a charge of +11) and 10 electrons (a charge of −10), yielding an
overall +1 charge.
Na → Na+^ + e− 11p/11e−^ 11p/10e−
- 11 + (−11) = 0 + 11 + (−10) = + 1 In contrast, an uncharged atom that gains an electron will have one more electron
than proton, so it forms an anion with a −1 charge. When a chlorine atom gains an
electron from a sodium atom, the chlorine atom changes from an uncharged atom
with 17 protons and 17 electrons to an anion with 17 protons and 18 electrons and an overall −1 charge.
Cl + e−^ → Cl− 17p/17e−^ 17p/18e−
- 17 + (−17) = 0 + 17 + (−18) = − Atoms can transfer one, two, or three electrons. Thus cations can have
a +1, +2, or +3 charge, and anions can have a −1, −2, or −3 charge.
Because particles with opposite charges attract each other, there is an
attraction between cations and anions. This attraction is called an ionic
bond. For example, when an electron is transferred from a sodium atom
to a chlorine atom, the attraction between the +1 sodium cation and the
−1 chlorine anion is an ionic bond (Figure 4.5).
You will see as you read more of this book that substances that have
ionic bonds are very different from those that have all covalent bonds. For
example, compounds that have ionic bonds, such as the sodium chloride
in table salt, are solids at room temperature and pressure, but compounds
with all covalent bonds, such as hydrogen chloride and water, can be gases
and liquids as well as solids.
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The Salt-Encrusted Shore of The Dead Sea Salt (sodium chloride) is an ionic compound. Water is molecular.
118 Chapter 4 Chemical Compounds
Summary of Covalent and Ionic Bond Formation
When atoms of different elements form chemical bonds, the electrons in the bonds can shift from one bonding atom to another. The atom that attracts electrons more strongly will acquire a negative charge, and the other atom will acquire a positive charge. The more the atoms differ in their electron‑attracting ability, the more the electron cloud shifts from one atom toward another. If there is a large enough difference in electron‑attracting ability, one, two, or three electrons can be viewed as shifting completely from one atom to another. The atoms become positive and negative ions, and the attraction between them is called an ionic bond. If the electron transfer is significant but not enough to form ions, the atoms acquire partial positive and partial negative charges. The bond in this situation is called a polar covalent bond. If there is no shift of electrons or if the shift is negligible, no significant charges will form, and the bond will be a nonpolar covalent bond. It might help, when thinking about these different kinds of bonds, to compare them to a game of tug‑of‑war between two people. The people are like the atoms with a chemical bond between them, and the rope is like the electrons in the bond. If the two people tugging have the same (or about the same) strength, the rope will not move (or not move much). This leads to a situation that is like the nonpolar covalent bond. If one person is stronger than the other person, the rope will shift toward that person, the way the electrons in a polar covalent bond shift toward the atom that attracts them more. If one person can pull a lot harder than the other person can, the stronger person pulls the rope right out of the hands of the weaker one. This is similar to the formation of ions and ionic bonds, when a nonmetallic atom pulls one or more electrons away from a metallic atom.
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Figure 4.7 Classifying Compounds objeCtive 9
Classifying Compounds Compounds can be classified as molecular or ionic. Molecular compounds are composed of molecules, which are collections of atoms held together by all covalent bonds. Ionic compounds contain cations and anions held together by ionic bonds (Figure 4.7). You will see some exceptions later in this text, but for now, if a formula for a compound indicates that all the elements in it are nonmetals, you can assume that all of the bonds are covalent bonds, which form molecules, and that the compound is a molecular compound. We will assume that metal‑nonmetal combinations lead to ionic bonds and ionic compounds.
example 4.2 - Classifying Compounds
Classify each of the following as either a molecular compound or an ionic compound. a. calcium chloride, CaCl 2 (used for de‑icing roads) b. ethanethiol, C (^) 2H5 SH (a foul‑smelling substance used to odorize natural gas) Solution a. Calcium, Ca, is a metal, and chlorine, Cl, is a nonmetal. We expect the bonds between them to be ionic, so calcium chloride is an ionic compound. b. Carbon, hydrogen, and sulfur are all nonmetallic elements, so we expect the bonds between them to be covalent bonds. The formula, C2H5SH, tells us that ethanethiol is composed of molecules that each contain two carbon atoms, six hydrogen atoms, and one sulfur atom. Ethanethiol is a molecular compound.
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exerCise 4.2 - Classifying Compounds
Classify each of the following substances as either a molecular compound or an ionic compound. a. formaldehyde, CH2O (used in embalming fluids) b. magnesium chloride, MgCl 2 (used in fireproofing wood and in paper manufacturing)
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HCl molecule
Covalent bond
Nonmetal Nonmetal
Metallic cations
Nonmetallic anions
Molecular compound Hydrogen chloride, HCl, gas
Ionic compound Sodium chloride, NaCl, solid
4.3 Molecular Compounds 121
4.3 Molecular Compounds
Have you ever wondered why salt dissolves so quickly in water but oil does not?…why bubbles form when you open a soft drink can?… why a glass of water fizzes when an Alka‑Seltzer tablet is plopped into it? What’s going on at the submicroscopic level that makes these things happen? To answer these questions, you need to know more about the structure of water, including the spatial arrangement of atoms in water molecules. The purpose of this section is to begin to describe the three‑dimensional structure of molecular compounds such as water. Earlier we saw that when some elements form ionic and covalent bonds, their atoms gain, lose, or share electrons. This suggests an important role for electrons in chemistry. However, chemists have also found that for most elements, some electrons are more influential in the formation of chemical bonds than others are. Of chlorine’s 17 electrons, for example, only seven are important in predicting how chlorine will bond. Of sulfur’s 16 electrons, only six are important; of phosphorus’s 15 electrons, only five are important. Chemists noticed that the important electrons, called valence electrons , are equal in number to the element’s “A‑group” number. For example, the nonmetallic elements in group 7A (F, Cl, Br, and I) have seven valence electrons, those in group 6A (O, S, and Se) have six valence electrons, those in group 5A (N and P) have five, and carbon (C) in group 4A has four. A more precise definition of valence electrons, and an explanation for why chlorine has seven, sulfur six, and so on, will have to wait until you learn more about atomic theory in Chapter 11. For now, it is enough to know the numbers of valence electrons for each nonmetallic atom and know how they are used to explain the bonding patterns of nonmetallic atoms. The valence electrons for an element can be depicted visually in an electron-dot symbol. (Electron‑dot symbols are known by other names, including electron‑dot structures, electron‑dot diagrams, and Lewis electron‑dot symbols.) An electron‑dot symbol that shows chlorine’s seven valence electrons is
Cl
Electron‑dot symbols are derived by placing valence electrons (represented by dots) to the right, left, top, and bottom of the element’s symbol. Starting on any of these four sides, we place one dot at a time until there are up to four unpaired electrons around the symbol. If there are more than four valence electrons for an atom, the remaining electrons are added one by one to the unpaired electrons to form up to four pairs.
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There is no set convention for the placement of the paired and unpaired electrons around the symbol. For example, the electron‑dot symbol for chlorine atoms could be
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4.3 Molecular Compounds 123
structure for a hydrogen chloride molecule, HCl:
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Like chlorine, the other elements in group 7A also have seven valence electrons, so their electron‑dot symbols are similar to that of chlorine. The unpaired dot can be placed on any of the four sides of each symbol.
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In order to obtain octets of electrons, these atoms tend to form compounds in which they have one bond and three lone pairs. Note how the Lewis structures of hydrogen fluoride, HF (used in the refining of uranium), hydrogen bromide, HBr (a pharmaceutical intermediate), and hydrogen iodide, HI (used to make iodine salts) resemble the structure of hydrogen chloride.
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The nonmetallic elements in group 6A (oxygen, sulfur, and selenium) have atoms with six valence electrons:
O S Se
(The unpaired dots can be placed on any two of the four sides of each symbol.) These elements usually gain an octet by forming two covalent bonds and two lone pairs, as in water, H2O, and hydrogen sulfide, H2S.
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Nitrogen and phosphorus, which are in group 5A, have atoms with five valence electrons:
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They form three covalent bonds to pair their three unpaired electrons and achieve an octet of electrons around each atom. Ammonia, NH3, and phosphorus trichloride, PCl3, molecules are examples.
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PCl 3 is used to make pesticides and gasoline additives.
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124 Chapter 4 Chemical Compounds
Carbon, in group 4A, has four unpaired electrons in its electron‑dot symbol.
Predictably, carbon atoms are capable of forming four covalent bonds (with no lone pairs). Examples include methane, CH4, the primary component of natural gas, and ethane, C2H (^) 6, and propane, C3H8, which are also found in natural gas, but in smaller quantities.
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Methane, ethane, and propane are hydrocarbons , compounds that contain only carbon and hydrogen. Fossil fuels that we burn to heat our homes, cook our food, and power our cars, are primarily hydrocarbons. For example, natural gas is a mixture of hydrocarbons with from one to four carbons, and gasoline contains hydrocarbon molecules with from six to twelve carbons. Like the hydrocarbons described above, many of the important compounds in nature contain a backbone of carbon‑carbon bonds. These compounds are called organic compounds, and the study of carbon‑ based compounds is called organic chemistry.
Figure 4. Household Hydrocarbon Liquid petroleum gas is a mixture of the hydrocarbons propane and butane.
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Table 4.1 shows electron‑dot symbols for the nonmetallic atoms and lists their most common bonding patterns. Note that the sum of the numbers of bonds and lone pairs is always four for the elements in this table.
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When drawing its Lewis structure, we assume that the carbon atom will have four bonds (represented by four lines), the oxygen atom will have two bonds and two lone pairs, and each hydrogen atom will have one bond. The Lewis structure below meets these criteria.
Methanol, CH3OH (methyl alcohol)
Methanol is an alcohol, which is a category of organic compounds, not just the intoxicating compound in certain drinks. Alcohols are organic compounds that possess one or more –OH groups attached to a hydrocarbon group (a group that contains only carbon and hydrogen). Ethanol, C (^) 2H5OH, is the alcohol in alcoholic beverages (see Special Topic 4.1: Intoxicating Liquids and the Brain ), while the alcohol in rubbing alcohol is usually 2‑propanol (Figure 4.9). These alcohols are also called methyl alcohol, CH3OH, ethyl alcohol, C2H5OH, and isopropyl alcohol, C3H7OH.
Ethanol, C2H (^) 5OH 2 ‑propanol, C3H7OH (ethyl alcohol) (isopropyl alcohol)
Check to see that each of these compounds follows our guidelines for drawing Lewis structures.
Figure 4. Products Containing Alcohols
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exerCise 4.3 - Drawing Lewis Structures from Formulas
Draw a Lewis structure for each of the following formulas:
a. nitrogen triiodide, NI 3 (explodes at the slightest touch) b. hexachloroethane, C (^) 2Cl 6 (used to make explosives) c. hydrogen peroxide, H2O2 (a common antiseptic) d. ethylene (or ethene), C2H4 (used to make polyethylene)
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example 4.3 - Drawing Lewis Structures from Formulas
Draw a Lewis structure for each of the following formulas:
a. phosphine, PH 3 (used to make semiconductors) b. hypochlorous acid, HOCl (used to bleach textiles) c. CFC‑11, CCl (^) 3F (used as a refrigerant) d. C (^) 2H2, acetylene (burned in oxyacetylene torches)
Solution
a. Phosphorus atoms usually have three covalent bonds and one lone pair, and hydrogen atoms have one covalent bond and no lone pairs. The following Lewis structure for PH 3 gives each of these atoms its most common bonding pattern.
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b. Hydrogen atoms have one covalent bond and no lone pairs, oxygen atoms usually have two covalent bonds and two lone pairs, and chlorine atoms usually have one covalent bond and three lone pairs.
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c. Carbon atoms usually have four covalent bonds and no lone pairs. Fluorine and chlorine atoms usually have one covalent bond and three lone pairs. The fluorine atom can be put in any of the four positions around the carbon atom.
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d. Carbon atoms form four bonds with no lone pairs, and hydrogen atoms form one bond with no lone pairs. To achieve these bonding patterns, there must be a triple bond between the carbon atoms.
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4.3 Molecular Compounds 129
Figure 4. Three Ways to Describe a Methane Molecule
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The nitrogen atom in an ammonia molecule, NH3, forms three covalent bonds and in addition has a lone pair of electrons. A lone pair on a central atom must be considered in predicting a molecule’s shape.
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Like the carbon atom in a methane molecule, the nitrogen atom has four electron‑groups around it, so the ammonia molecule has a shape that is very similar to the shape of a CH 4 molecule. However, the lone pair on the nitrogen atom repels neighboring electron‑groups more strongly than the bond pairs do, so the lone pair in the ammonia molecule pushes the bond pairs closer together than the bond pairs for methane. The bond angle is about 107° instead of 109.5°. Figure 4.11 shows three ways to represent the ammonia molecule.
Figure 4. Three Ways to Describe an Ammonia Molecule
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Liquid Water
A chemist’s‑eye view of the structure of liquid water starts with the prediction of the molecular shape of each water molecule. The Lewis structure of water shows that the oxygen atom has four electron‑groups around it: two covalent bonds and two lone pairs.
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We predict that the four groups would be distributed in a tetrahedral arrangement to keep their negative charges as far apart as possible. Because the lone pairs are more repulsive than the bond pairs, the angle between the bond pairs is less than 109.5°. In fact, it is about 105° (Figure 4.12).
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Figure 4. Three ways to Describe a Water Molecule
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Because oxygen atoms attract electrons much more strongly than do hydrogen atoms, the O‑H covalent bond is very polar, leading to a relatively large partial minus charge on the oxygen atom (represented by a δ−) and a relatively large partial plus charge on the hydrogen atom (represented by a δ+).
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The attraction between the region of partial positive charge on one water molecule and the region of partial negative charge on another water molecule tends to hold water molecules close together (Figure 4.13). Remember that opposite charges attract each other and like charges repel each other.
Figure 4. Attractions Between water Molecules
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As in other liquids, the attractions between water molecules are strong enough to keep them the same average distance apart, but they are weak enough to allow each molecule to be constantly breaking the attractions that momentarily connect it to some molecules and forming new attractions to other molecules (Figure 4.14). In other chapters, you will find this image of the structure of water useful in developing your understanding of what is happening when salt dissolves in your pasta water and when bubbles form in a soft drink or in a glass of Alka‑Seltzer and water.
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