Chapter 2 Molecules, Ions, and Chemical Formulas

Covalent Molecules and Compounds

Just as an atom is the simplest unit that has the fundamental chemical properties of an element, a molecule is the simplest unit that has the fundamental chemical properties of a covalent compound. Some pure elements exist as covalent molecules. Hydrogen, nitrogen, oxygen, and the halogens occur naturally as the diatomic (“two atoms”) molecules H₂, N₂, O₂, F₂, Cl₂, Br₂, and I₂ (part (a) in Figure 2.1 "Elements That Exist as Covalent Molecules"). Similarly, a few pure elements are polyatomicMolecules that contain more than two atoms. (“many atoms”) molecules, such as elemental phosphorus and sulfur, which occur as P₄ and S₈ (part (b) in Figure 2.1 "Elements That Exist as Covalent Molecules").

Each covalent compound is represented by a molecular formulaA representation of a covalent compound that consists of the atomic symbol for each component element (in a prescribed order) accompanied by a subscript indicating the number of atoms of that element in the molecule. The subscript is written only if the number is greater than 1., which gives the atomic symbol for each component element, in a prescribed order, accompanied by a subscript indicating the number of atoms of that element in the molecule. The subscript is written only if the number of atoms is greater than 1. For example, water, with two hydrogen atoms and one oxygen atom per molecule, is written as H₂O. Similarly, carbon dioxide, which contains one carbon atom and two oxygen atoms in each molecule, is written as CO₂.

Figure 2.1 Elements That Exist as Covalent Molecules

(a) Several elements naturally exist as diatomic molecules, in which two atoms (E) are joined by one or more covalent bonds to form a molecule with the general formula E₂. (b) A few elements naturally exist as polyatomic molecules, which contain more than two atoms. For example, phosphorus exists as P₄ tetrahedra—regular polyhedra with four triangular sides—with a phosphorus atom at each vertex. Elemental sulfur consists of a puckered ring of eight sulfur atoms connected by single bonds. Selenium is not shown due to the complexity of its structure.

Covalent compounds that contain predominantly carbon and hydrogen are called organic compoundsA covalent compound that contains predominantly carbon and hydrogen.. The convention for representing the formulas of organic compounds is to write carbon first, followed by hydrogen and then any other elements in alphabetical order (e.g., CH₄O is methyl alcohol, a fuel). Compounds that consist primarily of elements other than carbon and hydrogen are called inorganic compoundsAn ionic or covalent compound that consists primarily of elements other than carbon and hydrogen.; they include both covalent and ionic compounds. In inorganic compounds, the component elements are listed beginning with the one farthest to the left in the periodic table (see Chapter 32 "Appendix H: Periodic Table of Elements"), such as we see in CO₂ or SF₆. Those in the same group are listed beginning with the lower element and working up, as in ClF. By convention, however, when an inorganic compound contains both hydrogen and an element from groups 13–15, the hydrogen is usually listed last in the formula. Examples are ammonia (NH₃) and silane (SiH₄). Compounds such as water, whose compositions were established long before this convention was adopted, are always written with hydrogen first: Water is always written as H₂O, not OH₂. The conventions for inorganic acids, such as hydrochloric acid (HCl) and sulfuric acid (H₂SO₄), are described in Section 2.5 "Acids and Bases".

Representations of Molecular Structures

Molecular formulas give only the elemental composition of molecules. In contrast, structural formulasA representation of a molecule that shows which atoms are bonded to one another and, in some cases, the approximate arrangement of atoms in space. show which atoms are bonded to one another and, in some cases, the approximate arrangement of the atoms in space. Knowing the structural formula of a compound enables chemists to create a three-dimensional model, which provides information about how that compound will behave physically and chemically.

The structural formula for H₂ can be drawn as H–H and that for I₂ as I–I, where the line indicates a single pair of shared electrons, a single bondA chemical bond formed when two atoms share a single pair of electrons.. Two pairs of electrons are shared in a double bondA chemical bond formed when two atoms share two pairs of electrons., which is indicated by two lines— for example, O₂ is O=O. Three electron pairs are shared in a triple bondA chemical bond formed when two atoms share three pairs of electrons., which is indicated by three lines—for example, N₂ is N≡N (see Figure 2.2 "Molecules That Contain Single, Double, and Triple Bonds"). Carbon is unique in the extent to which it forms single, double, and triple bonds to itself and other elements. The number of bonds formed by an atom in its covalent compounds is not arbitrary. As you will learn in Chapter 8 "Ionic versus Covalent Bonding", hydrogen, oxygen, nitrogen, and carbon have a very strong tendency to form substances in which they have one, two, three, and four bonds to other atoms, respectively (Table 2.1 "The Number of Bonds That Selected Atoms Commonly Form to Other Atoms").

Figure 2.2 Molecules That Contain Single, Double, and Triple Bonds

Hydrogen (H₂) has a single bond between atoms. Oxygen (O₂) has a double bond between atoms, indicated by two lines (=). Nitrogen (N₂) has a triple bond between atoms, indicated by three lines (≡). Each bond represents an electron pair.

The structural formula for water can be drawn as follows:

Because the latter approximates the experimentally determined shape of the water molecule, it is more informative. Similarly, ammonia (NH₃) and methane (CH₄) are often written as planar molecules:

As shown in Figure 2.3 "The Three-Dimensional Structures of Water, Ammonia, and Methane", however, the actual three-dimensional structure of NH₃ looks like a pyramid with a triangular base of three hydrogen atoms. The structure of CH₄, with four hydrogen atoms arranged around a central carbon atom as shown in Figure 2.3 "The Three-Dimensional Structures of Water, Ammonia, and Methane", is tetrahedral. That is, the hydrogen atoms are positioned at every other vertex of a cube. Many compounds—carbon compounds, in particular—have four bonded atoms arranged around a central atom to form a tetrahedron.

Figure 2.3 The Three-Dimensional Structures of Water, Ammonia, and Methane

(a) Water is a V-shaped molecule, in which all three atoms lie in a plane. (b) In contrast, ammonia has a pyramidal structure, in which the three hydrogen atoms form the base of the pyramid and the nitrogen atom is at the vertex. (c) The four hydrogen atoms of methane form a tetrahedron; the carbon atom lies in the center.

CH₄. Methane has a three-dimensional, tetrahedral structure.

Figure 2.1 "Elements That Exist as Covalent Molecules", Figure 2.2 "Molecules That Contain Single, Double, and Triple Bonds", and Figure 2.3 "The Three-Dimensional Structures of Water, Ammonia, and Methane" illustrate different ways to represent the structures of molecules. It should be clear that there is no single “best” way to draw the structure of a molecule; the method you use depends on which aspect of the structure you want to emphasize and how much time and effort you want to spend. Figure 2.4 "Different Ways of Representing the Structure of a Molecule" shows some of the different ways to portray the structure of a slightly more complex molecule: methanol. These representations differ greatly in their information content. For example, the molecular formula for methanol (part (a) in Figure 2.4 "Different Ways of Representing the Structure of a Molecule") gives only the number of each kind of atom; writing methanol as CH₄O tells nothing about its structure. In contrast, the structural formula (part (b) in Figure 2.4 "Different Ways of Representing the Structure of a Molecule") indicates how the atoms are connected, but it makes methanol look as if it is planar (which it is not). Both the ball-and-stick model (part (c) in Figure 2.4 "Different Ways of Representing the Structure of a Molecule") and the perspective drawing (part (d) in Figure 2.4 "Different Ways of Representing the Structure of a Molecule") show the three-dimensional structure of the molecule. The latter (also called a wedge-and-dash representation) is the easiest way to sketch the structure of a molecule in three dimensions. It shows which atoms are above and below the plane of the paper by using wedges and dashes, respectively; the central atom is always assumed to be in the plane of the paper. The space-filling model (part (e) in Figure 2.4 "Different Ways of Representing the Structure of a Molecule") illustrates the approximate relative sizes of the atoms in the molecule, but it does not show the bonds between the atoms. Also, in a space-filling model, atoms at the “front” of the molecule may obscure atoms at the “back.”

Figure 2.4 Different Ways of Representing the Structure of a Molecule

(a) The molecular formula for methanol gives only the number of each kind of atom present. (b) The structural formula shows which atoms are connected. (c) The ball-and-stick model shows the atoms as spheres and the bonds as sticks. (d) A perspective drawing (also called a wedge-and-dash representation) attempts to show the three-dimensional structure of the molecule. (e) The space-filling model shows the atoms in the molecule but not the bonds. (f) The condensed structural formula is by far the easiest and most common way to represent a molecule.

Although a structural formula, a ball-and-stick model, a perspective drawing, and a space-filling model provide a significant amount of information about the structure of a molecule, each requires time and effort. Consequently, chemists often use a condensed structural formula (part (f) in Figure 2.4 "Different Ways of Representing the Structure of a Molecule"), which omits the lines representing bonds between atoms and simply lists the atoms bonded to a given atom next to it. Multiple groups attached to the same atom are shown in parentheses, followed by a subscript that indicates the number of such groups. For example, the condensed structural formula for methanol is CH₃OH, which tells us that the molecule contains a CH₃ unit that looks like a fragment of methane (CH₄). Methanol can therefore be viewed either as a methane molecule in which one hydrogen atom has been replaced by an –OH group or as a water molecule in which one hydrogen atom has been replaced by a –CH₃ fragment. Because of their ease of use and information content, we use condensed structural formulas for molecules throughout this text. Ball-and-stick models are used when needed to illustrate the three-dimensional structure of molecules, and space-filling models are used only when it is necessary to visualize the relative sizes of atoms or molecules to understand an important point.

Example 2

Write the molecular formula for each compound. The condensed structural formula is given.

1. Sulfur monochloride (also called disulfur dichloride) is a vile-smelling, corrosive yellow liquid used in the production of synthetic rubber. Its condensed structural formula is ClSSCl.
2. Ethylene glycol is the major ingredient in antifreeze. Its condensed structural formula is HOCH₂CH₂OH.
3. Trimethylamine is one of the substances responsible for the smell of spoiled fish. Its condensed structural formula is (CH₃)₃N.

Given: condensed structural formula

Asked for: molecular formula

Strategy:

A Identify every element in the condensed structural formula and then determine whether the compound is organic or inorganic.

B As appropriate, use either organic or inorganic convention to list the elements. Then add appropriate subscripts to indicate the number of atoms of each element present in the molecular formula.

Solution:

The molecular formula lists the elements in the molecule and the number of atoms of each.

1. A Each molecule of sulfur monochloride has two sulfur atoms and two chlorine atoms. Because it does not contain mostly carbon and hydrogen, it is an inorganic compound. B Sulfur lies to the left of chlorine in the periodic table, so it is written first in the formula. Adding subscripts gives the molecular formula S₂Cl₂.
2. A Counting the atoms in ethylene glycol, we get six hydrogen atoms, two carbon atoms, and two oxygen atoms per molecule. The compound consists mostly of carbon and hydrogen atoms, so it is organic. B As with all organic compounds, C and H are written first in the molecular formula. Adding appropriate subscripts gives the molecular formula C₂H₆O₂.
3. A The condensed structural formula shows that trimethylamine contains three CH₃ units, so we have one nitrogen atom, three carbon atoms, and nine hydrogen atoms per molecule. Because trimethylamine contains mostly carbon and hydrogen, it is an organic compound. B According to the convention for organic compounds, C and H are written first, giving the molecular formula C₃H₉N.

Exercise

Write the molecular formula for each molecule.

1. Chloroform, which was one of the first anesthetics and was used in many cough syrups until recently, contains one carbon atom, one hydrogen atom, and three chlorine atoms. Its condensed structural formula is CHCl₃.
2. Hydrazine is used as a propellant in the attitude jets of the space shuttle. Its condensed structural formula is H₂NNH₂.
3. Putrescine is a pungent-smelling compound first isolated from extracts of rotting meat. Its condensed structural formula is H₂NCH₂CH₂CH₂CH₂NH₂. This is often written as H₂N(CH₂)₄NH₂ to indicate that there are four CH₂ fragments linked together.

Answer:

1. CHCl₃
2. N₂H₄
3. C₄H₁₂N₂

Ionic Compounds

The substances described in the preceding discussion are composed of molecules that are electrically neutral; that is, the number of positively charged protons in the nucleus is equal to the number of negatively charged electrons. In contrast, ions are atoms or assemblies of atoms that have a net electrical charge. Ions that contain fewer electrons than protons have a net positive charge and are called cationsAn ion that has fewer electrons than protons, resulting in a net positive charge.. Conversely, ions that contain more electrons than protons have a net negative charge and are called anionsAn ion that has fewer protons than electrons, resulting in a net negative charge.. Ionic compounds contain both cations and anions in a ratio that results in no net electrical charge.

In covalent compounds, electrons are shared between bonded atoms and are simultaneously attracted to more than one nucleus. In contrast, ionic compounds contain cations and anions rather than discrete neutral molecules. Ionic compounds are held together by the attractive electrostatic interactions between cations and anions. In an ionic compound, the cations and anions are arranged in space to form an extended three-dimensional array that maximizes the number of attractive electrostatic interactions and minimizes the number of repulsive electrostatic interactions (Figure 2.5 "Covalent and Ionic Bonding"). As shown in Equation 2.1, the electrostatic energy of the interaction between two charged particles is proportional to the product of the charges on the particles and inversely proportional to the distance between them:

where Q₁ and Q₂ are the electrical charges on particles 1 and 2, and r is the distance between them. When Q₁ and Q₂ are both positive, corresponding to the charges on cations, the cations repel each other and the electrostatic energy is positive. When Q₁ and Q₂ are both negative, corresponding to the charges on anions, the anions repel each other and the electrostatic energy is again positive. The electrostatic energy is negative only when the charges have opposite signs; that is, positively charged species are attracted to negatively charged species and vice versa. As shown in Figure 2.6 "The Effect of Charge and Distance on the Strength of Electrostatic Interactions", the strength of the interaction is proportional to the magnitude of the charges and decreases as the distance between the particles increases. We will return to these energetic factors in Chapter 8 "Ionic versus Covalent Bonding", where they are described in greater quantitative detail.

Figure 2.5 Covalent and Ionic Bonding

(a) In molecular hydrogen (H₂), two hydrogen atoms share two electrons to form a covalent bond. (b) The ionic compound NaCl forms when electrons from sodium atoms are transferred to chlorine atoms. The resulting Na⁺ and Cl⁻ ions form a three-dimensional solid that is held together by attractive electrostatic interactions.

Figure 2.6 The Effect of Charge and Distance on the Strength of Electrostatic Interactions

As the charge on ions increases or the distance between ions decreases, so does the strength of the attractive (−…+) or repulsive (−…− or +…+) interactions. The strength of these interactions is represented by the thickness of the arrows.

One example of an ionic compound is sodium chloride (NaCl; Figure 2.7 "Sodium Chloride: an Ionic Solid"), formed from sodium and chlorine. In forming chemical compounds, many elements have a tendency to gain or lose enough electrons to attain the same number of electrons as the noble gas closest to them in the periodic table. When sodium and chlorine come into contact, each sodium atom gives up an electron to become a Na⁺ ion, with 11 protons in its nucleus but only 10 electrons (like neon), and each chlorine atom gains an electron to become a Cl⁻ ion, with 17 protons in its nucleus and 18 electrons (like argon), as shown in part (b) in Figure 2.5 "Covalent and Ionic Bonding". Solid sodium chloride contains equal numbers of cations (Na⁺) and anions (Cl⁻), thus maintaining electrical neutrality. Each Na⁺ ion is surrounded by 6 Cl⁻ ions, and each Cl⁻ ion is surrounded by 6 Na⁺ ions. Because of the large number of attractive Na⁺Cl⁻ interactions, the total attractive electrostatic energy in NaCl is great.

Figure 2.7 Sodium Chloride: an Ionic Solid

The planes of an NaCl crystal reflect the regular three-dimensional arrangement of its Na⁺ (purple) and Cl⁻ (green) ions.

Consistent with a tendency to have the same number of electrons as the nearest noble gas, when forming ions, elements in groups 1, 2, and 3 tend to lose one, two, and three electrons, respectively, to form cations, such as Na⁺ and Mg²⁺. They then have the same number of electrons as the nearest noble gas: neon. Similarly, K⁺, Ca²⁺, and Sc³⁺ have 18 electrons each, like the nearest noble gas: argon. In addition, the elements in group 13 lose three electrons to form cations, such as Al³⁺, again attaining the same number of electrons as the noble gas closest to them in the periodic table. Because the lanthanides and actinides formally belong to group 3, the most common ion formed by these elements is M³⁺, where M represents the metal. Conversely, elements in groups 17, 16, and 15 often react to gain one, two, and three electrons, respectively, to form ions such as Cl⁻, S²⁻, and P³⁻. Ions such as these, which contain only a single atom, are called monatomic ionsAn ion with only a single atom.. You can predict the charges of most monatomic ions derived from the main group elements by simply looking at the periodic table and counting how many columns an element lies from the extreme left or right. For example, you can predict that barium (in group 2) will form Ba²⁺ to have the same number of electrons as its nearest noble gas, xenon, that oxygen (in group 16) will form O²⁻ to have the same number of electrons as neon, and cesium (in group 1) will form Cs⁺ to also have the same number of electrons as xenon. Note that this method does not usually work for most of the transition metals, as you will learn in Section 2.3 "Naming Ionic Compounds". Some common monatomic ions are in Table 2.2 "Some Common Monatomic Ions and Their Names".

Physical Properties of Ionic and Covalent Compounds

In general, ionic and covalent compounds have different physical properties. Ionic compounds usually form hard crystalline solids that melt at rather high temperatures and are very resistant to evaporation. These properties stem from the characteristic internal structure of an ionic solid, illustrated schematically in part (a) in Figure 2.8 "Interactions in Ionic and Covalent Solids", which shows the three-dimensional array of alternating positive and negative ions held together by strong electrostatic attractions. In contrast, as shown in part (b) in Figure 2.8 "Interactions in Ionic and Covalent Solids", most covalent compounds consist of discrete molecules held together by comparatively weak intermolecular forces (the forces between molecules), even though the atoms within each molecule are held together by strong intramolecular covalent bonds (the forces within the molecule). Covalent substances can be gases, liquids, or solids at room temperature and pressure, depending on the strength of the intermolecular interactions. Covalent molecular solids tend to form soft crystals that melt at rather low temperatures and evaporate relatively easily.Some covalent substances, however, are not molecular but consist of infinite three-dimensional arrays of covalently bonded atoms and include some of the hardest materials known, such as diamond. This topic will be addressed in Chapter 12 "Solids". The covalent bonds that hold the atoms together in the molecules are unaffected when covalent substances melt or evaporate, so a liquid or vapor of discrete, independent molecules is formed. For example, at room temperature, methane, the major constituent of natural gas, is a gas that is composed of discrete CH₄ molecules. A comparison of the different physical properties of ionic compounds and covalent molecular substances is given in Table 2.3 "The Physical Properties of Typical Ionic Compounds and Covalent Molecular Substances".

Figure 2.8 Interactions in Ionic and Covalent Solids

(a) The positively and negatively charged ions in an ionic solid such as sodium chloride (NaCl) are held together by strong electrostatic interactions. (b) In this representation of the packing of methane (CH₄) molecules in solid methane, a prototypical molecular solid, the methane molecules are held together in the solid only by relatively weak intermolecular forces, even though the atoms within each methane molecule are held together by strong covalent bonds.

Binary Ionic Compounds

An ionic compound that contains only two elements, one present as a cation and one as an anion, is called a binary ionic compoundAn ionic compound that contains only two elements, one present as a cation and one as an anion.. One example is MgCl_2, a coagulant used in the preparation of tofu from soybeans. For binary ionic compounds, the subscripts in the empirical formula can also be obtained by crossing charges: use the absolute value of the charge on one ion as the subscript for the other ion. This method is shown schematically as follows:

Crossing charges. One method for obtaining subscripts in the empirical formula is by crossing charges.

When crossing charges, you will sometimes find it necessary to reduce the subscripts to their simplest ratio to write the empirical formula. Consider, for example, the compound formed by Mg²⁺ and O²⁻. Using the absolute values of the charges on the ions as subscripts gives the formula Mg₂O₂:

This simplifies to its correct empirical formula MgO. The empirical formula has one Mg²⁺ ion and one O²⁻ ion.

Polyatomic Ions

Polyatomic ionsA group of two or more atoms that has a net electrical charge. are groups of atoms that bear a net electrical charge, although the atoms in a polyatomic ion are held together by the same covalent bonds that hold atoms together in molecules. Just as there are many more kinds of molecules than simple elements, there are many more kinds of polyatomic ions than monatomic ions. Two examples of polyatomic cations are the ammonium (NH₄⁺) and the methylammonium (CH₃NH₃⁺) ions. Polyatomic anions are much more numerous than polyatomic cations; some common examples are in Table 2.4 "Common Polyatomic Ions and Their Names".

The method we used to predict the empirical formulas for ionic compounds that contain monatomic ions can also be used for compounds that contain polyatomic ions. The overall charge on the cations must balance the overall charge on the anions in the formula unit. Thus K⁺ and NO₃⁻ ions combine in a 1:1 ratio to form KNO₃ (potassium nitrate or saltpeter), a major ingredient in black gunpowder. Similarly, Ca²⁺ and SO₄²⁻ form CaSO₄ (calcium sulfate), which combines with varying amounts of water to form gypsum and plaster of Paris. The polyatomic ions NH₄⁺ and NO₃⁻ form NH₄NO₃ (ammonium nitrate), which is a widely used fertilizer and, in the wrong hands, an explosive. One example of a compound in which the ions have charges of different magnitudes is calcium phosphate, which is composed of Ca²⁺ and PO₄³⁻ ions; it is a major component of bones. The compound is electrically neutral because the ions combine in a ratio of three Ca²⁺ ions [3(+2) = +6] for every two ions [2(−3) = −6], giving an empirical formula of Ca₃(PO₄)₂; the parentheses around PO₄ in the empirical formula indicate that it is a polyatomic ion. Writing the formula for calcium phosphate as Ca₃P₂O₈ gives the correct number of each atom in the formula unit, but it obscures the fact that the compound contains readily identifiable PO₄³⁻ ions.

Hydrates

Many ionic compounds occur as hydratesA compound that contains specific ratios of loosely bound water molecules, called waters of hydration., compounds that contain specific ratios of loosely bound water molecules, called waters of hydrationThe loosely bound water molecules in hydrate compounds. These waters of hydration can often be removed by simply heating the compound.. Waters of hydration can often be removed simply by heating. For example, calcium dihydrogen phosphate can form a solid that contains one molecule of water per Ca(H₂PO₄)₂ unit and is used as a leavening agent in the food industry to cause baked goods to rise. The empirical formula for the solid is Ca(H₂PO₄)₂·H₂O. In contrast, copper sulfate usually forms a blue solid that contains five waters of hydration per formula unit, with the empirical formula CuSO₄·5H₂O. When heated, all five water molecules are lost, giving a white solid with the empirical formula CuSO₄ (Figure 2.9 "Loss of Water from a Hydrate with Heating").

Figure 2.9 Loss of Water from a Hydrate with Heating

When blue CuSO₄·5H₂O is heated, two molecules of water are lost at 30°C, two more at 110°C, and the last at 250°C to give white CuSO₄.

Compounds that differ only in the numbers of waters of hydration can have very different properties. For example, CaSO₄·½H₂O is plaster of Paris, which was often used to make sturdy casts for broken arms or legs, whereas CaSO₄·2H₂O is the less dense, flakier gypsum, a mineral used in drywall panels for home construction. When a cast would set, a mixture of plaster of Paris and water crystallized to give solid CaSO₄·2H₂O. Similar processes are used in the setting of cement and concrete.

Binary Inorganic Compounds

Binary covalent compounds—that is, covalent compounds that contain only two elements—are named using a procedure similar to that used to name simple ionic compounds, but prefixes are added as needed to indicate the number of atoms of each kind. The procedure, diagrammed in Figure 2.13 "Naming a Covalent Inorganic Compound", uses the following steps:

Figure 2.13 Naming a Covalent Inorganic Compound

1. Place the elements in their proper order.

   1. The element farthest to the left in the periodic table is usually named first. If both elements are in the same group, the element closer to the bottom of the column is named first.
   2. The second element is named as if it were a monatomic anion in an ionic compound (even though it is not), with the suffix -ide attached to the root of the element name.

2. Identify the number of each type of atom present.

   1. Prefixes derived from Greek stems are used to indicate the number of each type of atom in the formula unit (Table 2.6 "Prefixes for Indicating the Number of Atoms in Chemical Names"). The prefix mono- (“one”) is used only when absolutely necessary to avoid confusion, just as we omit the subscript 1 when writing molecular formulas.

      To demonstrate steps 1 and 2a, we name HCl as hydrogen chloride (because hydrogen is to the left of chlorine in the periodic table) and PCl₅ as phosphorus pentachloride. The order of the elements in the name of BrF₃, bromine trifluoride, is determined by the fact that bromine lies below fluorine in group 17.

      Table 2.6 Prefixes for Indicating the Number of Atoms in Chemical Names

      Prefix	Number
      mono-	1
      di-	2
      tri-	3
      tetra-	4
      penta-	5
      hexa-	6
      hepta-	7
      octa-	8
      nona-	9
      deca-	10
      undeca-	11
      dodeca-	12

   2. If a molecule contains more than one atom of both elements, then prefixes are used for both. Thus N₂O₃ is dinitrogen trioxide, as shown in Figure 2.13 "Naming a Covalent Inorganic Compound".
   3. In some names, the final a or o of the prefix is dropped to avoid awkward pronunciation. Thus OsO₄ is osmium tetroxide rather than osmium tetraoxide.

3. Write the name of the compound.

   1. Binary compounds of the elements with oxygen are generally named as “element oxide,” with prefixes that indicate the number of atoms of each element per formula unit. For example, CO is carbon monoxide. The only exception is binary compounds of oxygen with fluorine, which are named as oxygen fluorides. (The reasons for this convention will become clear in Chapter 7 "The Periodic Table and Periodic Trends" and Chapter 8 "Ionic versus Covalent Bonding".)
   2. Certain compounds are always called by the common names that were assigned long ago when names rather than formulas were used. For example, H₂O is water (not dihydrogen oxide); NH₃ is ammonia; PH₃ is phosphine; SiH₄ is silane; and B₂H₆, a dimer of BH₃, is diborane. For many compounds, the systematic name and the common name are both used frequently, so you must be familiar with them. For example, the systematic name for NO is nitrogen monoxide, but it is much more commonly called nitric oxide. Similarly, N₂O is usually called nitrous oxide rather than dinitrogen monoxide. Notice that the suffixes -ic and -ous are the same ones used for ionic compounds.

The structures of some of the compounds in Example 8 and Example 9 are shown in Figure 2.14 "The Structures of Some Covalent Inorganic Compounds and the Locations of the “Central Atoms” in the Periodic Table", along with the location of the “central atom” of each compound in the periodic table. It may seem that the compositions and structures of such compounds are entirely random, but this is not true. After you have mastered the material in Chapter 7 "The Periodic Table and Periodic Trends" and Chapter 8 "Ionic versus Covalent Bonding", you will be able to predict the compositions and structures of compounds of this type with a high degree of accuracy.

Figure 2.14 The Structures of Some Covalent Inorganic Compounds and the Locations of the “Central Atoms” in the Periodic Table

The compositions and structures of covalent inorganic compounds are not random. As you will learn in Chapter 7 "The Periodic Table and Periodic Trends" and Chapter 8 "Ionic versus Covalent Bonding", they can be predicted from the locations of the component atoms in the periodic table.

Hydrocarbons

Approximately one-third of the compounds produced industrially are organic compounds. All living organisms are composed of organic compounds, as is most of the food you consume, the medicines you take, the fibers in the clothes you wear, and the plastics in the materials you use. Section 2.1 "Chemical Compounds" introduced two organic compounds: methane (CH₄) and methanol (CH₃OH). These and other organic compounds appear frequently in discussions and examples throughout this text.

The detection of organic compounds is useful in many fields. In one recently developed application, scientists have devised a new method called “material degradomics” to make it possible to monitor the degradation of old books and historical documents. As paper ages, it produces a familiar “old book smell” from the release of organic compounds in gaseous form. The composition of the gas depends on the original type of paper used, a book’s binding, and the applied media. By analyzing these organic gases and isolating the individual components, preservationists are better able to determine the condition of an object and those books and documents most in need of immediate protection.

The simplest class of organic compounds is the hydrocarbonsThe simplest class of organic molecules, consisting of only carbon and hydrogen., which consist entirely of carbon and hydrogen. Petroleum and natural gas are complex, naturally occurring mixtures of many different hydrocarbons that furnish raw materials for the chemical industry. The four major classes of hydrocarbons are the alkanesA saturated hydrocarbon with only carbon–hydrogen and carbon–carbon single bonds., which contain only carbon–hydrogen and carbon–carbon single bonds; the alkenesAn unsaturated hydrocarbon with at least one carbon–carbon double bond., which contain at least one carbon–carbon double bond; the alkynesAn unsaturated hydrocarbon with at least one carbon–carbon triple bond., which contain at least one carbon–carbon triple bond; and the aromatic hydrocarbonsAn unsaturated hydrocarbon consisting of a ring of six carbon atoms with alternating single and double bonds., which usually contain rings of six carbon atoms that can be drawn with alternating single and double bonds. Alkanes are also called saturated hydrocarbons, whereas hydrocarbons that contain multiple bonds (alkenes, alkynes, and aromatics) are unsaturated.

Alkanes

The simplest alkane is methane (CH₄), a colorless, odorless gas that is the major component of natural gas. In larger alkanes whose carbon atoms are joined in an unbranched chain (straight-chain alkanes), each carbon atom is bonded to at most two other carbon atoms. The structures of two simple alkanes are shown in Figure 2.15 "Straight-Chain Alkanes with Two and Three Carbon Atoms", and the names and condensed structural formulas for the first 10 straight-chain alkanes are in Table 2.7 "The First 10 Straight-Chain Alkanes". The names of all alkanes end in -ane, and their boiling points increase as the number of carbon atoms increases.

Figure 2.15 Straight-Chain Alkanes with Two and Three Carbon Atoms

Table 2.7 The First 10 Straight-Chain Alkanes

Name	Number of Carbon Atoms	Molecular Formula	Condensed Structural Formula	Boiling Point (°C)	Uses
methane	1	CH4	CH4	−162	natural gas constituent
ethane	2	C2H6	CH3CH3	−89	natural gas constituent
propane	3	C3H8	CH3CH2CH3	−42	bottled gas
butane	4	C4H10	CH3CH2CH2CH3 or CH3(CH2)2CH3	0	lighters, bottled gas
pentane	5	C5H12	CH3(CH2)3CH3	36	solvent, gasoline
hexane	6	C6H14	CH3(CH2)4CH3	69	solvent, gasoline
heptane	7	C7H16	CH3(CH2)5CH3	98	solvent, gasoline
octane	8	C8H18	CH3(CH2)6CH3	126	gasoline
nonane	9	C9H20	CH3(CH2)7CH3	151	gasoline
decane	10	C10H22	CH3(CH2)8CH3	174	kerosene

Alkanes with four or more carbon atoms can have more than one arrangement of atoms. The carbon atoms can form a single unbranched chain, or the primary chain of carbon atoms can have one or more shorter chains that form branches. For example, butane (C₄H₁₀) has two possible structures. Normal butane (usually called n-butane) is CH₃CH₂CH₂CH₃, in which the carbon atoms form a single unbranched chain. In contrast, the condensed structural formula for isobutane is (CH₃)₂CHCH₃, in which the primary chain of three carbon atoms has a one-carbon chain branching at the central carbon. Three-dimensional representations of both structures are as follows:

The systematic names for branched hydrocarbons use the lowest possible number to indicate the position of the branch along the longest straight carbon chain in the structure. Thus the systematic name for isobutane is 2-methylpropane, which indicates that a methyl group (a branch consisting of –CH₃) is attached to the second carbon of a propane molecule. Similarly, you will learn in Section 2.6 "Industrially Important Chemicals" that one of the major components of gasoline is commonly called isooctane; its structure is as follows:

As you can see, the compound has a chain of five carbon atoms, so it is a derivative of pentane. There are two methyl group branches at one carbon atom and one methyl group at another. Using the lowest possible numbers for the branches gives 2,2,4-trimethylpentane for the systematic name of this compound.

Alkenes

The simplest alkenes are ethylene, C₂H₄ or CH₂=CH₂, and propylene, C₃H₆ or CH₃CH=CH₂ (part (a) in Figure 2.16 "Some Simple (a) Alkenes, (b) Alkynes, and (c) Cyclic Hydrocarbons"). The names of alkenes that have more than three carbon atoms use the same stems as the names of the alkanes (Table 2.7 "The First 10 Straight-Chain Alkanes") but end in -ene instead of -ane.

Once again, more than one structure is possible for alkenes with four or more carbon atoms. For example, an alkene with four carbon atoms has three possible structures. One is CH₂=CHCH₂CH₃ (1-butene), which has the double bond between the first and second carbon atoms in the chain. The other two structures have the double bond between the second and third carbon atoms and are forms of CH₃CH=CHCH₃ (2-butene). All four carbon atoms in 2-butene lie in the same plane, so there are two possible structures (part (a) in Figure 2.16 "Some Simple (a) Alkenes, (b) Alkynes, and (c) Cyclic Hydrocarbons"). If the two methyl groups are on the same side of the double bond, the compound is cis-2-butene (from the Latin cis, meaning “on the same side”). If the two methyl groups are on opposite sides of the double bond, the compound is trans-2-butene (from the Latin trans, meaning “across”). These are distinctly different molecules: cis-2-butene melts at −138.9°C, whereas trans-2-butene melts at −105.5°C.

Figure 2.16 Some Simple (a) Alkenes, (b) Alkynes, and (c) Cyclic Hydrocarbons

The positions of the carbon atoms in the chain are indicated by C1 or C2.

Just as a number indicates the positions of branches in an alkane, the number in the name of an alkene specifies the position of the first carbon atom of the double bond. The name is based on the lowest possible number starting from either end of the carbon chain, so CH₃CH₂CH=CH₂ is called 1-butene, not 3-butene. Note that CH₂=CHCH₂CH₃ and CH₃CH₂CH=CH₂ are different ways of writing the same molecule (1-butene) in two different orientations.

The name of a compound does not depend on its orientation. As illustrated for 1-butene, both condensed structural formulas and molecular models show different orientations of the same molecule. Don’t let orientation fool you; you must be able to recognize the same structure no matter what its orientation.

Note the Pattern

The positions of groups or multiple bonds are always indicated by the lowest number possible.

Cyclic Hydrocarbons

In a cyclic hydrocarbonA hydrocarbon in which the ends of the carbon chain are connected to form a ring of covalently bonded carbon atoms., the ends of a hydrocarbon chain are connected to form a ring of covalently bonded carbon atoms. Cyclic hydrocarbons are named by attaching the prefix cyclo- to the name of the alkane, the alkene, or the alkyne. The simplest cyclic alkanes are cyclopropane (C₃H₆) a flammable gas that is also a powerful anesthetic, and cyclobutane (C₄H₈) (part (c) in Figure 2.16 "Some Simple (a) Alkenes, (b) Alkynes, and (c) Cyclic Hydrocarbons"). The most common way to draw the structures of cyclic alkanes is to sketch a polygon with the same number of vertices as there are carbon atoms in the ring; each vertex represents a CH₂ unit. The structures of the cycloalkanes that contain three to six carbon atoms are shown schematically in Figure 2.17 "The Simple Cycloalkanes".

Figure 2.17 The Simple Cycloalkanes

Aromatic Hydrocarbons

Alkanes, alkenes, alkynes, and cyclic hydrocarbons are generally called aliphatic hydrocarbonsAlkanes, alkenes, alkynes, and cyclic hydrocarbons (hydrocarbons that are not aromatic).. The name comes from the Greek aleiphar, meaning “oil,” because the first examples were extracted from animal fats. In contrast, the first examples of aromatic hydrocarbons, also called arenes, were obtained by the distillation and degradation of highly scented (thus aromatic) resins from tropical trees.

The simplest aromatic hydrocarbon is benzene (C₆H₆), which was first obtained from a coal distillate. The word aromatic now refers to benzene and structurally similar compounds. As shown in part (a) in Figure 2.18 "Two Aromatic Hydrocarbons: (a) Benzene and (b) Toluene", it is possible to draw the structure of benzene in two different but equivalent ways, depending on which carbon atoms are connected by double bonds or single bonds. Toluene is similar to benzene, except that one hydrogen atom is replaced by a –CH₃ group; it has the formula C₇H₈ (part (b) in Figure 2.18 "Two Aromatic Hydrocarbons: (a) Benzene and (b) Toluene"). As you will soon learn, the chemical behavior of aromatic compounds differs from the behavior of aliphatic compounds. Benzene and toluene are found in gasoline, and benzene is the starting material for preparing substances as diverse as aspirin and nylon.

Figure 2.18 Two Aromatic Hydrocarbons: (a) Benzene and (b) Toluene

Figure 2.19 "Two Hydrocarbons with the Molecular Formula C" illustrates two of the molecular structures possible for hydrocarbons that have six carbon atoms. As you can see, compounds with the same molecular formula can have very different structures.

Figure 2.19 Two Hydrocarbons with the Molecular Formula C₆H₁₂

Example 10

Write the condensed structural formula for each hydrocarbon.

1. n-heptane
2. 2-pentene
3. 2-butyne
4. cyclooctene

Given: name of hydrocarbon

Asked for: condensed structural formula

Strategy:

A Use the prefix to determine the number of carbon atoms in the molecule and whether it is cyclic. From the suffix, determine whether multiple bonds are present.

B Identify the position of any multiple bonds from the number(s) in the name and then write the condensed structural formula.

Solution:

1. A The prefix hept- tells us that this hydrocarbon has seven carbon atoms, and n- indicates that the carbon atoms form a straight chain. The suffix -ane tells that it is an alkane, with no carbon–carbon double or triple bonds. B The condensed structural formula is CH₃CH₂CH₂CH₂CH₂CH₂CH₃, which can also be written as CH₃(CH₂)₅CH₃.

2. A The prefix pent- tells us that this hydrocarbon has five carbon atoms, and the suffix -ene indicates that it is an alkene, with a carbon–carbon double bond. B The 2- tells us that the double bond begins on the second carbon of the five-carbon atom chain. The condensed structural formula of the compound is therefore CH₃CH=CHCH₂CH₃.

   

3. A The prefix but- tells us that the compound has a chain of four carbon atoms, and the suffix -yne indicates that it has a carbon–carbon triple bond. B The 2- tells us that the triple bond begins on the second carbon of the four-carbon atom chain. So the condensed structural formula for the compound is CH₃C≡CCH₃.

   

4. A The prefix cyclo- tells us that this hydrocarbon has a ring structure, and oct- indicates that it contains eight carbon atoms, which we can draw as

   

   The suffix -ene tells us that the compound contains a carbon–carbon double bond, but where in the ring do we place the double bond? B Because all eight carbon atoms are identical, it doesn’t matter. We can draw the structure of cyclooctene as

   

Exercise

Write the condensed structural formula for each hydrocarbon.

1. n-octane
2. 2-hexene
3. 1-heptyne
4. cyclopentane

Answer:

1. CH₃(CH₂)₆CH₃
2. CH₃CH=CHCH₂CH₂CH₃
3. HC≡C(CH₂)₄CH₃

4.  

   

The general name for a group of atoms derived from an alkane is an alkyl group. The name of an alkyl group is derived from the name of the alkane by adding the suffix -yl. Thus the –CH₃ fragment is a methyl group, the –CH₂CH₃ fragment is an ethyl group, and so forth, where the dash represents a single bond to some other atom or group. Similarly, groups of atoms derived from aromatic hydrocarbons are aryl groups, which sometimes have unexpected names. For example, the –C₆H₅ fragment is derived from benzene, but it is called a phenyl group. In general formulas and structures, alkyl and aryl groups are often abbreviated as RThe abbreviation used for alkyl groups and aryl groups in general formulas and structures..

Structures of alkyl and aryl groups. The methyl group is an example of an alkyl group, and the phenyl group is an example of an aryl group.

Alcohols

Replacing one or more hydrogen atoms of a hydrocarbon with an –OH group gives an alcoholA class of organic compounds obtained by replacing one or more of the hydrogen atoms of a hydrocarbon with an −OH group., represented as ROH. The simplest alcohol (CH₃OH) is called either methanol (its systematic name) or methyl alcohol (its common name) (see Figure 2.4 "Different Ways of Representing the Structure of a Molecule"). Methanol is the antifreeze in automobile windshield washer fluids, and it is also used as an efficient fuel for racing cars, most notably in the Indianapolis 500. Ethanol (or ethyl alcohol, CH₃CH₂OH) is familiar as the alcohol in fermented or distilled beverages, such as beer, wine, and whiskey; it is also used as a gasoline additive (Section 2.6 "Industrially Important Chemicals"). The simplest alcohol derived from an aromatic hydrocarbon is C₆H₅OH, phenol (shortened from phenyl alcohol), a potent disinfectant used in some sore throat medications and mouthwashes.

Ethanol, which is easy to obtain from fermentation processes, has successfully been used as an alternative fuel for several decades. Although it is a “green” fuel when derived from plants, it is an imperfect substitute for fossil fuels because it is less efficient than gasoline. Moreover, because ethanol absorbs water from the atmosphere, it can corrode an engine’s seals. Thus other types of processes are being developed that use bacteria to create more complex alcohols, such as octanol, that are more energy efficient and that have a lower tendency to absorb water. As scientists attempt to reduce mankind’s dependence on fossil fuels, the development of these so-called biofuels is a particularly active area of research.

Summary

Covalent inorganic compounds are named by a procedure similar to that used for ionic compounds, using prefixes to indicate the numbers of atoms in the molecular formula. The simplest organic compounds are the hydrocarbons, which contain only carbon and hydrogen. Alkanes contain only carbon–hydrogen and carbon–carbon single bonds, alkenes contain at least one carbon–carbon double bond, and alkynes contain one or more carbon–carbon triple bonds. Hydrocarbons can also be cyclic, with the ends of the chain connected to form a ring. Collectively, alkanes, alkenes, and alkynes are called aliphatic hydrocarbons. Aromatic hydrocarbons, or arenes, are another important class of hydrocarbons that contain rings of carbon atoms related to the structure of benzene (C₆H₆). A derivative of an alkane or an arene from which one hydrogen atom has been removed is called an alkyl group or an aryl group, respectively. Alcohols are another common class of organic compound, which contain an –OH group covalently bonded to either an alkyl group or an aryl group (often abbreviated R).

Key Takeaway

• Covalent inorganic compounds are named using a procedure similar to that used for ionic compounds, whereas hydrocarbons use a system based on the number of bonds between carbon atoms.

Conceptual Problems

1. Benzene (C₆H₆) is an organic compound, and KCl is an ionic compound. The sum of the masses of the atoms in each empirical formula is approximately the same. How would you expect the two to compare with regard to each of the following? What species are present in benzene vapor?

   1. melting point
   2. type of bonding
   3. rate of evaporation
   4. structure

2. Can an inorganic compound be classified as a hydrocarbon? Why or why not?

3. Is the compound NaHCO₃ a hydrocarbon? Why or why not?

4. Name each compound.

   1. NiO
   2. TiO₂
   3. N₂O
   4. CS₂
   5. SO₃
   6. NF₃
   7. SF₆

5. Name each compound.

   1. HgCl₂
   2. IF₅
   3. N₂O₅
   4. Cl₂O
   5. HgS
   6. PCl₅

6. For each structural formula, write the condensed formula and the name of the compound.

   1.  

      

   2.  

      

   3.  

      

   4.  

      

   5.  

      

7. For each structural formula, write the condensed formula and the name of the compound.

   1.  

      

   2.  

      

   3.  

      

   4.  

      

8. Would you expect PCl₃ to be an ionic compound or a covalent compound? Explain your reasoning.

9. What distinguishes an aromatic hydrocarbon from an aliphatic hydrocarbon?

10. The following general formulas represent specific classes of hydrocarbons. Refer to Table 2.7 "The First 10 Straight-Chain Alkanes" and Table 2.8 "Some Common Acids That Do Not Contain Oxygen" and Figure 2.16 "Some Simple (a) Alkenes, (b) Alkynes, and (c) Cyclic Hydrocarbons" and identify the classes.

    1. C_nH_{2n + 2}
    2. C_nH_2n
    3. C_nH_{2n − 2}

11. Using R to represent an alkyl or aryl group, show the general structure of an

    1. alcohol.
    2. phenol.

Answer

11. 1. ROH (where R is an alkyl group)
    2. ROH (where R is an aryl group)

Numerical Problems

1. Write the formula for each compound.

   1. dinitrogen monoxide
   2. silicon tetrafluoride
   3. boron trichloride
   4. nitrogen trifluoride
   5. phosphorus tribromide

2. Write the formula for each compound.

   1. dinitrogen trioxide
   2. iodine pentafluoride
   3. boron tribromide
   4. oxygen difluoride
   5. arsenic trichloride

3. Write the formula for each compound.

   1. thallium(I) selenide
   2. neptunium(IV) oxide
   3. iron(II) sulfide
   4. copper(I) cyanide
   5. nitrogen trichloride

4. Name each compound.

   1. RuO₄
   2. PbO₂
   3. MoF₆
   4. Hg₂(NO₃)₂·2H₂O
   5. WCl₄

5. Name each compound.

   1. NbO₂
   2. MoS₂
   3. P₄S₁₀
   4. Cu₂O
   5. ReF₅

6. Draw the structure of each compound.

   1. propyne
   2. ethanol
   3. n-hexane
   4. cyclopropane
   5. benzene

7. Draw the structure of each compound.

   1. 1-butene
   2. 2-pentyne
   3. cycloheptane
   4. toluene
   5. phenol

Answers

1. 1. N₂O
   2. SiF₄
   3. BCl₃
   4. NF₃
   5. PBr₃

3. 1. Tl₂Se
   2. NpO₂
   3. FeS
   4. CuCN
   5. NCl₃

5. 1. niobium (IV) oxide
   2. molybdenum (IV) sulfide
   3. tetraphosphorus decasulfide
   4. copper(I) oxide
   5. rhenium(V) fluoride

7. 1.  

      

   2.  

      

   3.  

      

   4.  

      

   5.  

      

Acids

The names of acids differentiate between (1) acids in which the H⁺ ion is attached to an oxygen atom of a polyatomic anion (these are called oxoacidsAn acid in which the dissociable ${\text{H}}^{+}$ ion is attached to an oxygen atom of a polyatomic anion., or occasionally oxyacids) and (2) acids in which the H⁺ ion is attached to some other element. In the latter case, the name of the acid begins with hydro- and ends in -ic, with the root of the name of the other element or ion in between. Recall that the name of the anion derived from this kind of acid always ends in -ide. Thus hydrogen chloride (HCl) gas dissolves in water to form hydrochloric acid (which contains H⁺ and Cl⁻ ions), hydrogen cyanide (HCN) gas forms hydrocyanic acid (which contains H⁺ and CN⁻ ions), and so forth (Table 2.8 "Some Common Acids That Do Not Contain Oxygen"). Examples of this kind of acid are commonly encountered and very important. For instance, your stomach contains a dilute solution of hydrochloric acid to help digest food. When the mechanisms that prevent the stomach from digesting itself malfunction, the acid destroys the lining of the stomach and an ulcer forms.

If an acid contains one or more H⁺ ions attached to oxygen, it is a derivative of one of the common oxoanions, such as sulfate (SO₄²⁻) or nitrate (NO₃⁻). These acids contain as many H⁺ ions as are necessary to balance the negative charge on the anion, resulting in a neutral species such as H₂SO₄ and HNO₃.

The names of acids are derived from the names of anions according to the following rules:

1. If the name of the anion ends in -ate, then the name of the acid ends in -ic. For example, because NO₃⁻ is the nitrate ion, HNO₃ is nitric acid. Similarly, ClO₄⁻ is the perchlorate ion, so HClO₄ is perchloric acid. Two important acids are sulfuric acid (H₂SO₄) from the sulfate ion (SO₄²⁻) and phosphoric acid (H₃PO₄) from the phosphate ion (PO₄³⁻). These two names use a slight variant of the root of the anion name: sulfate becomes sulfuric and phosphate becomes phosphoric.
2. If the name of the anion ends in -ite, then the name of the acid ends in -ous. For example, OCl⁻ is the hypochlorite ion, and HOCl is hypochlorous acid; NO₂⁻ is the nitrite ion, and HNO₂ is nitrous acid; and SO₃²⁻ is the sulfite ion, and H₂SO₃ is sulfurous acid. The same roots are used whether the acid name ends in -ic or -ous; thus, sulfite becomes sulfurous.

The relationship between the names of the oxoacids and the parent oxoanions is illustrated in Figure 2.20 "The Relationship between the Names of the Oxoacids and the Names of the Parent Oxoanions", and some common oxoacids are in Table 2.9 "Some Common Oxoacids".

Figure 2.20 The Relationship between the Names of the Oxoacids and the Names of the Parent Oxoanions

Many organic compounds contain the carbonyl groupA carbon atom double-bonded to an oxygen atom. It is a characteristic feature of many organic compounds, including carboxylic acids., in which there is a carbon–oxygen double bond. In carboxylic acidsAn organic compound that contains an −OH group covalently bonded to the carbon atom of a carbonyl group. The general formula of a carboxylic acid is ${\text{RCO}}_{\text{2}}\text{H}$. In water a carboxylic acid dissociates to produce an acidic solution., an –OH group is covalently bonded to the carbon atom of the carbonyl group. Their general formula is RCO₂H, sometimes written as RCOOH:

where R can be an alkyl group, an aryl group, or a hydrogen atom. The simplest example, HCO₂H, is formic acid, so called because it is found in the secretions of stinging ants (from the Latin formica, meaning “ant”). Another example is acetic acid (CH₃CO₂H), which is found in vinegar. Like many acids, carboxylic acids tend to have sharp odors. For example, butyric acid (CH₃CH₂CH₂CO₂H), is responsible for the smell of rancid butter, and the characteristic odor of sour milk and vomit is due to lactic acid [CH₃CH(OH)CO₂H]. Some common carboxylic acids are shown in Figure 2.21 "Some Common Carboxylic Acids".

Figure 2.21 Some Common Carboxylic Acids

Although carboxylic acids are covalent compounds, when they dissolve in water, they dissociate to produce H⁺ ions (just like any other acid) and RCO₂⁻ ions. Note that only the hydrogen attached to the oxygen atom of the CO₂ group dissociates to form an H⁺ ion. In contrast, the hydrogen atom attached to the oxygen atom of an alcohol does not dissociate to form an H⁺ ion when an alcohol is dissolved in water. The reasons for the difference in behavior between carboxylic acids and alcohols will be discussed in Chapter 8 "Ionic versus Covalent Bonding".

Bases

We will present more comprehensive definitions of bases in later chapters, but virtually every base you encounter in the meantime will be an ionic compound, such as sodium hydroxide (NaOH) and barium hydroxide [Ba(OH)₂], that contain the hydroxide ion and a metal cation. These have the general formula M(OH)_n. It is important to recognize that alcohols, with the general formula ROH, are covalent compounds, not ionic compounds; consequently, they do not dissociate in water to form a basic solution (containing OH⁻ ions). When a base reacts with any of the acids we have discussed, it accepts a proton (H⁺). For example, the hydroxide ion (OH⁻) accepts a proton to form H₂O. Thus bases are also referred to as proton acceptors.

Concentrated aqueous solutions of ammonia (NH₃) contain significant amounts of the hydroxide ion, even though the dissolved substance is not primarily ammonium hydroxide (NH₄OH) as is often stated on the label. Thus aqueous ammonia solution is also a common base. Replacing a hydrogen atom of NH₃ with an alkyl group results in an amineAn organic compound that has the general formula ${\text{RNH}}_{\text{2}}$, where R is an alkyl group. Amines, like ammonia, are bases. (RNH₂), which is also a base. Amines have pungent odors—for example, methylamine (CH₃NH₂) is one of the compounds responsible for the foul odor associated with spoiled fish. The physiological importance of amines is suggested in the word vitamin, which is derived from the phrase vital amines. The word was coined to describe dietary substances that were effective at preventing scurvy, rickets, and other diseases because these substances were assumed to be amines. Subsequently, some vitamins have indeed been confirmed to be amines.

Petroleum

The petroleum that is pumped out of the ground at locations around the world is a complex mixture of several thousand organic compounds, including straight-chain alkanes, cycloalkanes, alkenes, and aromatic hydrocarbons with four to several hundred carbon atoms. The identities and relative abundances of the components vary depending on the source. So Texas crude oil is somewhat different from Saudi Arabian crude oil. In fact, the analysis of petroleum from different deposits can produce a “fingerprint” of each, which is useful in tracking down the sources of spilled crude oil. For example, Texas crude oil is “sweet,” meaning that it contains a small amount of sulfur-containing molecules, whereas Saudi Arabian crude oil is “sour,” meaning that it contains a relatively large amount of sulfur-containing molecules.

Gasoline

Petroleum is converted to useful products such as gasoline in three steps: distillation, cracking, and reforming. Recall from Chapter 1 "Introduction to Chemistry" that distillation separates compounds on the basis of their relative volatility, which is usually inversely proportional to their boiling points. Part (a) in Figure 2.23 "The Distillation of Petroleum" shows a cutaway drawing of a column used in the petroleum industry for separating the components of crude oil. The petroleum is heated to approximately 400°C (750°F), at which temperature it has become a mixture of liquid and vapor. This mixture, called the feedstock, is introduced into the refining tower. The most volatile components (those with the lowest boiling points) condense at the top of the column where it is cooler, while the less volatile components condense nearer the bottom. Some materials are so nonvolatile that they collect at the bottom without evaporating at all. Thus the composition of the liquid condensing at each level is different. These different fractions, each of which usually consists of a mixture of compounds with similar numbers of carbon atoms, are drawn off separately. Part (b) in Figure 2.23 "The Distillation of Petroleum" shows the typical fractions collected at refineries, the number of carbon atoms they contain, their boiling points, and their ultimate uses. These products range from gases used in natural and bottled gas to liquids used in fuels and lubricants to gummy solids used as tar on roads and roofs.

Figure 2.23 The Distillation of Petroleum

(a) This is a diagram of a distillation column used for separating petroleum fractions. (b) Petroleum fractions condense at different temperatures, depending on the number of carbon atoms in the molecules, and are drawn off from the column. The most volatile components (those with the lowest boiling points) condense at the top of the column, and the least volatile (those with the highest boiling points) condense at the bottom.

The economics of petroleum refining are complex. For example, the market demand for kerosene and lubricants is much lower than the demand for gasoline, yet all three fractions are obtained from the distillation column in comparable amounts. Furthermore, most gasolines and jet fuels are blends with very carefully controlled compositions that cannot vary as their original feedstocks did. To make petroleum refining more profitable, the less volatile, lower-value fractions must be converted to more volatile, higher-value mixtures that have carefully controlled formulas. The first process used to accomplish this transformation is crackingA process in petroleum refining in which the larger and heavier hydrocarbons in kerosene and higher-boiling-point fractions are heated to high temperatures, causing the carbon–carbon bonds to break (“crack”), thus producing a more volatile mixture., in which the larger and heavier hydrocarbons in the kerosene and higher-boiling-point fractions are heated to temperatures as high as 900°C. High-temperature reactions cause the carbon–carbon bonds to break, which converts the compounds to lighter molecules similar to those in the gasoline fraction. Thus in cracking, a straight-chain alkane with a number of carbon atoms corresponding to the kerosene fraction is converted to a mixture of hydrocarbons with a number of carbon atoms corresponding to the lighter gasoline fraction. The second process used to increase the amount of valuable products is called reformingThe second process used in petroleum refining, which is the chemical conversion of straight-chain alkanes to either branched-chain alkanes or mixtures of aromatic hydrocarbons.; it is the chemical conversion of straight-chain alkanes to either branched-chain alkanes or mixtures of aromatic hydrocarbons. Using metals such as platinum brings about the necessary chemical reactions. The mixtures of products obtained from cracking and reforming are separated by fractional distillation.

Octane Ratings

The quality of a fuel is indicated by its octane ratingA measure of a fuel’s ability to burn in a combustion engine without knocking or pinging (indications of premature combustion). The higher the octane rating, the higher quality the fuel., which is a measure of its ability to burn in a combustion engine without knocking or pinging. Knocking and pinging signal premature combustion (Figure 2.24 "The Burning of Gasoline in an Internal Combustion Engine"), which can be caused either by an engine malfunction or by a fuel that burns too fast. In either case, the gasoline-air mixture detonates at the wrong point in the engine cycle, which reduces the power output and can damage valves, pistons, bearings, and other engine components. The various gasoline formulations are designed to provide the mix of hydrocarbons least likely to cause knocking or pinging in a given type of engine performing at a particular level.

Figure 2.24 The Burning of Gasoline in an Internal Combustion Engine

(a) Normally, fuel is ignited by the spark plug, and combustion spreads uniformly outward. (b) Gasoline with an octane rating that is too low for the engine can ignite prematurely, resulting in uneven burning that causes knocking and pinging.

The octane scale was established in 1927 using a standard test engine and two pure compounds: n-heptane and isooctane (2,2,4-trimethylpentane). n-Heptane, which causes a great deal of knocking on combustion, was assigned an octane rating of 0, whereas isooctane, a very smooth-burning fuel, was assigned an octane rating of 100. Chemists assign octane ratings to different blends of gasoline by burning a sample of each in a test engine and comparing the observed knocking with the amount of knocking caused by specific mixtures of n-heptane and isooctane. For example, the octane rating of a blend of 89% isooctane and 11% n-heptane is simply the average of the octane ratings of the components weighted by the relative amounts of each in the blend. Converting percentages to decimals, we obtain the octane rating of the mixture:

0.89(100) + 0.11(0) = 89

A gasoline that performs at the same level as a blend of 89% isooctane and 11% n-heptane is assigned an octane rating of 89; this represents an intermediate grade of gasoline. Regular gasoline typically has an octane rating of 87; premium has a rating of 93 or higher.

As shown in Figure 2.25 "The Octane Ratings of Some Hydrocarbons and Common Additives", many compounds that are now available have octane ratings greater than 100, which means they are better fuels than pure isooctane. In addition, antiknock agents, also called octane enhancers, have been developed. One of the most widely used for many years was tetraethyllead [(C₂H₅)₄Pb], which at approximately 3 g/gal gives a 10–15-point increase in octane rating. Since 1975, however, lead compounds have been phased out as gasoline additives because they are highly toxic. Other enhancers, such as methyl t-butyl ether (MTBE), have been developed to take their place. They combine a high octane rating with minimal corrosion to engine and fuel system parts. Unfortunately, when gasoline containing MTBE leaks from underground storage tanks, the result has been contamination of the groundwater in some locations, resulting in limitations or outright bans on the use of MTBE in certain areas. As a result, the use of alternative octane enhancers such as ethanol, which can be obtained from renewable resources such as corn, sugar cane, and, eventually, corn stalks and grasses, is increasing.

Figure 2.25 The Octane Ratings of Some Hydrocarbons and Common Additives

Example 12

You have a crude (i.e., unprocessed or straight-run) petroleum distillate consisting of 10% n-heptane, 10% n-hexane, and 80% n-pentane by mass, with an octane rating of 52. What percentage of MTBE by mass would you need to increase the octane rating of the distillate to that of regular-grade gasoline (a rating of 87), assuming that the octane rating is directly proportional to the amounts of the compounds present? Use the information presented in Figure 2.25 "The Octane Ratings of Some Hydrocarbons and Common Additives".

Given: composition of petroleum distillate, initial octane rating, and final octane rating

Asked for: percentage of MTBE by mass in final mixture

Strategy:

A Define the unknown as the percentage of MTBE in the final mixture. Then subtract this unknown from 100% to obtain the percentage of petroleum distillate.

B Multiply the percentage of MTBE and the percentage of petroleum distillate by their respective octane ratings; add these values to obtain the overall octane rating of the new mixture.

C Solve for the unknown to obtain the percentage of MTBE needed.

Solution:

A The question asks what percentage of MTBE will give an overall octane rating of 87 when mixed with the straight-run fraction. From Figure 2.25 "The Octane Ratings of Some Hydrocarbons and Common Additives", the octane rating of MTBE is 116. Let x be the percentage of MTBE, and let 100 − x be the percentage of petroleum distillate.

B Multiplying the percentage of each component by its respective octane rating and setting the sum equal to the desired octane rating of the mixture (87) times 100 gives

$$\begin{array}{ll}\text{final octane rating of mixture }\hfill & =87(100)\hfill \\ \hfill & =52(100-x)+116x\hfill \\ \hfill & =5200-52x+116x\hfill \\ \hfill & =5200+64x\hfill \end{array}$$

C Solving the equation gives x = 55%. Thus the final mixture must contain 55% MTBE by mass.

To obtain a composition of 55% MTBE by mass, you would have to add more than an equal mass of MTBE (actually 0.55/0.45, or 1.2 times) to the straight-run fraction. This is 1.2 tons of MTBE per ton of straight-run gasoline, which would be prohibitively expensive. Thus there are sound economic reasons for reforming the kerosene fractions to produce toluene and other aromatic compounds, which have high octane ratings and are much cheaper than MTBE.

Exercise

As shown in Figure 2.25 "The Octane Ratings of Some Hydrocarbons and Common Additives", toluene is one of the fuels suitable for use in automobile engines. How much toluene would have to be added to a blend of the petroleum fraction in this example containing 15% MTBE by mass to increase the octane rating to that of premium gasoline (93)?

Answer: The final blend is 56% toluene by mass, which requires a ratio of 56/44, or 1.3 tons of toluene per ton of blend.

Sulfuric Acid

Sulfuric acid is one of the oldest chemical compounds known. It was probably first prepared by alchemists who burned sulfate salts such as FeSO₄·7H₂O, called green vitriol from its color and glassy appearance (from the Latin vitrum, meaning “glass”). Because pure sulfuric acid was found to be useful for dyeing textiles, enterprising individuals looked for ways to improve its production. By the mid-18th century, sulfuric acid was being produced in multiton quantities by the lead-chamber process, which was invented by John Roebuck in 1746. In this process, sulfur was burned in a large room lined with lead, and the resulting fumes were absorbed in water.

Production

The production of sulfuric acid today is likely to start with elemental sulfur obtained through an ingenious technique called the Frasch process, which takes advantage of the low melting point of elemental sulfur (115.2°C). Large deposits of elemental sulfur are found in porous limestone rocks in the same geological formations that often contain petroleum. In the Frasch process, water at high temperature (160°C) and high pressure is pumped underground to melt the sulfur, and compressed air is used to force the liquid sulfur-water mixture to the surface (Figure 2.26 "Extraction of Elemental Sulfur from Underground Deposits"). The material that emerges from the ground is more than 99% pure sulfur. After it solidifies, it is pulverized and shipped in railroad cars to the plants that produce sulfuric acid, as shown here.

Transporting sulfur. A train carries elemental sulfur through the White Canyon of the Thompson River in British Columbia, Canada.

Figure 2.26 Extraction of Elemental Sulfur from Underground Deposits

In the Frasch process for extracting sulfur, very hot water at high pressure is injected into the sulfur-containing rock layer to melt the sulfur. The resulting mixture of liquid sulfur and hot water is forced up to the surface by compressed air.

An increasing number of sulfuric acid manufacturers have begun to use sulfur dioxide (SO₂) as a starting material instead of elemental sulfur. Sulfur dioxide is recovered from the burning of oil and gas, which contain small amounts of sulfur compounds. When not recovered, SO₂ is released into the atmosphere, where it is converted to an environmentally hazardous form that leads to acid rain (Chapter 4 "Reactions in Aqueous Solution").

If sulfur is the starting material, the first step in the production of sulfuric acid is the combustion of sulfur with oxygen to produce SO₂. Next, SO₂ is converted to SO₃ by the contact process, in which SO₂ and O₂ react in the presence of V₂O₅ to achieve about 97% conversion to SO₃. The SO₃ can then be treated with a small amount of water to produce sulfuric acid. Usually, however, the SO₃ is absorbed in concentrated sulfuric acid to produce oleum, a more potent form called fuming sulfuric acid. Because of its high SO₃ content (approximately 99% by mass), oleum is cheaper to ship than concentrated sulfuric acid. At the point of use, the oleum is diluted with water to give concentrated sulfuric acid (very carefully because dilution generates enormous amounts of heat). Because SO₂ is a pollutant, the small amounts of unconverted SO₂ are recovered and recycled to minimize the amount released into the air.