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Organic Chemistry: A Brief Course, 13th Edition, International Edition

David J. Hart, Christopher M. Hadad, Leslie E. Craine, and Harold Hart

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Bonding and Isomerism

Why does sucrose (table sugar) melt at 185°C, while sodium

chloride (table salt)—melts at a much higher temperature, 801°C? Why do both of these substances dissolve in water, while olive oil does not? Why does the molecule methyl butyrate smell like apples, while the molecule propyl acetate, which contains the same number and kind of atoms, smells like pears? To answer questions such as these, you must understand how atoms bond with one another and how molecules interact with one another. Bonding is the key to the structure, physical properties, and chemical behavior of different kinds of matter.

Perhaps you have already studied bonding and related concepts in a begin- ning chemistry course. Browse through each section of this chapter to see whether it is familiar, and try to work the problems. If you can work the problems, you can safely skip that section. But if you have difficulty with any of the problems within or at the end of this chapter, study the entire chapter carefully because we will use the ideas developed here throughout the rest of the book.

1.1How Electrons Are Arranged in Atoms

1.2Ionic and Covalent Bonding

1.3Carbon and the Covalent Bond

1.4Carbon–Carbon Single Bonds

1.5Polar Covalent Bonds

1.6Multiple Covalent Bonds

1.7Valence

A CLOSER LOOK AT… Green Chemistry

1.8Isomerism

1.9Writing Structural Formulas

1.10Abbreviated Structural Formulas

1.11Formal Charge

1.12Resonance

1.13Arrow Formalism

1.14The Orbital View of Bonding; the Sigma Bond

1.15Carbon sp3 Hybrid Orbitals

1.16Tetrahedral Carbon; the Bonding in Methane

1.17Classification According to Molecular Framework

1.18Classification According to Functional Group

An atom consists of a small, dense nucleus containing positively charged protons and neutral neutrons and surrounded by negatively charged electrons. The atomic number of an element equals the number of protons in its nucleus; its atomic weight is the sum of the number of protons and neutrons in its nucleus.

Electrons are located in orbitals. Orbitals are grouped in shells. An orbital can hold a maximum of two electrons.

Valence electrons are located in the outermost shell. The kernel of the atom contains the nucleus and the inner electrons.

Atoms contain a small, dense nucleus surrounded by electrons. The nucleus is positively charged and contains most of the mass of the atom. The nucleus consists of protons, which are positively charged, and neutrons, which are neutral. (The only exception is hydrogen, whose nucleus consists of only a single proton.) In a neutral atom, the positive charge of the nucleus is exactly balanced by the negative charge of the electrons that surround it. The atomic number of an element is equal to the number of protons in its nucleus (and to the number of electrons around the nucleus in a neutral atom). The atomic weight is approximately equal to the sum of the number of protons and the number of neutrons in the nucleus; the electrons are not counted because they are very light by comparison. The periodic table on the inside back cover of this book shows all the elements with their atomic numbers and weights.

We are concerned here mainly with the atom’s electrons because their number and arrangement provide the key to how a particular atom reacts with other atoms to form molecules. Also, we will deal only with electron arrangements in the lighter elements because these elements are the most important in organic molecules.

Electrons are concentrated in certain regions of space around the nucleus called orbitals. Each orbital can contain a maximum of two electrons. The orbitals, which differ in shape, are designated by the letters s, p, and d. In addition, orbitals are grouped in shells designated by the numbers 1, 2, 3, and so on. Each shell contains different types and numbers of orbitals, corresponding to the shell number. For example, shell 1 contains only one type of orbital, designated the 1s orbital. Shell 2 contains two types of orbitals, 2s and 2p, and shell 3 contains three types, 3s, 3p, and 3d. Within a particu- lar shell, the number of s, p, and d orbitals is 1, 3, and 5, respectively (Table 1.1). These rules permit us to count how many electrons each shell will contain when it is filled (last column in Table 1.1). Table 1.2 shows how the electrons of the first 18 elements are arranged.

The first shell is filled for helium (He) and all elements beyond, and the second shell is filled for neon (Ne) and all elements beyond. Filled shells play almost no role in chemical bonding. Rather, the outer electrons, or valence electrons, are mainly involved in chemical bonding, and we will focus our attention on them.

Table 1.3 shows the valence electrons, the electrons in the outermost shell, for the first 18 elements. The element’s symbol stands for the kernel of the element (the nucleus plus the filled electron shells), and the dots represent the valence electrons. The elements are arranged in groups according to the periodic table, and (except for helium) these group numbers correspond to the number of valence electrons.

Armed with this information about atomic structure, we are now ready to tackle the problem of how elements combine to form chemical bonds.

Table 1.1 Numbers of Orbitals and Electrons in the First Three Shells

Number of orbitals of

each type

1.2Ionic and Covalent Bonding

An early, but still useful, theory of chemical bonding was proposed in 1916 by Gilbert Newton Lewis, then a professor at the University of California, Berkeley. Lewis noticed that the inert gas helium had only two electrons surrounding its nucleus and that the next inert gas, neon, had 10 such electrons (2 1 8; see Table 1.2). He concluded that atoms of these gases must have very stable electron arrangements because these ele- ments do not combine with other atoms. He further suggested that other atoms might react in such a way in order to achieve these stable arrangements. This stability could be achieved in one of two ways: by complete transfer of electrons from one atom to another or by sharing of electrons between atoms.

An inert gas has a stable electron configuration.

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Ionic compounds are composed of positively charged cations and negatively charged anions.

Electropositive atoms give up electrons and form cations.

Electronegative atoms accept electrons and form anions.

Figure 1.1

Sodium chloride, NaCl, is an ionic crystal. The purple spheres represent sodium ions, Na1, and the green spheres are chloride ions, Cl2. Each ion is surrounded by six oppositely charged ions, except for those ions that are at the surface of the crystal.

Chapter1 Bonding and Isomerism

1.2.a Ionic Compounds

Ionic bonds are formed by the transfer of one or more valence electrons from one atom to another. Because electrons are negatively charged, the atom that gives up electrons becomes positively charged, a cation. The atom that receives electrons becomes negatively charged, an anion. The reaction between sodium and chlorine atoms to form sodium chloride (ordinary table salt) is a typical electron-transfer reaction.*

The sodium atom has only one valence electron (it is in the third shell; see Table 1.2). By giving up that electron, sodium achieves the electron arrangement of neon. At the same time, it becomes positively charged, a sodium cation. The chlorine atom has seven valence electrons. By accepting an additional electron, chlorine achieves the electron arrangement of argon and becomes negatively charged, a chloride anion. Atoms, such as sodium, that tend to give up electrons are said to be electropositive. Often such atoms are metals. Atoms, such as chlorine, that tend to accept electrons are said to be electronegative. Often such atoms are nonmetals.

EXAMPLE 1.1

Write an equation for the reaction of magnesium (Mg) with fluorine (F) atoms.

Mg F F Mg 2 2 F

Solution Magnesium has two valence electrons. Since each fluorine atom can accept only one electron (from the magnesium) to complete its valence shell, two fluorine atoms are needed to react with one magnesium atom.

PROBLEM 1.1 Write an equation for the reaction of lithium atoms (Li) with bromine atoms (Br).

The product of eq. 1.1 is sodium chloride, an ionic compound made up of equal numbers of sodium and chloride ions. In general, ionic compounds form when strongly electropositive atoms and strongly electronegative atoms interact. The ions in a crystal of an ionic substance are held together by the attractive force between their opposite charges, as shown in Figure 1.1 for a sodium chloride crystal.

In a sense, the ionic bond is not really a bond at all. Being oppositely charged, the ions attract one another like the opposite poles of a magnet. In the crystal, the ions are packed in a definite arrangement, but we cannot say that any particular ion is bonded or connected to any other particular ion. And, of course, when the sub- stance is dissolved, the ions separate and are able to move about in solution relatively freely.

*The curved arrow in eq. 1.1 shows the movement of one electron from the valence shell of the sodium atom to the valence shell of the chlorine atom. The use of curved arrows to show the movement of electrons is explained in greater detail in Section 1.13.

Solution

a.The lithium nucleus has less positive charge (13) to attract electrons than the beryllium nucleus (14). It takes less energy, therefore, to remove an electron from lithium than it does to remove one from beryllium. Since lithium loses an electron more easily than beryllium, lithium is the more electropositive atom.

b.The valence electron in the sodium atom is shielded from the positive charge of the nucleus by two inner shells of electrons, whereas the valence electron of lithium is shielded by only one inner shell. It takes less energy, therefore, to remove an electron from sodium; so, sodium is the more electropositive element.

PROBLEM 1.3 Using Table 1.3, determine which is the more electropositive element: sodium or aluminum, boron or carbon, boron or aluminum.

PROBLEM 1.4 Using Table 1.3, determine which is the more electronegative element: oxygen or fluorine, oxygen or nitrogen, fluorine or chlorine.

PROBLEM 1.5 Judging from its position in Table 1.3, do you expect carbon to be electropositive or electronegative?

1.2.b The Covalent Bond

Elements that are neither strongly electronegative nor strongly electropositive, or that have similar electronegativities, tend to form bonds by sharing electron pairs rather than completely transferring electrons. A covalent bond involves the mutual sharing of one or more electron pairs between atoms. Two (or more) atoms joined by covalent bonds constitute a molecule. When the two atoms are identical or have

A covalent bond is formed when two atoms share one or more electron pairs. A molecule consists of two or more atoms joined by covalent bonds.

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Bond energy (BE) is the energy necessary to break a mole of covalent bonds. The amount of energy depends on the type of bond broken.

The bond length is the average distance between two covalently bonded atoms.

Chapter1 Bonding and Isomerism

equal electronegativities, the electron pairs are shared equally. The hydrogen mol- ecule is an example.

Each hydrogen atom can be considered to have filled its first electron shell by the sharing process. That is, each atom is considered to “own” all of the electrons it shares with the other atom, as shown by the loops in these structures.

H H H H

EXAMPLE 1.4

Write an equation similar to eq. 1.2 for the formation of a chlorine molecule from two chlorine atoms.

Cl Cl Cl Cl heat

Solution One electron pair is shared by the two chlorine atoms. In that way, each chlorine completes its valence shell with eight electrons (three unshared pairs and one shared pair).

PROBLEM 1.6 Write an equation similar to eq. 1.2 for the formation of a fluorine molecule from two fluorine atoms.

When two hydrogen atoms combine to form a molecule, heat is liberated. Con- versely, this same amount of heat (energy) has to be supplied to a hydrogen molecule to break it apart into atoms. To break apart 1 mole (2 g) of hydrogen molecules into atoms requires 104 kcal (or 435 kJ*) of heat, quite a lot of energy. This energy is called the bond energy, or BE, and is different for bonds between different atoms (see Table A in the Appendix).

The H!H bond is a very strong bond. The main reason for this is that the shared electron pair is attracted to both hydrogen nuclei, whereas in a hydrogen atom, the valence electron is associated with only one nucleus. But other forces in the hydro- gen molecule tend to counterbalance the attraction between the electron pair and the nuclei. These forces are the repulsion between the two like-charged nuclei and the repulsion between the two like-charged electrons. A balance is struck between the attractive and the repulsive forces. The hydrogen atoms neither fly apart nor do they fuse together. Instead, they remain connected, or bonded, and vibrate about some equi- librium distance, which we call the bond length. For a hydrogen molecule, the bond length (that is, the average distance between the two hydrogen nuclei) is 0.74 Å.** The length of a covalent bond depends on the atoms that are bonded and the number of electron pairs shared between the atoms. Bond lengths for some typical covalent bonds are given in Table B in the Appendix.

*Although most organic chemists use the kilocalorie as the unit of heat energy, the currently used international unit is the kilojoule; 1 kcal 5 4.184 kJ. In this text, the kilocalorie will be used. If your instructor prefers to use kJ, multiply kcal 3 4.184 (or 3 4 for a rough estimate) to convert to kJ.

**Å, or angstrom unit, is 1028 cm, so the H!H bond length is 0.74 3 1028 cm. Although the angstrom is com- monly used by organic chemists, another unit often used for bond lengths is the picometer (pm; 1 Å 5 100 pm). To convert the H!H bond length from Å to pm, multiply 0.74 3 100. The H!H bond length is 74 pm. In this text, the angstrom will be used as the unit for bond lengths.

A radical is a molecular fragment with an odd number of unshared electrons.

The carbon–carbon bond in ethane, like the hydrogen–hydrogen bond in a hydrogen molecule, is a purely covalent bond, with the electrons being shared equally between the two identical carbon atoms. As with the hydrogen molecule, heat is required to break the carbon–carbon bond of ethane to give two CH3 fragments (called methyl radicals). A radical is a molecular fragment with an odd number of unshared electrons.

However, less heat is required to break the carbon–carbon bond in ethane than is required to break the hydrogen–hydrogen bond in a hydrogen molecule. The actual amount is 88 kcal (or 368 kJ) per mole of ethane. The carbon–carbon bond in ethane is longer (1.54 Å) than the hydrogen–hydrogen bond (0.74 Å) and also somewhat weaker. Breaking carbon–carbon bonds by heat, as represented in eq. 1.3, is the first step in the cracking of petroleum, an important process in the manufacture of gasoline (see “A Word About Petroleum, Gasoline, and Octane Number” on pages 102–103).

There is almost no limit to the number of carbon atoms that can be linked, and some molecules contain as many as 100 or more carbon–carbon bonds. This ability of an element to form chains as a result of bonding between the same atoms is called catenation.

PROBLEM 1.9 Using the structure of ethane as a guide, draw the structure for propane, C3H8.

Catenation is the ability of an element to form chains of its own atoms through covalent bonding.

1.5Polar Covalent Bonds

As we have seen, covalent bonds can be formed not only between identical atoms (H!H, C!C) but also between different atoms (C!H, C!Cl), provided that the atoms do not differ too greatly in electronegativity. However, if the atoms are different from one another, the electron pair may not be shared equally between them. Such a bond is sometimes called a polar covalent bond because the atoms that are linked carry a partial negative and a partial positive charge.

The hydrogen chloride molecule provides an example of a polar covalent bond. Chlorine atoms are more electronegative than hydrogen atoms, but even so, the bond that they form is covalent rather than ionic. However, the shared electron pair is attracted more toward the chlorine, which therefore is slightly negative with respect to the hydrogen. This bond polarization is indicated by an arrow whose head is negative and whose tail is marked with a plus sign. Alternatively, a partial charge, written as d1 or d2 (read as “delta plus” or “delta minus”), may be shown:

The bonding electron pair, which is shared unequally, is displaced toward the chlorine.

You can usually rely on the periodic table to determine which end of a polar covalent bond is more negative and which end is more positive. As we proceed from left to right across the table within a given period, the elements become more electro- negative, owing to increasing atomic number or charge on the nucleus. The increasing nuclear charge attracts valence electrons more strongly. As we proceed from the top to the bottom of the table within a given group (down a column), the elements become less electronegative because the valence electrons are shielded from the nucleus by an increasing number of inner-shell electrons. From these generalizations, we can safely predict that the atom on the right in each of the following bonds will be negative with respect to the atom on the left:

The carbon–hydrogen bond, which is so common in organic compounds, requires special mention. Carbon and hydrogen have nearly identical electronegativities, so the C!H bond is almost purely covalent. The electronegativities of some common elements are listed in Table 1.4.

A polar covalent bond is a covalent bond in which the electron pair is not shared equally between the two atoms.

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Chapter1 Bonding and Isomerism

Table 1.4 Electronegativities of Some Common Elements

Group

I

2.7

EXAMPLE 1.7

Indicate any bond polarization in the structure of tetrachloromethane.

Solution

Cl

A

Cl O CO Cl

A

Cl

Chlorine is more electronegative than carbon. The electrons in each C!Cl bond are therefore displaced toward the chlorine.

PROBLEM 1.10 Predict the polarity of the N!Cl bond and of the S!O bond.

PROBLEM 1.11 Draw the structure of the refrigerant dichlorodifluo- romethane, CCl2F2 (CFC-12), and indicate the polarity of the bonds. (The C atom is the central atom.)

PROBLEM 1.12 Draw the formula for methanol, CH3OH, and (where appro- priate) indicate the bond polarity with an arrow, ?: . (The C atom is bonded to three H atoms and the O atom.)

To complete their valence shells, atoms may sometimes share more than one electron pair. Carbon dioxide, CO2, is an example. The carbon atom has four valence electrons, and each oxygen has six valence electrons. A structure that allows each atom to com- plete its valence shell with eight electrons is

In structure A, the dots represent the electrons from carbon, and the x’s are the electrons from the oxygens. Structure B shows the bonds’ and oxygens’ unshared electrons, and structure C shows only the covalent bonds. Two electron pairs are shared between carbon and oxygen. Consequently, the bond is called a double bond. Each oxygen atom also has two pairs of nonbonding electrons, or unshared electron pairs. The loops in the following structures show that each atom in carbon dioxide has a complete valence shell of eight electrons:

Hydrogen cyanide, HCN, is an example of a simple compound with a triple bond, a bond in which three electron pairs are shared.

hydrogen cyanide

PROBLEM 1.13 Show with loops how each atom in hydrogen cyanide completes its valence shell.

EXAMPLE 1.8

Determine what, if anything, is wrong with the following electron arrange- ment for carbon dioxide:

O C O

Solution The formula contains the correct total number of valence electrons (16), and each oxygen is surrounded by 8 valence electrons, which is correct. However, what is wrong is that the carbon atom has 10 valence electrons, 2 more than is allowable.

PROBLEM 1.14 Show what is wrong with each of the following electron arrangements for carbon dioxide:

PROBLEM 1.15 Methanal (formaldehyde) has the formula H2CO. Draw a structure that shows how the valence electrons are arranged.

PROBLEM 1.16 Draw an electron-dot structure for carbon monoxide, CO.

In a double bond, two electron pairs are shared between two atoms.

Nonbonding electrons, or unshared electron pairs, reside on one atom.

In a triple bond, three electron pairs are shared between two atoms.

Hydrocarbons are compounds composed of just hydrogen and carbon atoms.

The valence of an element is the number of bonds that an atom of the element can form.

Carbon atoms can be connected to one another by double bonds or triple bonds, as well as by single bonds. Thus, there are three hydrocarbons (compounds with just carbon and hydrogen atoms) that have two carbon atoms per molecule: ethane, ethene, and ethyne.

They differ in that the carbon–carbon bond is single, double, or triple, respectively. They also differ in the number of hydrogens. As we will see later, these compounds have different chemical reactivities because of the different types of bonds between the carbon atoms.

EXAMPLE 1.9

Draw the structure for C3H6 having one carbon–carbon double bond.

Solution First, draw the three carbons with one double bond.

C " C ! C

Then add the hydrogens in such a way that each carbon has eight electrons around it (or in such a way that each carbon has four bonds).

H H H

A A A

H O C P C O C O H

A

H

PROBLEM 1.17 Draw three different structures that have the formula C4H8 and have one carbon–carbon double bond.

1.7Valence

The valence of an element is simply the number of bonds that an atom of the ele- ment can form. The number is usually equal to the number of electrons needed to fill the valence shell. Table 1.5 gives the common valences of several elements. Notice the difference between the number of valence electrons and the valence. Oxygen, for example, has six valence electrons but a valence of only 2. The sum of the two numbers is equal to the number of electrons in the filled shell.

The valences in Table 1.5 apply whether the bonds are single, double, or triple. For example, carbon has four bonds in each of the structures we have written so far: methane, tetrachloromethane, ethane, ethene, ethyne, carbon dioxide, and so on. These common valences are worth remembering, because they will help you to write correct structures.

Table 1.5 Valences of Common Elements

1.8Isomerism

The molecular formula of a substance tells us the numbers of different atoms present, but a structural formula tells us how those atoms are arranged. For example, H2O is the molecular formula for water. It tells us that each water molecule contains two hydrogen atoms and one oxygen atom. But the structural formula H!O!H tells us

The molecular formula of a

substance gives the number of different atoms present; the

structural formula indicates how

those atoms are arranged.

Conduct research on the Internet to find more informa- tion on green chemistry and to answer the following questions.

Green Chemistry

1. What is green chemistry?

Hangers Cleaners

1.How does Hangers Cleaners make use of carbon dioxide?

2.What are the advantages of using CO2 over traditional dry cleaning solvents? Explain why you think so.

2.What are the 12 principles of green chemistry?

3.What is the Pollution Prevention Act of 1990?

4.What are some organizations involved in promoting green chemistry? How are these organizations promoting green chemistry?

Ionic Liquids and Their Uses

1.What are some of the advantages of using ionic liquids as solvents in chemical reactions?

2.What environmental problems posed by traditional solvents are avoided by using ionic liquids?

Isomers are molecules with the same number and kinds of atoms but different arrangements of the atoms. Structural (or constitutional) isomers have the same molecular formula but different structural formulas.

more than that. The structural formula gives us the connectivity between atoms and tells us that the hydrogens are connected to the oxygen (and not to each other).

It is sometimes possible to arrange the same atoms in more than one way and still satisfy their valences. Molecules that have the same kinds and numbers of atoms but different arrangements are called isomers, a term that comes from the Greek (isos, equal, and meros, part). Structural (or constitutional) isomers are compounds that have the same molecular formula, but different structural formulas. Let us look at a particular pair of isomers.

Two very different chemical substances are known, each with the molecular formula C2H6O. One of these substances is a colorless liquid that boils at 78.5°C, whereas the other is a colorless gas at ordinary temperatures (boiling point (bp) 223.6°C). The only possible explanation is that the atoms must be arranged differ- ently in the molecules of each substance and that these arrangements are somehow responsible for the fact that one substance is a liquid and the other is a gas.

For the molecular formula C2H6O, two (and only two) structural formulas are possible that satisfy the valence requirement of 4 for carbon, 2 for oxygen, and 1 for hydrogen. They are:

In one formula, the two carbons are connected to one another by a single covalent bond; in the other formula, each carbon is connected to the oxygen. When we complete the valences by adding hydrogens, each arrangement requires six hydrogens. Many kinds of experimental evidence verify these structural assignments. We leave for later chapters (Chapters 7 and 8) an explanation of why these arrangements of atoms produce sub- stances that are so different from one another.

Ethanol and methoxymethane are structural isomers. They have the same molecular formula but different structural formulas. Ethanol and methoxymethane differ in physical and chemical properties as a consequence of their different molecular structures. In general, structural isomers are different compounds. They differ in phys- ical and chemical properties as a consequence of their different molecular structures.

PROBLEM 1.20 Draw structural formulas for the three possible isomers of C3H8O.

In a continuous chain, atoms are bonded one after another.

1.9Writing Structural Formulas

You will be writing structural formulas throughout this course. Perhaps a few hints about how to do so will be helpful. Let’s look at another case of isomerism. Suppose we want to write out all possible structural formulas that correspond to the molecular formula C5H12. We begin by writing all five carbons in a continuous chain.

C ! C ! C ! C ! C

a continuous chain

This chain uses up one valence for each of the end carbons and two valences for the carbons in the middle of the chain. Each end carbon therefore has three valences

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1.9 Writing Structural Formulas

left for bonds to hydrogens. Each middle carbon has only two valences for bonds to hydrogens. As a consequence, the structural formula in this case is written as:

pentane, bp 36C

To find structural formulas for the other isomers, we must consider branched chains. For example, we can reduce the longest chain to only four carbons and connect the fifth carbon to one of the middle carbons, as in the following structural formula:

C O C O C O C

A

C

a branched chain

If we add the remaining bonds so that each carbon has a valence of 4, we see that three of the carbons have three hydrogens attached, but the other carbons have only one or two hydrogens. The molecular formula, however, is still C5H12.

H O C O H

A

H

2-methylbutane, bp 28C

(isopentane)

Suppose we keep the chain of four carbons and try to connect the fifth carbon somewhere else. Consider the following chains:

C O C OC O C C O C OC O C C O C OC O C

AAA

CCC

Do we have anything new here? No! The first two structures have five-carbon chains, exactly as in the formula for pentane, and the third structure is identical to the branched chain we have already drawn for 2-methylbutane—a four-carbon chain with a one-carbon branch attached to the second carbon in the chain (counting now from the right instead of from the left). Notice that for every drawing of pentane, you can draw a line through all five carbon atoms without lifting your pencil from the paper. For every drawing of 2-methylbutane, a continuous line can be drawn through exactly four carbon atoms.*

But there is a third isomer of C5H12. We can find it by reducing the longest chain to only three carbons and connecting two one-carbon branches to the middle carbon.

C

A

C O C O C

A

C

*Using a molecular model kit (see note on p. 37) to construct the carbon chains as drawn will help you to see which representations are identical and which are different.

15

In a branched chain, some atoms form branches from the longest continuous chain.

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Chapter1 Bonding and Isomerism

If we fill in the hydrogens, we see that the middle carbon has no hydrogens attached to it.

H

A

H O C O H

HH

AA

H O C OO C OO C O H

AA

HH

H OC O H

A

H

2,2-dimethylpropane, bp 10C

(neopentane)

So we can draw three (and only three) different structural formulas that corre- spond to the molecular formula C5H12, and in fact, we find that only three different chemical substances with this formula exist. They are commonly called n-pentane (n for normal, with an unbranched carbon chain), isopentane, and neopentane.

PROBLEM 1.21 To which isomer of C5H12 does each of the following structural formulas correspond?

H

1.10Abbreviated Structural Formulas

Structural formulas like the ones we have written so far are useful, but they are also some- what cumbersome. They take up a lot of space and are tiresome to write out. Consequently, we often take some shortcuts that still convey the meaning of structural formulas. For example, we may abbreviate the structural formula of ethanol (ethyl alcohol) from:

HH

A A

Each formula clearly represents ethanol rather than its isomer methoxymethane (dimethyl ether), which can be represented by any of the following structures:

H H

AA

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Chapter1 Bonding and Isomerism

bond between the second and third carbon atoms (that is, CH3CH" CHCH2CH3) is represented as follows:

Three line segments emanate from this point; therefore, this carbon has one hydrogen (4 3 1) attached to it.

Two line segments emanate from this point; therefore, this carbon has two hydrogens (4 2 2) attached to it.

One line segment emanates from this point; therefore, this carbon has three hydrogens (4 1 3) attached to it.

EXAMPLE 1.13

Write a line-segment formula for CH3CH2CH" CHCH2CH(CH3)2 .

Solution

PROBLEM 1.24 Write a line-segment formula for (CH3)2CHCH2CH(CH3)2.

1.11Formal Charge

So far, we have considered only molecules whose atoms are neutral. But in some molecules, one or more atoms may be charged, either positively or negatively. Because such charges usually affect the chemical reactions of such molecules, it is important to know how to tell where the charge is located.

Consider the formula for hydronium ion, H3O1, the product of the reaction of a water molecule with a proton.

hydronium ion

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1.11 Formal Charge

The structure has eight electrons around the oxygen and two electrons around each hydrogen, so that all valence shells are complete. Note that there are eight valence electrons altogether. Oxygen contributes six, and each hydrogen contributes one, for a total of nine, but the ion has a single positive charge, so one electron must have been given away, leaving eight. Six of these eight electrons are used to form three O!H single bonds, leaving one unshared electron pair on the oxygen.

Although the entire hydronium ion carries a positive charge, we can ask, “Which atom, in a formal sense, bears the charge?” To determine formal charge, we consider each atom to “own” all of its unshared electrons plus only half of its shared electrons (one electron from each covalent bond). We then subtract this total from the number of valence electrons in the neutral atom to get the formal charge. This definition can be expressed in equation form as follows:

Let us apply this definition to the hydronium ion.

Thus, it is the oxygen atom that formally carries the 11 charge in the hydronium ion.

EXAMPLE 1.14

On which atom is the formal charge in the hydroxide ion, OH2?

Solution The electron-dot formula is

O H

Oxygen contributes six electrons, hydrogen contributes one, and there is one more for the negative charge, for a total of eight electrons. The formal charge on oxygen is 6 2 (6 1 1) 5 21, so the oxygen carries the nega- tive charge. (So instead, you might see hydroxide written as HO2 to reflect the negative charge on oxygen.) The hydrogen is neutral.

PROBLEM 1.25 Calculate the formal charge on the nitrogen atom in ammonia, NH3; in the ammonium ion, NH41; and in the amide ion, NH22.

Now let us look at a slightly more complex situation involving electron-dot structures and formal charge.

19

The formal charge on an atom in a covalently bonded molecule or ion is the number of valence electrons in the neutral atom minus the number of covalent bonds to the atom and the number of unshared electrons on the atom.

Resonance structures of a molecule or ion are two or more structures with identical arrangements of the atoms but different arrangements of the electrons. If resonance structures can be written, the true structure of the molecule or ion is a resonance hybrid of the contributing resonance structures.

In electron-dot structures, a pair of dots or a dash represents a bond between just two atoms. But sometimes, an electron pair is involved with more than two atoms in the process of forming bonds. Molecules and ions in which this occurs cannot be adequately represented by a single electron-dot structure. As an example, consider the structure of the carbonate ion, CO322.

The total number of valence electrons in the carbonate ion is 24 (4 from the carbon, 3 3 6 5 18 from the three oxygens, plus 2 more electrons that give the ion its negative charge; these 2 electrons presumably have been donated by some metal, per- haps one each from two sodium atoms). An electron-dot structure that completes the valence shell of eight electrons around the carbon and each oxygen is

carbonate ion, CO32

The structure contains two carbon–oxygen single bonds and one carbon–oxygen double bond. Application of the definition for formal charge shows that the carbon is formally neutral, each singly bonded oxygen has a formal charge of 21, and the doubly bonded oxygen is formally neutral.

PROBLEM 1.26 Show that the last sentence of the preceding paragraph is correct.

When we wrote the electron-dot structure for the carbonate ion, our choice of which oxygen atom would be doubly bonded to the carbon atom was purely arbitrary. There are in fact three exactly equivalent structures that we might write.

three equivalent structures for the carbonate ion

In each structure there is one C " O bond and there are two C!O bonds. These structures have the same arrangement of the atoms. They differ from one another only in the arrangement of the electrons.

The three structures for the carbonate ion are redrawn below, with curved arrows to show how electron pairs can be moved to convert one structure to another:

Chemists use curved arrows to keep track of a change in the location of electrons. A detailed explanation of the use of curved arrows is given in Section 1.13.

Physical measurements tell us that none of the foregoing structures accurately describes the real carbonate ion. For example, although each structure shows two dif- ferent types of bonds between carbon and oxygen, we find experimentally that all three carbon–oxygen bond lengths are identical: 1.31 Å. This distance is intermediate between the normal C " O (1.20 Å) and C!O (1.41 Å) bond lengths. To explain this fact, we usually say that the real carbonate ion has a structure that is a resonance hy- brid of the three contributing resonance structures. It is as if we could take an average

O

C O

O

2

1.13Arrow Formalism

Arrows in chemical drawings have specific meanings. For example, in Section 1.12 we used curved arrows to move electrons to show the relatedness of the three resonance structures of the carbonate ion. Just as it is important to learn the structural representa- tions and names of molecules, it is important to learn the language of arrow formalism in organic chemistry.

1.Curved arrows are used to show how electrons are moved in resonance structures and in reactions. Therefore, curved arrows always start at the initial position of electrons and end at their final position. In the example given below, the arrow that points from the C " O bond to the oxygen atom in the structure on the left indicates that the two electrons in one of the covalent bonds between carbon and oxygen are moved onto the oxygen atom:

Note that the carbon atom in the structure on the right now has a formal positive charge, and the oxygen has a formal negative charge. Notice also that when a pair of electrons in a polar covalent bond is moved to one of the bonded atoms, it is moved to the more electronegative atom, in this case oxygen. In the following example,

Curved arrows show how electrons are moved in resonance structures and in reactions.

Fishhook arrows indicate the movement of only a single electron.

Straight arrows point from reactants to products in chemical reaction equations.

A double-headed straight arrow between two structures indicates resonance structures.

the arrow that points from the unshared pair of electrons on the oxygen atom to a point between the carbon and oxygen atoms in the structure on the left indicates that the unshared pair of electrons on the oxygen atom moves between the oxygen and carbon atoms to form a covalent bond:

Note that both carbon and oxygen have formal charges of 0 in the structure on the right.

A curved arrow with half a head is called a fishhook. This kind of arrow is used to indicate the movement of a single electron. In eq. 1.6, two fishhooks are used to show the movement of each of the two electrons in the C—C bond of ethane to a carbon atom, forming two methyl radicals (see eq. 1.3):

2.Straight arrows point from reactants to products in chemical reaction equations. An example is the straight arrow pointing from ethane to the two methyl radicals in eq. 1.6. Straight arrows with half-heads are commonly used in pairs to indicate that a reaction is reversible.

A 1 B EF C 1 D

Adouble-headed straight arrow (G) between two structures indicates that they are resonance structures. Such an arrow does not indicate the occurrence of a chemical reaction. The double-headed arrows between resonance structures (Sec. 1.12) for the C " O bond are shown above.

EXAMPLE 1.15

Using correct arrow formalism, write the contributors to the resonance hybrid structure of the acetate ion, CH3CO22. Indicate any formal charges.

Solution There are two equivalent resonance structures for the acetate ion. Each one has a formal negative charge on one of the oxygen atoms.

Notice that when one pair of electrons from oxygen is moved to form a covalent bond with carbon, a pair of electrons in a covalent bond between carbon and the other oxygen atom is moved to oxygen. This is necessary to ensure that the carbon atom does not exceed its valence of 4.

PROBLEM 1.28 Using correct arrow formalism, write the contributors to the resonance hybrid of azide ion, a linear ion with three connected nitrogens, N32. Indicate the formal charge on each nitrogen atom.

We will use curved arrows throughout this text as a way of keeping track of elec- tron movement. Several curved-arrow problems are included at the end of this chapter to help you get used to drawing them.

1.14The Orbital View of Bonding; the Sigma Bond

Although electron-dot structures are often useful, they have some limitations. The Lewis theory of bonding itself has some limitations, especially in explaining the three- dimensional geometries of molecules. For this purpose in particular, we will discuss how another theory of bonding, involving orbitals, is more useful.

The atomic orbitals named in Section 1.1 have definite shapes. The s orbitals are spherical. The electrons that fill an s orbital confine their movement to a spherical region of space around the nucleus. The three p orbitals are dumbbell shaped and mu- tually perpendicular, oriented along the three coordinate axes, x, y, and z. Figure 1.2 shows the shapes of these orbitals.

Figure 1.2

The shapes of the s and p orbitals used by the valence electrons of carbon. The nucleus is at the origin of the three coordinate axes.

Z

X Y

In the orbital view of bonding, atoms approach each other in such a way that their atomic orbitals can overlap to form a bond. For example, if two hydrogen atoms form a hydrogen molecule, their two spherical 1s orbitals combine to form a new orbital that encompasses both of the atoms (see Figure 1.3). This orbital contains both valence electrons (one from each hydrogen). Like atomic orbitals, each molecular orbital can contain no more than two electrons. In the hydrogen molecule, these electrons mainly occupy the space between the two nuclei.

A molecular orbital is the space occupied by electrons in a molecule.

Figure 1.3

The molecular orbital representation of covalent bond formation between two hydrogen atoms.

The orbital in the hydrogen molecule is cylindrically symmetric along the H!H internuclear axis. Such orbitals are called sigma (S) orbitals, and the bond is referred to as a sigma bond. Sigma bonds may also be formed by the overlap of an s and a p orbital or of two p orbitals, as shown in Figure 1.4.*

*Two properly aligned p orbitals can also overlap to form another type of bond, called a p (pi) bond. We discuss this type of bond in Chapter 3.

A sigma (S) orbital lies along the axis between two bonded atoms; a pair of electrons in a sigma orbital is called a sigma bond.

Figure 1.4

Orbital overlap to form s bonds.

+

Figure 1.5

Distribution of the six electrons in a carbon atom. Each dot stands for an electron.

An sp3 hybrid orbital is a p-shaped orbital that is one part s and three parts p in character.

Figure 1.6

Unhybridized vs. sp3 hybridized orbitals on carbon. The dots stand for electrons. (Only the electrons in the valence shell are shown; the electrons in the 1s orbital are omitted because they are not involved in bonding.)

Let us see how these ideas apply to bonding in carbon compounds.

1.15Carbon sp3 Hybrid Orbitals

In a carbon atom, the six electrons are arranged as shown in Figure 1.5 (compare with carbon in Table 1.2). The 1s shell is filled, and the four valence electrons are in the 2s orbital and two different 2p orbitals. There are a few things to notice about Figure 1.5. The energy scale at the left represents the energy of electrons in the various orbitals. The farther the electron is from the nucleus, the greater its potential energy, because it takes energy to keep the electron (negatively charged) and the nucleus (positively charged) apart. The 2s orbital has a slightly lower energy than the three 2p orbitals, which have equal energies (they differ from one another only in orientation around the nucleus, as shown in Figure 1.2). The two highest energy electrons are placed in dif- ferent 2p orbitals rather than in the same orbital, because this keeps them farther apart and thus reduces the repulsion between these like-charged particles. One p orbital is vacant.

We might get a misleading idea about the bonding of carbon from Figure 1.5. For example, we might think that carbon should form only two bonds (to complete the partially filled 2p orbitals) or perhaps three bonds (if some atom donated two elec- trons to the empty 2p orbital). But we know from experience that this picture is wrong. Carbon usually forms four single bonds, and often these bonds are all equivalent, as in CH4 or CCl4. How can this discrepancy between theory and fact be resolved?

One solution, illustrated in Figure 1.6, is to mix or combine the four atomic orbitals of the valence shell to form four identical hybrid orbitals, each containing one valence electron. In this model, the hybrid orbitals are called SP3 hybrid orbitals because each one has one part s character and three parts p character. As shown in Figure 1.6, each sp3 orbital has the same energy: less than that of the 2p orbitals but greater than that of the 2s orbital. The shape of sp3 orbitals resembles the shape of p orbitals, except that the dumbbell is lopsided, and the electrons are more likely to be found in the lobe that extends out the greater distance from the nucleus, as shown in Figure 1.7. The four sp3 hybrid orbitals of a single carbon atom are directed toward the corners of a regular tetrahedron, also shown in Figure 1.7. This particular geometry puts each orbital as far from the other three orbitals as it can be and thus minimizes repulsion when the orbitals are filled with electron pairs. The angle between any two of the four bonds formed from sp3 orbitals is approximately 109.5°, the angle made by lines drawn from the center to the corners of a regular tetrahedron.

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1.16 Tetrahedral Carbon; the Bonding in Methane

109.5˚

SP3

25

Figure 1.7

An sp3 orbital extends mainly in one direction from the nucleus and forms bonds with other atoms in that direction. The four sp3 orbitals of any particular carbon atom are directed toward the corners of a regular tetrahedron, as shown in the right-hand part of the figure (in this part of the drawing, the small “back” lobes of the orbitals have been omitted for simplification, although they can be important in chemical reactions).

1.16Tetrahedral Carbon; the Bonding in Methane

We can now describe how a carbon atom combines with four hydrogen atoms to form methane. This process is pictured in Figure 1.9. The carbon atom is joined to each hydrogen atom by a sigma bond, which is formed by the overlap of a carbon sp3 orbital with a hydrogen 1s orbital. The four sigma bonds are directed from the carbon nucleus to the corners of a regular tetrahedron. In this way, the electron pair in any one bond experiences minimum repulsion from the electrons in the other bonds. Each H!C!H bond angle is the same, 109.5°. To summarize, in methane, there are four sp3–s C!H sigma bonds, each directed from the carbon atom to one of the four corners of a regular tetrahedron.

H

H

H

A bond angle is the angle made by two covalent bonds to the same atom.

Figure 1.9

A molecule of methane, CH4, is formed by the overlap of the four sp3 carbon orbitals with the 1s orbitals of four hydrogen atoms. The resulting molecule has the geometry of a regular tetrahedron and contains four sigma bonds of the sp3–s type.

Figure 1.10

The carbon and two of the hydrogens in methane form a plane that perpendicularly bisects the plane formed by the carbon and the other two hydrogens.

PROBLEM 1.29 Considering the repulsion that exists between electrons in different bonds, give a reason why a planar geometry for methane would be less stable than the tetrahedral geometry.

Because the tetrahedral geometry of carbon plays such an important role in organic chemistry, it is a good idea to become familiar with the features of a regular tetrahedron. One feature is that the center and any two corners of a tetrahedron form a plane that is the perpendicular bisector of a similar plane formed by the center and the other two corners. In methane, for example, any two hydrogens and the carbon form a plane that perpendicularly bisects the plane formed by the carbon and the other two hydrogens. These planes are illustrated in Figure 1.10.

The geometry of carbon with four single bonds, as in methane, is commonly rep- resented as shown in Figure 1.11a, in which the solid lines lie in the plane of the page, the dashed wedge goes behind the plane of the paper, and the solid wedge extends out of the plane of the paper toward you. Structures drawn in this way are sometimes called 3D (that is, three-dimensional) structures.

Two other 3D representations commonly used are the ball-and-stick model (Figure 1.11b) and the space-filling model (Figure 1.11c). The ball-and-stick model emphasizes the bonds that connect the atoms, while the space-filling model empha- sizes the space occupied by the atoms.

In addition, a 3D representation, called an electrostatic potential map, is sometimes used to show the distribution of electrons in a molecule (Figure 1.11d). Red indicates partial negative charge (greater electron density), and blue indicates partial positive

(a)In a 3D structure, solid lines lie in the plane of the page (C and H in C—H lie in the plane). Dashed wedges extend behind the plane (H in C^H lies behind the plane). Solid wedges project out toward you (H in CH is in front of the plane).

(b)A ball-and-stick model of a molecule emphasizes the bonds that connect atoms.

(c)A space-filling model emphasizes the space occupied by the atoms.

(d)An electrostatic potential map shows the distribution of electrons in a molecule. Red indicates partial negative charge, and blue indicates partial positive charge.

Figure 1.11

Four representations of methane.

109.5

A H

H

1.17Classification According to Molecular Framework

The three main classes of molecular frameworks for organic structures are acyclic, carbocyclic, and heterocyclic compounds.

1.17.a Acyclic Compounds

By acyclic (pronounced a´-cyclic), we mean not cyclic. Acyclic organic molecules have chains of carbon atoms but no rings. As we have seen, the chains may be unbranched or branched.

Acyclic compounds contain no rings. Carbocyclic compounds

contain rings of carbon atoms.

Heterocyclic compounds have rings

containing at least one atom that is not carbon.

Figure 1.12

O

CH3C(CH2)4CH3

2-heptanone (oil of cloves) bp 151.5C

A colorless liquid with a fruity odor, in part responsible for the “peppery” odor of blue cheese

Foodcollection/Getty Images

2-Heptanone contributes to the

Examples of natural acyclic compounds, their sources (in parentheses), and selected characteristics.

“peppery” odor of blue cheese.

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Figure 1.13

Examples of natural carbocyclic compounds with rings of various sizes and shapes. The source and special features of each structure are indicated below it.

Chapter1 Bonding and Isomerism

1.17.b Carbocyclic Compounds

Carbocyclic compounds contain rings of carbon atoms. The smallest possible car- bocyclic ring has three carbon atoms, but carbon rings come in many sizes and shapes. The rings may have chains of carbon atoms attached to them and may con- tain multiple bonds. Many compounds with more than one carbocyclic ring are known. Figure 1.13 shows the structures of a few carbocyclic compounds that occur in nature. Five- and six-membered rings are most common, but smaller and larger rings are also found.

O

© Dani, Jeske/Animals Animals

Musk deer, source of muscone.

-pinene (turpentine) bp 156.2C

A bicyclic molecule; one would have to break TWO bonds to make it acyclic

OH

O

testosterone

(testes)

mp 155C

A male sex hormone in which several rings of common sizes are FUSED together; that is, they share two adjacent carbon atoms

1.17 Classification According to Molecular Framework

carbon chains or rings attached to them, and in short may exhibit a great variety of structures. Figure 1.14 shows the structures of a few natural products that contain het- erocyclic rings. In these abbreviated structural formulas, the symbols for the hetero- atoms are shown, but the carbons are indicated using lines only.

The structures in Figures 1.12 through 1.14 show not only the molecular frame- works, but also various groups of atoms that may be part of or attached to the frameworks. Fortunately, these groups can also be classified in a way that helps simplify the study of organic chemistry.

29

© John Pontier / Animals Animals

Clover, a source of coumarin.

S

i

C O OH

B

O

nicotine

bp 246C

Present in tobacco, nicotine has two heterocyclic rings of different sizes, each containing one nitrogen.

adenine

mp 360–365C (decomposes)

One of the four hetero- cyclic bases of DNA, adenine contains two fused heterocyclic rings, each of

which contains two heteroatoms (nitrogen).

penicillin-G

(amorphous solid)

One of the most widely used antibiotics, penicillin has two heterocyclic rings, the smaller of which is crucial to biological activity.

coumarin mp 71C

Found in clover and grasses, coumarin produces the pleasant odor of new-mown hay.

Figure 1.14

S

SS

-terthienyl mp 92–93C

This compound, with three linked sulfur- containing rings, is present in certain marigold species.

OO B

O

B O

cantharidin

mp 218C

This compound, an oxygen heterocycle, is the active principle in cantharis (also known as Spanish fly), a material isolated from certain dried beetles of the

species CANTHARIS VESI-

CATORIA and incorrectly thought by some to increase sexual desire.

Functional groups are groups of atoms that have characteristic chemical properties regardless of the molecular framework to which they are attached.

Certain groups of atoms have chemical properties that depend only moderately on the molecular framework to which they are attached. These groups of atoms are called functional groups. The hydroxyl group, !OH, is an example of a functional group, and compounds with this group attached to a carbon framework are called alcohols. In most organic reactions, some chemical change occurs at the functional group, but the rest of the molecule keeps its original structure. This maintenance of most of the struc- tural formula throughout a chemical reaction greatly simplifies our study of organic chemistry. It allows us to focus attention on the chemistry of the various functional groups. We can study classes of compounds instead of having to learn the chemistry of each individual compound.

Some of the main functional groups that we will study are listed in Table 1.6, together with a typical compound of each type. Although we will describe these classes of compounds in greater detail in later chapters, it would be a good idea for you to become familiar with their names and structures now. If a particular functional group is mentioned before its chemistry is discussed in detail, and you forget what it is, you can refer to Table 1.6 or to the inside front cover of this book.

PROBLEM 1.30 What functional groups can you find in the following natural products? (Their formulas are given in Figures 1.12, 1.13, and 1.14.)

A. Functional groups that are a part of the molecular framework

B. Functional groups containing oxygen

1.With carbon–oxygen single bonds

A

ethyl acetate, a solvent for nail polish and model airplane glue

ethylamine, smells like ammonia

acrylonitrile, raw material for making Orlon

formamide, a softener for paper

methyl chloride, refrigerant and local anesthetic

G

*The DC P O group, present in several functional (a contraction of carbonyl and hydroxyl).

**The —NH2 group is called an amino group.

O

B

groups, is called a carbonyl group. The O C O OH group of acids is called a carboxyl group