Chemistry: Chapter Three
Composition of Compounds
A chemical formula, in combination with the molar masses of its constituent elements, indicates the relative quantities of each element in a compound, which is extremely useful information. For example, about 35 years ago, scientists began to suspect that synthetic compounds known as chlorofluorocarbons (or CFCs) were destroying ozone (O3) in Earth's upper atmosphere. Upper atmospheric ozone is important because it shields life on Earth from the sun's harmful ultraviolet light. CFCs are chemically inert compounds that were used primarily as refrigerants and industrial solvents. Over time, however, CFCs began to accumulate in the atmosphere. In the upper atmosphere sunlight breaks bonds within CFCs, resulting in the release of chlorine atoms. The chlorine atoms then react with ozone, converting it into O2. Therefore, the harmful part of CFCs is the chlorine atoms that they carry. How can we determine the mass of chlorine in a given mass of a CFC? One way we can express how much of an element is in a given compound is by using the element's mass percent composition for that compound. The mass percent composition, or more simply mass percent of an element, is that element's percentage of the compound's total mass. We calculate the mass percent of element X in a compound from the chemical formula as follows: mass percent of element x = mass of element x in 1 mol of compound/mass of 1 mol of the compound x 100% Suppose, for example, that we want to calculate the mass percent composition of Cl in the chlorofluorocarbon CCl2F2. The mass percent Cl is Mass percent Cl = 2 × Molar mass Cl × 100% Molar mass CCl2F2 CCl2F2 We multiply the molar mass of Cl by two because the chemical formula has a sub- script of 2 for Cl, indicating that 1 mol of CCl2F2 contains 2 mol of Cl atoms. We calculate the molar mass of CCl2F2 as follows: Molar mass = 12.01 g/mol + 2(35.45 g/mol) + 2(19.00 g/mol) = 120.91 g/mol So the mass percent of Cl in CCl2F2 is 58.64%
Molecular Models
A molecular model is a more accurate and complete way to specify a compound. Ball-and-stick models represent atoms as balls and chemical bonds as sticks; how the two connect reflects a molecule's shape. The balls are normally color-coded to specific elements. For example, carbon is customarily black, hydrogen is white, nitrogen is blue, and oxygen is red. (For a complete list of colors of elements in the molecular models used in this book see Appendix IIA.) In space-filling molecular models, atoms fill the space between each other to more closely represent our best estimates for how a molecule might appear if scaled to a visible size. For example, consider the following ways to represent a molecule of methane, the main component of natural gas: The molecular formula of methane shows the number and type of each atom in the molecule: one carbon atom and four hydrogen atoms. The structural formula indicates how the atoms are connected: The carbon atom is bonded to the four hydrogen atoms. The ball-and-stick model clearly shows the geometry of the molecule: The carbon atom sits in the center of a tetrahedron formed by the four hydrogen atoms. The space-filling model gives the best sense of the relative sizes of the atoms and how they merge together in bonding. Throughout this book, you will see molecules represented in all of these ways. As you look at these representations, keep in mind what you learned in Chapter 1: The details about a molecule—the atoms that compose it, the lengths of the bonds between atoms, the angles of the bonds between atoms, and its overall shape—determine the properties of the substance that the molecule composes. If any of these details were to change, the properties of the substance would change. Table 3.1 (on the next page) shows various compounds represented in the different ways we have just discussed.
Naming Binary Ionic Compounds with only one type of Cationn
Binary compounds contain only two different elements. The names for binary ionic compounds take the form: name of cation(metal) + base name of anion (nonmetal) + -ide For example, the name for KCl consists of the name of the cation, potassium, fol- lowed by the base name of the anion, chlor, with the ending -ide. Its full name is potassium chloride. KCl potassium chloride The name for CaO consists of the name of the cation, calcium, followed by the base name of the anion, ox, with the ending -ide. Its full name is calcium oxide. CaO calcium oxide The base names for various nonmetals and their most common charges in ionic compounds are shown in Table 3.2. More than one kind of Cation For these types of metals, the name of the cation is followed by a Roman numeral (in parentheses) indicating its charge in that particular compound. For example, we distinguish between Fe2+ and Fe3+ as follows: Fe2+ iron(II) Fe3+ iron(III) The full names therefore have the form name of cation(metal) (charge of cation (metal) in Roman numerals in parentheses) base name of anion (nonmetal) + -ide The charge of the metal cation is obtained by inference from the sum of the charges of the nonmetal anions—remember that the sum of all the charges must be zero. Table 3.3 shows some of the metals that form more than one cation and the values of their most common charges. For example, in CrBr3, the charge of chromium must be 3+ in order for the compound to be charge-neutral with three Br- anions. The cation is therefore named as Cr3+ chromium(III) The full name of the compound is CrBr3 chromium(III) bromide Similarly, in CuO, the charge of copper must be 2+ in order for the compound to be neutral with one O2- anion. The cation is therefore named as follows: Cu2+ copper(II) The full name of the compound is CuO copper(II) oxide Naming Ionic Compounds Containing Polyatomic Ions: We name ionic compounds containing polyatomic ions in the same way as other ionic compounds, except that we use the name of the polyatomic ion whenever it occurs. Tables 3.4 lists common polyatomic ions and their formulas. For example, NaNO2 is named according to its cation, Na+, sodium, and its polyatomic anion, NO2-, nitrite. Its full name is sodium nitrite. NaNO2 sodium nitrite FeSO4 is named according to its cation, iron, its charge (II), and its polyatomic ion sul- fate. The full name is iron(II) sulfate. FeSO4 iron(II) sulfate If the compound contains both a polyatomic cation and a polyatomic anion, use the names of both polyatomic ions. For example, NH4NO3 is ammonium nitrate. NH4NO3 ammonium nitrate You must be able to recognize polyatomic ions in a chemical formula, so become familiar with Table 3.4. Most polyatomic ions are oxyanions, anions containing oxygen and another element. Notice that when a series of oxyanions contains different numbers of oxygen atoms, they are named systematically according to the number of oxygen atoms in the ion. If there are only two ions in the series, the one with more oxygen atoms is given the ending -ate and the one with fewer is given the ending -ite. For example, NO3- is nitrate and NO2- is nitrite. If there are more than two ions in the series, then the prefixes hypo-, meaning less than, and per-, meaning more than, are used. So ClO- is hypochlorite, meaning less oxygen than chlorite, and ClO4 - is perchlorate, meaning more oxygen than chlorate. ClO- hypochlorite ClO2 - chlorite ClO3 - chlorate ClO4 - perchlorate
Types of Chemical Formula
Chemical formulas are generally divided into three different types: empirical, molecular, and structural. An empirical formula simply lists the relative number of atoms of each element in a compound. A molecular formula lists the actual number of atoms of each element in a molecule of a compound. For example, the empirical formula for hydrogen peroxide is HO, but its molecular formula is H2O2. The molecular formula is always a whole-number multiple of the empirical formula. For some compounds, the empirical formula and the molecular formula are identical. For example, both the empirical and molecular formulas for water are H2O because water molecules contain 2 hydrogen atoms and 1 oxygen atom, and no simpler whole-number ratio can express the relative number of hydrogen atoms to oxygen atoms. A structural formula, which uses lines to represent covalent bonds, shows how atoms in a molecule are connected or bonded to each other. The structural formula for H2O2 is H¬O¬O¬H We can also write structural formulas to give a sense of the molecule's geometry. For example, the structural formula for hydrogen peroxide can be written: hydrogen peroxide can be written: H \ O--- O = H Writing the formula this way shows the approximate angles between bonds, giving a sense of the molecule's shape. Structural formulas can also show different types of bonds that occur between molecules. For example, the structural formula for carbon dioxide is O=C=O The two lines between the carbon and oxygen atoms in this formula represent a double bond, which is generally stronger and shorter than a single bond (represented by a single line). A single bond corresponds to one shared electron pair, while a double bond corresponds to two shared electron pairs. We will discuss single, double, and even triple bonds in more detail in Chapter 9. The type of formula we use depends on how much we know about the compound and how much we want to communicate. Notice that a structural formula communicates the most information, while an empirical formula communicates the least.
Chemical Bonds
Compounds are composed of atoms held together by chemical bonds. Chemical bonds form because of the attractions between the charged particles—electrons and protons— that compose atoms. We can broadly classify most chemical bonds into two types: ionic and covalent. Ionic bonds—which occur between metals and nonmetals—involve the transfer of electrons from one atom to another. Covalent bonds—which occur between two or more nonmetals—involve the sharing of electrons between two atoms. Ionic Bonds Recall from Chapter 2 that metals have a tendency to lose electrons and that nonmetals have a tendency to gain them. Therefore, when a metal interacts with a nonmetal, it can transfer one or more of its electrons to the nonmetal. The metal atom then becomes a cation (a positively charged ion), and the nonmetal atom becomes an anion (a negatively charged ion) as shown in FiguRE 3.1▼. These oppositely charged ions are attracted to one another by electrostatic forces, and they form an ionic bond. The result is an ionic compound, which in the solid phase is composed of a lattice—a regular three-dimensional array—of alternating cations and anions.
Organic Compounds
Early chemists divided compounds into two types: organic and inorganic. Organic compounds came from living things. Sugar—obtained from sugarcane or the sugar beet—is a common example of an organic compound. Inorganic compounds, on the other hand, came from the Earth. Salt—mined from the ground or from the ocean—is a common example of an inorganic compound. Eighteenth-century chemists could synthesize inorganic compounds in the laboratory, but not organic compounds, so a clear division existed between the two different types of compounds. Today, chemists can synthesize both organic and inorganic compounds, and even though organic chemistry is a subfield of chemistry, the differences between organic and inorganic compounds are primarily organizational, not fundamental. Organic compounds are composed of carbon and hydrogen and a few other elements, including nitrogen, oxygen, and sulfur. The key element in organic chemistry, however, is carbon. Carbon always forms four bonds in compounds. For example, the simplest organic compound is methane, CH4. The chemistry of carbon is unique and complex because carbon frequently bonds to itself to form chain, branched, and ring structures. This versatility allows carbon to be the backbone of millions of different chemical compounds, which is why even a survey of organic chemistry requires a year-long course. For now, all you really need to know is that the simplest organic compounds are called hydrocarbons and that they are composed of only carbon and hydrogen. Hydrocarbons compose common fuels such as oil, gasoline, liquid propane gas, and natural gas. Table 3.6 lists some common hydrocarbons and their names.
Hydrogen, Oxygen, Water
Hydrogen (H2) is an explosive gas used as a fuel in rocket engines. Oxygen (O2), also a gas, is a natural component of air. Oxygen is not itself flammable but must be present for combustion (burning) to occur. Hydrogen and oxygen both have extremely low boiling points (see the table on the next page). When hydrogen and oxygen combine to form the compound water (H2O), however, a dramatically different substance results. First of all, water is a liquid rather than a gas at room temperature, and its boiling point is hundreds of degrees higher than the boiling points of hydrogen and oxygen. Second, instead of being flammable (like hydrogen gas) or supporting combustion (like oxygen gas), water actually smothers flames. Water is nothing like the hydrogen and oxygen from which it was formed. The dramatic difference between the elements hydrogen and oxygen and the compound water is typical of the differences between elements and the compounds that they form. When two elements combine to form a compound, an entirely new substance results. Common table salt, for example, is a stable compound composed of sodium and chlorine. Elemental sodium, by contrast, is a highly reactive, silvery metal that can explode on contact with water, and elemental chlorine is a corrosive, greenish-yellow gas that can be fatal if inhaled. Yet the compound that results from the combination of these two elements is sodium chloride (or table salt), a flavor enhancer that we sprinkle on our food. Although some of the substances that we encounter in everyday life are elements, most are compounds. Free atoms are rare on Earth. As we discussed in Chapter 1, a compound is different from a mixture of elements. In a compound, elements combine in fixed, definite proportions; in a mixture, elements can mix in any proportions. For example, consider the difference between a hydrogen-oxygen mixture and water. A hydrogen-oxygen mixture can have any proportion of hydrogen and oxygen gas. Water, by contrast, is composed of water molecules that always contain two hydrogen atoms to every one oxygen atom. Water has a definite proportion of hydrogen to oxygen. In this chapter you will learn about compounds: how to represent them, how to name them, how to distinguish among their different types, and how to write chemical equations showing how they form and change. You will also learn how to quantify the composition of a compound according to its constituent elements. This is important whenever you want to know how much of a particular element is contained within a particular compound. For example, patients with high blood pressure often have to reduce their sodium ion intake. Since the sodium ion is normally consumed in the form of sodium chloride, a high-blood-pressure patient needs to know how much sodium is in a given amount of sodium chloride. Similarly, an iron-mining company needs to know how much iron it can recover from a given amount of iron ore. This chapter provides the tools to understand and solve these kinds of problems.
Molar Mass of a Compound
In Chapter 2 (Section 2.8), we saw that an element's molar mass—the mass in grams of one mole of its atoms—is numerically equivalent to its atomic mass. We then used the molar mass in combination with Avogadro's number to determine the number of atoms in a given mass of the element. The same concept applies to compounds. The molar mass of a compound—the mass in grams of 1 mol of its molecules or formula units—is numerically equivalent to its formula mass. For example, we just calculated the formula mass of CO2 to be 44.01 amu. The molar mass is, therefore, CO2 molar mass = 44.01 g/mol Using Molar Mass to Count Molecules by Weighing The molar mass of CO2 provides us with a conversion factor between mass (in grams) and amount (in moles) of CO2. Suppose we want to find the number of CO2 molecules in a sample of dry ice (solid CO2) with a mass of 10.8 g. This calculation is analogous to Example 2.8, where we found the number of atoms in a sample of copper of a given mass. We begin with the mass of 10.8 g and use the molar mass to convert to the amount in moles. Then we use Avogadro's number to convert to number of molecules. The conceptual plan is as follows: g CO2 ---> mol CO2 ---> molecules
Formula Mass and Mole Concept for Compounds
In Chapter 2, we defined the average mass of an atom of an element as the atomic mass for that element. Similarly, we now define the average mass of a molecule (or a formula unit) of a compound as the formula mass for that compound. The terms molecular mass and molecular weight have the same meaning as formula mass. For any compound, the formula mass is the sum of the atomic masses of all the atoms in its chemical formula. Formula = (Number of atoms in 1st element in chemical formula x Atomic mass of 1st element) + (Number of atoms of 2nd element in chemical formula x Atomic mass of 2nd element) For example, the formula mass of carbon dioxide, CO2, is Formula mass = 12.01 amu + 2(16.00 amu) = 44.01 amu and that of sodium oxide, Na2O, is Formula mass = 2(22.99 amu) + 16.00 amu = 61.98 amu
Determining a Chemical Formula from Experimental Data
In Section 3.8, we calculated mass percent composition from a chemical formula. Can we also do the reverse? Can we calculate a chemical formula from mass percent composition? This question is important because laboratory analyses of compounds do not often give chemical formulas directly, but only the relative masses of each element present in a compound. For example, if we decompose water into hydrogen and oxygen in the laboratory, we can measure the masses of hydrogen and oxygen produced. Can we determine a chemical formula from these kind of data? The answer is a qualified yes. We can calculate a chemical formula, but it is an empirical formula (not a molecular formula). To arrive at a molecular formula, we need additional information, such as the molar mass of the compound. Suppose we decompose a sample of water in the laboratory and it produces 0.857 g of hydrogen and 6.86 g of oxygen. How do we determine an empirical formula from these data? We know that an empirical formula represents a ratio of atoms or moles of atoms, but not a ratio of masses. So the first thing we must do is convert our data from mass (in grams) to amount (in moles). How many moles of each element are present in the sample? To convert to moles, use the molar mass of each element: Moles H = 0.857 g H * 1 mol H/1.008 g H = 0.850 mol H Moles O = 6.86 g O * 1 mol O/16.00 g O = 0.429 mol O From these data, we know there are 0.850 mol H for every 0.429 mol O. We can now write a pseudoformula for water: Hydrogen 0.850 Oxygen 0.429 To get the smallest whole-number subscripts in our formula, we divide all the subscripts by the smallest one, in this case 0.429: H0.850/0.429 O0.429/0.429 = H1.98O = H2O Our empirical formula for water, which also happens to be the molecular formula, is H2O. We can use the following procedure to obtain the empirical formula of any compound from experimental data giving the relative masses of the constituent elements. The left column outlines the procedure, and the examples in the center and right column demonstrate how to apply the procedure.
Molecular Compounds: Formulas and Names
In contrast to ionic compounds, the formula for a molecular compound cannot easily be determined based on its constituent elements because the same group of elements may form many different molecular compounds, each with a different formula. Recall from Chapter 1, for example, that carbon and oxygen form both CO and CO2, and that hydrogen and oxygen form both H2O and H2O2. Nitrogen and oxygen form all of the following molecular compounds: NO, NO2, N2O, N2O3, N2O4, and N2O5. In Chapter 9, we will discuss the stability of these various combinations of the same elements. For now, we focus on naming a molecular compound based on its formula or writing its formula based on its name. Naming Molecular Compounds: Like ionic compounds, many molecular compounds have common names. For example, H2O and NH3 have the common names water and ammonia, respectively. However, the sheer number of existing molecular compounds—numbering in the millions—demands a systematic approach to naming them. The first step in naming a molecular compound is identifying it as one. Remember, molecular compounds form between two or more nonmetals. In this section, we learn how to name binary (two-element) molecular compounds. Their names have the form. When writing the name of a molecular compound, as when writing the formula, we first list the more metal-like element (toward the left and bottom of the periodic table). We always write the name of the element with the smallest group number first. If the two elements lie in the same group, then we write the element with the greatest row number first. The prefixes given to each element indicate the number of atoms present: mono = 1 hexa = 6 di = 2 hepta = 7 tri = 3 octa = 8 tetra = 4 nona = 9 penta = 5 deca = 10 If there is only one atom of the first element in the formula, the prefix mono- is normally omitted. For example, we name NO2 according to the first element, nitrogen, with no prefix because mono- is omitted for the first element, followed by the prefix di, to indicate two oxygen atoms, followed by the base name of the second element, ox, with the ending -ide. Its full name is nitrogen dioxide. NO2 nitrogen dioxide We name the compound N2O, sometimes called laughing gas, similarly except that we use the prefix di- before nitrogen to indicate two nitrogen atoms and the prefix mono- before oxide to indicate one oxygen atom. Its full name is dinitrogen monoxide. N2O dinitrogen monoxide
Combustion Analysis
In the previous section, we discussed how to determine the empirical formula of a compound from the relative masses of its constituent elements. Another common (and related) way to obtain empirical formulas for unknown compounds, especially those containing carbon and hydrogen, is combustion analysis. In combustion analysis, the unknown compound under- goes combustion (or burning) in the presence of pure oxygen, as shown in FiguRE 3.9▶ on the next page. All of the carbon in the sample is converted to CO2, and all of the hydrogen is converted to H2O. The CO2 and H2O produced are weighed, and the numerical relationships between moles inherent in the formulas for CO2 and H2O (1 mol CO2 : 1 mol C and 1 mol H2O : 2 mol H) are used to determine the amounts of C and H in the original sample. Any other elemental constituents, such as O, Cl, or N, can be determined by subtracting the original mass of the sample from the sum of the masses of C and H. The examples that follow demonstrate these calculations for a sample containing only C and H and for a sample containing C, H, and O.
Ionic Compounds: Formulas and Names
Ionic compounds occur throughout Earth's crust as minerals. Examples include limestone (CaCO3), a type of sedimentary rock; gibbsite 3Al(OH)3 4 , an aluminum- containing mineral; and soda ash (Na2CO3), a natural deposit. Ionic compounds are also in the foods that we eat. Examples include table salt (NaCl), the most common flavor enhancer, calcium carbonate (CaCO3), a source of calcium necessary for bone health, and potassium chloride (KCl), a source of potassium necessary for fluid balance and muscle function. Ionic compounds are generally very stable because the attractions between cations and anions within ionic compounds are strong, and because each ion interacts with several oppositely charged ions in the crystalline lattice. Writing Formulas for Ionic Compounds Since ionic compounds are charge-neutral, and since many elements form only one type of ion with a predictable charge, we can deduce the formulas for many ionic compounds from their constituent elements. For example, the formula for the ionic compound composed of sodium and chlorine must be NaCl because, in compounds, Na always forms 1+ cations and Cl always forms 1- anions. In order for the compound to be charge- neutral, it must contain one Na+ cation to every one Cl- anion. The formula for the ionic compound composed of calcium and chlorine can only be CaCl2 because Ca always forms 2+ cations and Cl always forms 1- anions. In order for this compound to be charge-neutral, it must contain one Ca2+ cation to every two Cl- anions. Summarizing Ionic Compound Formulas: ▶ Ionic compounds always contain positive and negative ions. ▶ In a chemical formula, the sum of the charges of the positive ions (cations) must equal the sum of the charges of the negative ions (anions). ▶ The formula reflects the smallest whole-number ratio of ions. To write the formula for an ionic compound, follow the procedure in the left column of the following table. Examples 3.3 and 3.4, shown in the center and right columns, illustrate how to apply the procedure.
Conversion Factors from Chemical Formulas
Mass percent composition is one way to understand how much chlorine is in a par- ticular chlorofluorocarbon or, more generally, how much of a constituent element is present in a given mass of any compound. However, we can also approach this question another way. Chemical formulas indicate the inherent relationships between atoms (or moles of atoms) and molecules (or moles of molecules). For example, the formula for CCl2F2 tells us that 1 mol of CCl2F2 contains 2 mol of Cl atoms. We write the ratio as follows: 1 mol CCl2F2:2 mol Cl With ratios such as these—that come from the chemical formula—we can directly determine the amounts of the constituent elements present in a given amount of a compound without calculating mass percent composition. For example, we calculate the number of moles of Cl in 38.5 mol of CCl2F2 as follows: We often want to know, however, not the amount in moles of an element in a certain number of moles of compound, but the mass in grams (or other units) of a constituent element in a given mass of the compound. For example, suppose we want to know the mass (in grams) of Cl contained in 25.0 g CCl2F2. The relationship inherent in the chemical formula (2 mol Cl:1 mol CCl2F2) applies to amount in moles, not to mass. Therefore, we must first convert the mass of CCl2F2 to moles CCl2F2. Then we use the conversion factor from the chemical formula to convert to moles Cl. Finally, we use the molar mass of Cl to convert to grams Cl. The calculation proceeds as follows: Notice that we must convert g CCl2F2 to mol CCl2F2 before we can use the chemical formula as a conversion factor. The general form for solving problems where we are asked to find the mass of an element present in a given mass of a compound is Mass compound S moles compound S moles element S mass element We use the molar mass to convert between mass and moles, and we use relationships inherent in the chemical to convert between moles and moles.
An Atomic-Level View of Elements and Compounds
Recall from Chapter 1 that we can categorize pure substances as either elements or compounds. We can further subcategorize elements and compounds according to the basic units that compose them, as shown in FiguRE 3.3▶. Elements may be either atomic or molecular. Compounds may be either molecular or ionic. Atomic elements are elements that exist in nature with single atoms as their basic units. Most elements fall into this category. For example, helium is composed of helium atoms, aluminum is composed of aluminum atoms, and iron is composed of iron atoms. Molecular elements do not normally exist in nature with single atoms as their basic units; instead, they exist as molecules—two or more atoms of the element bonded together. Most molecular elements exist as diatomic molecules. For example, hydrogen is composed of H2 molecules, nitrogen is composed of N2 molecules, and chlorine is composed of Cl2 molecules. A few molecular elements exist as polyatomic molecules. Phosphorus, for example, exists as P4 and sulfur exists as S8. The elements that exist primarily as diatomic or polyatomic molecules are shown in the periodic table in FiguRE 3.4▲. Molecular compounds are usually composed of two or more covalently bonded nonmetals. The basic units of molecular compounds are molecules composed of the constituent atoms. For example, water is composed of H2O molecules, dry ice is composed of CO2 molecules, and propane (often used as a fuel for grills) is composed of C3H8 molecules as shown in the diagram. Ionic compounds are usually composed of a metal ionically bonded to one or more nonmetals. The basic unit of an ionic compound is the formula unit, the smallest electrically neutral collection of ions. A formula unit is not a molecule—it does not usually exist as a discrete entity, but rather as part of a larger lattice. The ionic compound table salt, for example, which has the formula unit NaCl, is composed of Na+ and Cl- ions in a one-to-one ratio. In table salt, Na+ and Cl- ions exist in a three-dimensional array. However, because ionic bonds are not directional, no one Na+ ion pairs with a specific Cl- ion. Rather, as you can see in FiguRE 3.5(b)◀, any individual Na+ cation is surrounded by Cl- anions and vice versa. Many common ionic compounds contain ions that are themselves composed of a group of covalently bonded atoms with an overall charge. For example, the active ingredient in household bleach is sodium hypochlorite, which acts to chemically alter color- causing molecules in clothes (bleaching action) and to kill bacteria (disinfection). Hypochlorite is a polyatomic ion—an ion composed of two or more atoms—with the formula ClO-. (Note that the charge on the hypochlorite ion is a property of the whole ion, not just the oxygen atom. This is true for all polyatomic ions.) The hypochlorite ion is often found as a unit in other compounds as well [such as KClO and Mg(ClO)2]. Other polyatomic ion-containing compounds found in everyday products include sodium bicarbonate (NaHCO3), also known as baking soda; sodium nitrite (NaNO2), an inhibitor of bacterial growth in packaged meats; and calcium carbonate (CaCO3), the active ingredient in antacids such as Tums®.
Hydrated Ionic Compounds
Some ionic compounds—called hydrates—contain a specific number of water molecules associated with each formula unit. For example, Epsom salts has the formula MgSO4 x7 H2O and the systematic name magnesium sulfate heptahydrate. The seven H2O molecules associated with the formula unit are waters of hydration. Waters of hydration can usually be removed by heating the compound. FiguRE 3.8◀ shows a sample of copper(II) sulfate pentahydrate being heated. The hydrate is blue, and the anhydrous salt (the salt without the associated water molecules) is white. Hydrates are named just as other ionic compounds, but they are given the additional name "prefixhydrate," where the prefix indicates the number of water molecules associated with each formula unit. Some other common examples of hydrated ionic compounds and their names are as follows: CaSO4 x1/2 H2O calcium sulfate hemihydrate BaCl2 x 2 H2O barium chloride dihydrate CuSO4 x 5 H2O copper(II) sulfate pentahydrate
Naming Ionic Compounds
Some ionic compounds—such as NaCl (table salt) and NaHCO3 (baking soda)—have common names, which are nicknames of sorts that can be learned only through familiarity. However, chemists have developed systematic names for different types of compounds including ionic ones. We can determine systematic names by looking at the chemical formula of a compound. Conversely, we can deduce the formula of a compound from its systematic name. The first step in naming an ionic compound is identifying it as one. Remember, ionic compounds are usually formed between metals and nonmetals; when you see a metal and one or more nonmetals together in a chemical formula, you can assume you have an ionic compound. Ionic compounds can be categorized into two types, depending on the metal in the compound. The first type contains metals whose charge is invariant from one compound to another. In other words, whenever one of these metals forms an ion, the ion always has the same charge. Since the charge of the metal in this first type of ionic compound is always the same, it need not be specified in the name of the compound. Sodium, for instance, has a 1+ charge in all of its compounds. FiguRE 3.6▲ lists these types of metals; their charges can be generally determined from their group number in the periodic table. The second type of ionic compound contains metals whose charges can be different in different compounds—each of these metals can form more than one kind of cation, and the charge must therefore be specified for a given compound. Iron, for instance, has a 2+ charge in some of its compounds and a 3+ charge in others. Metals of this type are often in the section of the periodic table known as the transition metals. However, some transition metals, such as Sc, Zn, and Ag, have the same charge in all of their compounds, and some main-group metals, such as lead and tin, have charges that can vary from one compound to another.
Writing and Balance Chemical Equations
The method of combustion analysis we discussed in Section 3.9 involves a chemical reaction, a process in which one or more substances are converted into one or more different substances. Compounds form and change through chemical reactions. As we have seen, water can be produced by the reaction of hydrogen with oxygen. A combustion reaction is a particular type of chemical reaction in which a substance combines with oxygen to form one or more oxygen-containing compounds. Combustion reactions also emit heat, which provides the energy on which our society depends. The heat produced in the combustion of gasoline, for example, helps expand the gaseous combustion products in a car engine's cylinders, which push the pistons and propel the car. We use the heat released by the combustion of natural gas to cook food and to heat our homes. We represent a chemical reaction with a chemical equation. For example, we represent the combustion of natural gas by the equation shown here: The substances on the left side of the equation are the reactants, and the substances on the right side are the products. We often specify the states of each reactant or product in parentheses next to the formula as follows: CH4(g) + O2(g) S CO2(g) + H2O(g) The (g) indicates that these substances are gases in the reaction. Table 3.5 summarizes the common states of reactants and products and their symbols used in chemical equations. If you look more closely at the equation for the combustion of natural gas, you may notice a problem. CH4 (g) O2(g) ---> CO2 + H2O(g) The left side of the equation has two oxygen atoms, while the right side has three. The reaction as written violates the law of conservation of mass because an oxygen atom formed out of nothing. Notice also that there are four hydrogen atoms on the left and only two on the right. Two hydrogen atoms have vanished, again violating mass conservation. To correct these problems—that is, to write an equation that more closely represents what actually happens—we must balance the equation. We must change the coefficients (the numbers in front of the chemical formulas), not the subscripts (the numbers within the chemical formulas), to ensure that the number of each type of atom on the left side of the equation is equal to the number on the right side. New atoms do not form during a reaction, nor do atoms vanish—matter is always conserved. When we add coefficients to the reactants and products to balance an equation, we change the number of molecules in the equation but not the kind of molecules. The result is a balanced chemical equation. To balance the equation for the combustion of methane, we put the coefficient 2 before O2 in the reactants, and the coefficient 2 before H2O in the products. The equation is now balanced; the numbers of each type of atom on either side of the equation are equal. The balanced equation tells us that 1 CH4 molecule reacts with 2 O2 molecules to form 1 CO2 molecule and 2 H2O molecules. We can verify that the equation is balanced by summing the number of each type of atom on each side of the equation.
Representing Compounds: Chemical Formulas and Molecular Models
The quickest and easiest way to represent a compound is with its chemical formula, which, at minimum, indicates which elements are present in the compound and the relative number of atoms or ions of each. For example, H2O is the chemical formula for water—it indicates that water consists of hydrogen and oxygen atoms in a two-to-one ratio. The formula contains the symbol for each element and a subscript indicating the relative number of atoms of the element. A subscript of 1 is typically omitted. Chemical formulas normally list the more metallic (or more positively charged) elements first, followed by the less metallic (or more negatively charged) elements. Other examples of common chemical formulas include NaCl for sodium chloride, indicating sodium and chloride ions in a one-to-one ratio; CO2 for carbon dioxide, indicating carbon and oxygen atoms in a one-to-two ratio; and CCl4 for carbon tetrachloride, indicating carbon and chlorine in a one-to-four ratio.
Acids
We can define acids in a number of ways, as we will see in Chapter 15. For now, we define acids as molecular compounds that release hydrogen ions (H+) when dissolved in water. Acids are composed of hydrogen, usually written first in their formula, and one or more nonmetals, written second. For example, HCl is a molecular compound that, when dissolved in water, forms H+(aq) and Cl-(aq) ions, where aqueous (aq) means dissolved in water. Therefore, HCl is an acid when dissolved in water. To distinguish between gaseous HCl (which is named hydrogen monochloride because it is a molecular compound) and HCl in solution (which is named as an acid), we write the former as HCl(g) and the latter as HCl(aq). Acids are characterized by their sour taste and their ability to dissolve many metals. Since the acid HCl(aq) is present in stomach fluids, its sour taste is painfully obvious when you vomit. HCl(aq), hydrochloric acid, also dissolves some metals. If you put a strip of zinc into a test tube of HCl(aq), it slowly dissolves as the H+(aq) ions convert the zinc metal into Zn2+(aq) cations. Acids are present in many foods, such as lemons and limes, and in household products, such as bathroom cleaner and Lime-Away®. In this section, we discuss how to name acids; in Chapter 15 you will learn more about their properties. We categorize acids into two types: binary acids and oxyacids. Naming Binary Acids: Binary acids are composed of hydrogen and a nonmetal. The names for binary acids have the form hydro acid + base name of nonmetal + -ic + acid For example, HCl(aq) is hydrochloric acid and HBr(aq) is hydrobromic acid. HCl(aq) hydrochloric acid HBr(aq) hydrobromic acid Naming Oxyacids Oxyacids contain hydrogen and an oxyanion (an anion containing a nonmetal and oxygen). The common oxyanions are listed in the table of polyatomic ions (Table 3.4). For example, HNO3(aq) contains the nitrate (NO3-) ion, H2SO3(aq) contains the sul- fite (SO3 2-) ion, and H2SO4(aq) contains the sulfate (SO4 2-) ion. Notice that these acids are simply a combination of one or more H+ ions with an oxyanion. The number of H+ ions depends on the charge of the oxyanion so that the formula is always charge-neutral. The names of oxyacids depend on the ending of the oxyanion and have the following forms: oxyanions ending with -ate acid oxyanions ending with -ate base name of oxyanion + -ic + acid oxyanions ending with -ite +base name of oxyanion + -ous oxyanions ending with -ite + acid So HNO3(aq) is nitric acid (oxyanion is nitrate), and H2SO3(aq) is sulfurous acid (oxy-anion is sulfite). HNO3(aq) nitric acid H2SO3(aq) sulfurous acid
Calculating Molecular Formulas for Compounds
We can determine the molecular formula of a compound from the empirical formula if we also know the molar mass of the compound. Recall from Section 3.3 that the molecular formula is always a whole-number multiple of the empirical formula. Molecular formula = empirical formula * n, where n = 1, 2, 3, c Suppose we want to find the molecular formula for fructose (a sugar found in fruit) from its empirical formula, CH2O, and its molar mass, 180.2 g/mol. We know that the molecular formula is a whole-number multiple of CH2O: Molecular formula = (CH2O) * n = CnH2nOn We also know that the molar mass is a whole-number multiple of the empirical formula molar mass, the sum of the masses of all the atoms in the empirical formula. Molar mass = empirical formula molar mass * n For a particular compound, the value of n in both cases is the same. Therefore, we can find n by calculating the ratio of the molar mass to the empirical formula molar mass. n = molar mass empirical formula/molar mass For fructose, the empirical formula molar mass is empirical formula molar mass = 12.01 g/mol + 2(1.008 g/mol) + 16.00 g/mol = 30.03 g/mol Therefore, n is n = 180.2 g/mol/30.03 g/mol = 6 We can use this value of n to find the molecular formula. Molecular formula = (CH2O) * 6 = C6H12O6
Covalent Bonds
When a nonmetal bonds with another nonmetal, neither atom transfers its electron to the other. Instead the bonding atoms share some of their electrons. The shared electrons have lower potential energy than they would in the isolated atoms because they interact with the nuclei of both atoms. The bond is a covalent bond, and the covalently bound atoms compose a molecule. Each molecule is an independent entity (it is not covalently bound to other molecules.) Therefore, we call covalently bonded compounds molecular compounds. We can begin to understand the stability of a covalent bond by considering the most stable (or lowest potential energy) configuration of a negative charge interacting with two positive charges (which are separated by some small distance). FiguRE 3.2▼ shows that the lowest potential energy occurs when the negative charge lies between the two positive charges because in this arrangement the negative charge can interact with both positive charges. Similarly, shared electrons in a covalent chemical bond hold the bonding atoms together by attracting the positively charged nuclei of both bonding atoms.
