Single Bond Character Definition Essay

Matter is made of combinations of elements—substances such as hydrogen or carbon that cannot be broken down or converted into other substances by chemical means. The smallest particle of an element that still retains its distinctive chemical properties is an atom. However, the characteristics of substances other than pure elements—including the materials from which living cells are made—depend on the way their atoms are linked together in groups to form molecules. In order to understand how living organisms are built from inanimate matter, therefore, it is crucial to know how all of the chemical bonds that hold atoms together in molecules are formed.

Cells Are Made From a Few Types of Atoms

Each atom has at its center a positively charged nucleus, which is surrounded at some distance by a cloud of negatively charged electrons, held in a series of orbitals by electrostatic attraction to the nucleus. The nucleus in turn consists of two kinds of subatomic particles: protons, which are positively charged, and neutrons, which are electrically neutral. The number of protons in the atomic nucleus gives the atomic number. An atom of hydrogen has a nucleus composed of a single proton; so hydrogen, with an atomic number of 1, is the lightest element. An atom of carbon has six protons in its nucleus and an atomic number of 6 (Figure 2-1). The electric charge carried by each proton is exactly equal and opposite to the charge carried by a single electron. Since an atom as a whole is electrically neutral, the number of negatively charged electrons surrounding the nucleus is equal to the number of positively charged protons that the nucleus contains; thus the number of electrons in an atom also equals the atomic number. It is these electrons that determine the chemical behavior of an atom, and all of the atoms of a given element have the same atomic number.

Figure 2-1

Highly schematic representations of an atom of carbon and an atom of hydrogen. Although the electrons are shown here as individual particles, in reality their behavior is governed by the laws of quantum mechanics, and there is no way of predicting exactly (more...)

Neutrons are uncharged subatomic particles of essentially the same mass as protons. They contribute to the structural stability of the nucleus—if there are too many or too few, the nucleus may disintegrate by radioactive decay—but they do not alter the chemical properties of the atom. Thus an element can exist in several physically distinguishable but chemically identical forms, called isotopes, each isotope having a different number of neutrons but the same number of protons. Multiple isotopes of almost all the elements occur naturally, including some that are unstable. For example, while most carbon on Earth exists as the stable isotope carbon 12, with six protons and six neutrons, there are also small amounts of an unstable isotope, the radioactive carbon 14, whose atoms have six protons and eight neutrons. Carbon 14 undergoes radioactive decay at a slow but steady rate. This forms the basis for a technique known as carbon 14 dating, which is used in archaeology to determine the time of origin of organic materials.

The atomic weight of an atom, or the molecular weight of a molecule, is its mass relative to that of a hydrogen atom. This is essentially equal to the number of protons plus neutrons that the atom or molecule contains, since the electrons are much lighter and contribute almost nothing to the total. Thus the major isotope of carbon has an atomic weight of 12 and is symbolized as 12C, whereas the unstable isotope just discussed has an atomic weight of 14 and is written as 14C. The mass of an atom or a molecule is often specified in daltons, one dalton being an atomic mass unit approximately equal to the mass of a hydrogen atom.

Atoms are so small that it is hard to imagine their size. An individual carbon atom is roughly 0.2 nm in diameter, so that it would take about 5 million of them, laid out in a straight line, to span a millimeter. One proton or neutron weighs approximately 1/(6 × 1023) gram, so one gram of hydrogen contains 6 × 1023 atoms. This huge number (6 × 1023, called Avogadro's number) is the key scale factor describing the relationship between everyday quantities and quantities measured in terms of individual atoms or molecules. If a substance has a molecular weight of X, 6 × 1023 molecules of it will have a mass of X grams. This quantity is called one mole of the substance (Figure 2-2).

There are 92 naturally occurring elements, each differing from the others in the number of protons and electrons in its atoms. Living organisms, however, are made of only a small selection of these elements, four of which—carbon (C), hydrogen (H), nitrogen (N), and oxygen (O)—make up 96.5% of an organism's weight. This composition differs markedly from that of the nonliving inorganic environment (Figure 2-3) and is evidence of a distinctive type of chemistry. The most common elements in living organisms are listed in Table 2-1 with some of their atomic characteristics.

Figure 2-3

The abundances of some chemical elements in the nonliving world (the Earth's crust) compared with their abundances in the tissues of an animal. The abundance of each element is expressed as a percentage of the total number of atoms present in the sample. (more...)

Table 2-1

Atomic Characteristics of the Most Abundant Elements in Living Tissues.

The Outermost Electrons Determine How Atoms Interact

To understand how atoms bond together to form the molecules that make up living organisms, we have to pay special attention to their electrons. Protons and neutrons are welded tightly to one another in the nucleus and change partners only under extreme conditions—during radioactive decay, for example, or in the interior of the sun or of a nuclear reactor. In living tissues, it is only the electrons of an atom that undergo rearrangements. They form the exterior of an atom and specify the rules of chemistry by which atoms combine to form molecules.

Electrons are in continuous motion around the nucleus, but motions on this submicroscopic scale obey different laws from those we are familiar with in everyday life. These laws dictate that electrons in an atom can exist only in certain discrete states, called orbitals, and that there is a strict limit to the number of electrons that can be accommodated in an orbital of a given type—a so-called electron shell. The electrons closest on average to the positive nucleus are attracted most strongly to it and occupy the innermost, most tightly bound shell. This shell can hold a maximum of two electrons. The second shell is farther away from the nucleus, and its electrons are less tightly bound. This second shell can hold up to eight electrons. The third shell contains electrons that are even less tightly bound; it can also hold up to eight electrons. The fourth and fifth shells can hold 18 electrons each. Atoms with more than four shells are very rare in biological molecules.

The electron arrangement of an atom is most stable when all the electrons are in the most tightly bound states that are possible for them—that is, when they occupy the innermost shells. Therefore, with certain exceptions in the larger atoms, the electrons of an atom fill the orbitals in order—the first shell before the second, the second before the third, and so on. An atom whose outermost shell is entirely filled with electrons is especially stable and therefore chemically unreactive. Examples are helium with 2 electrons, neon with 2 + 8, and argon with 2 + 8 + 8; these are all inert gases. Hydrogen, by contrast, with only one electron and therefore only a half-filled shell, is highly reactive. Likewise, the other atoms found in living tissues all have incomplete outer electron shells and are therefore able to donate, accept, or share electrons with each other to form both molecules and ions (Figure 2-4).

Figure 2-4

Filled and unfilled electron shells in some common elements. All the elements commonly found in living organisms have unfilled outermost shells (red) and can thus participate in chemical reactions with other atoms. For comparison, some elements that have (more...)

Because an unfilled electron shell is less stable than a filled one, atoms with incomplete outer shells have a strong tendency to interact with other atoms in a way that causes them to either gain or lose enough electrons to achieve a completed outermost shell. This electron exchange can be achieved either by transferring electrons from one atom to another or by sharing electrons between two atoms. These two strategies generate two types of chemical bonds between atoms: an ionic bond is formed when electrons are donated by one atom to another, whereas a covalent bond is formed when two atoms share a pair of electrons (Figure 2-5). Often, the pair of electrons is shared unequally, with a partial transfer between the atoms; this intermediate strategy results in a polar covalent bond, as we shall discuss later.

Figure 2-5

Comparison of covalent and ionic bonds. Atoms can attain a more stable arrangement of electrons in their outermost shell by interacting with one another. An ionic bond is formed when electrons are transferred from one atom to the other. A covalent bond (more...)

An H atom, which needs only one more electron to fill its shell, generally acquires it by electron sharing, forming one covalent bond with another atom; in many cases this bond is polar. The other most common elements in living cells—C, N, and O, with an incomplete second shell, and P and S, with an incomplete third shell (see Figure 2-4)—generally share electrons and achieve a filled outer shell of eight electrons by forming several covalent bonds. The number of electrons that an atom must acquire or lose (either by sharing or by transfer) to attain a filled outer shell is known as its valence.

The crucial role of the outer electron shell in determining the chemical properties of an element means that, when the elements are listed in order of their atomic number, there is a periodic recurrence of elements with similar properties: an element with, say, an incomplete second shell containing one electron will behave in much the same way as an element that has filled its second shell and has an incomplete third shell containing one electron. The metals, for example, have incomplete outer shells with just one or a few electrons, whereas, as we have just seen, the inert gases have full outer shells.

Ionic Bonds Form by the Gain and Loss of Electrons

Ionic bonds are most likely to be formed by atoms that have just one or two electrons in addition to a filled outer shell or are just one or two electrons short of acquiring a filled outer shell. They can often attain a completely filled outer electron shell more easily by transferring electrons to or from another atom than by sharing electrons. For example, from Figure 2-4 we see that a sodium (Na) atom, with atomic number 11, can strip itself down to a filled shell by giving up the single electron external to its second shell. By contrast, a chlorine (Cl) atom, with atomic number 17, can complete its outer shell by gaining just one electron. Consequently, if a Na atom encounters a Cl atom, an electron can jump from the Na to the Cl, leaving both atoms with filled outer shells. The offspring of this marriage between sodium, a soft and intensely reactive metal, and chlorine, a toxic green gas, is table salt (NaCl).

When an electron jumps from Na to Cl, both atoms become electrically charged ions. The Na atom that lost an electron now has one less electron than it has protons in its nucleus; it therefore has a single positive charge (Na+). The Cl atom that gained an electron now has one more electron than it has protons and has a single negative charge (Cl-). Positive ions are called cations, and negative ions, anions. Ions can be further classified according to how many electrons are lost or gained. Thus sodium and potassium (K) have one electron to lose and form cations with a single positive charge (Na+ and K+), whereas magnesium and calcium have two electrons to lose and form cations with two positive charges (Mg2+ and Ca2+).

Because of their opposite charges, Na+ and Cl- are attracted to each other and are thereby held together in an ionic bond. A salt crystal contains astronomical numbers of Na+ and Cl- (about 2 × 1019 ions of each type in a crystal 1 mm across) packed together in a precise three-dimensional array with their opposite charges exactly balanced (Figure 2-6). Substances such as NaCl, which are held together solely by ionic bonds, are generally called salts rather than molecules. Ionic bonds are just one of several types of noncovalent bonds that can exist between atoms, and we shall meet other examples.

Figure 2-6

Sodium chloride: an example of ionic bond formation. (A) An atom of sodium (Na) reacts with an atom of chlorine (Cl). Electrons of each atom are shown schematically in their different energy levels; electrons in the chemically reactive (incompletely filled) (more...)

Because of a favorable interaction between water molecules and ions, ionic bonds are greatly weakened by water; thus many salts (including NaCl) are highly soluble in water—dissociating into individual ions (such as Na+ and Cl-), each surrounded by a group of water molecules. In contrast, covalent bond strengths are not affected in this way.

Covalent Bonds Form by the Sharing of Electrons

All the characteristics of a cell depend on the molecules it contains. A molecule is defined as a cluster of atoms held together by covalent bonds; here electrons are shared between atoms to complete the outer shells, rather than being transferred between them. In the simplest possible molecule—a molecule of hydrogen (H2)—two H atoms, each with a single electron, share two electrons, which is the number required to fill the first shell. These shared electrons form a cloud of negative charge that is densest between the two positively charged nuclei and helps to hold them together, in opposition to the mutual repulsion between like charges that would otherwise force them apart. The attractive and repulsive forces are in balance when the nuclei are separated by a characteristic distance, called the bond length.

A further crucial property of any bond—covalent or noncovalent—is its strength. Bond strength is measured by the amount of energy that must be supplied to break that bond. This is often expressed in units of kilocalories per mole (kcal/mole), where a kilocalorie is the amount of energy needed to raise the temperature of one liter of water by one degree centigrade. Thus if 1 kilocalorie must be supplied to break 6 × 1023 bonds of a specific type (that is, 1 mole of these bonds), then the strength of that bond is 1 kcal/mole. An equivalent, widely used measure of energy is the kilojoule, which is equal to 0.239 kilocalories.

To get an idea of what bond strengths mean, it is helpful to compare them with the average energies of the impacts that molecules are constantly undergoing from collisions with other molecules in their environment (their thermal, or heat, energy), as well as with other sources of biological energy such as light and glucose oxidation (Figure 2-7). Typical covalent bonds are stronger than the thermal energies by a factor of 100, so they are resistant to being pulled apart by thermal motions and are normally broken only during specific chemical reactions with other atoms and molecules. The making and breaking of covalent bonds are violent events, and in living cells they are carefully controlled by highly specific catalysts, called enzymes. Noncovalent bonds as a rule are much weaker; we shall see later that they are important in the cell in the many situations where molecules have to associate and dissociate readily to carry out their functions.

Figure 2-7

Some energies important for cells. Note that these energies are compared on a logarithmic scale.

Whereas an H atom can form only a single covalent bond, the other common atoms that form covalent bonds in cells—O, N, S, and P, as well as the all-important C atom—can form more than one. The outermost shell of these atoms, as we have seen, can accommodate up to eight electrons, and they form covalent bonds with as many other atoms as necessary to reach this number. Oxygen, with six electrons in its outer shell, is most stable when it acquires an extra two electrons by sharing with other atoms and therefore forms up to two covalent bonds. Nitrogen, with five outer electrons, forms a maximum of three covalent bonds, while carbon, with four outer electrons, forms up to four covalent bonds—thus sharing four pairs of electrons (see Figure 2-4).

When one atom forms covalent bonds with several others, these multiple bonds have definite orientations in space relative to one another, reflecting the orientations of the orbits of the shared electrons. Covalent bonds between multiple atoms are therefore characterized by specific bond angles as well as bond lengths and bond energies (Figure 2-8). The four covalent bonds that can form around a carbon atom, for example, are arranged as if pointing to the four corners of a regular tetrahedron. The precise orientation of covalent bonds forms the basis for the three-dimensional geometry of organic molecules.

Figure 2-8

The geometry of covalent bonds. (A) The spatial arrangement of the covalent bonds that can be formed by oxygen, nitrogen, and carbon. (B) Molecules formed from these atoms have a precise three-dimensional structure, as shown here by ball and stick models (more...)

There Are Different Types of Covalent Bonds

Most covalent bonds involve the sharing of two electrons, one donated by each participating atom; these are called single bonds. Some covalent bonds, however, involve the sharing of more than one pair of electrons. Four electrons can be shared, for example, two coming from each participating atom; such a bond is called a double bond. Double bonds are shorter and stronger than single bonds and have a characteristic effect on the three-dimensional geometry of molecules containing them. A single covalent bond between two atoms generally allows the rotation of one part of a molecule relative to the other around the bond axis. A double bond prevents such rotation, producing a more rigid and less flexible arrangement of atoms (Figure 2-9 and Panel 2-1, pp. 111–112).

Figure 2-9

Carbon-carbon double bonds and single bonds compared. (A) The ethane molecule, with a single covalent bond between the two carbon atoms, illustrates the tetrahedral arrangement of single covalent bonds formed by carbon. One of the CH3 groups joined by (more...)

Panel 2-1

Chemical Bonds and Groups Commonly Encountered in Biological Molecules.

Some molecules share electrons between three or more atoms, producing bonds that have a hybrid character intermediate between single and double bonds. The highly stable benzenemolecule, for example, comprises a ring of six carbon atoms in which the bonding electrons are evenly distributed (although usually depicted as an alternating sequence of single and double bonds, as shown in Panel 2-1).

When the atoms joined by a single covalent bond belong to different elements, the two atoms usually attract the shared electrons to different degrees. Compared with a C atom, for example, O and N atoms attract electrons relatively strongly, whereas an H atom attracts electrons more weakly. By definition, a polar structure (in the electrical sense) is one with positive charge concentrated toward one end (the positive pole) and negative charge concentrated toward the other (the negative pole). Covalent bonds in which the electrons are shared unequally in this way are therefore known as polar covalent bonds (Figure 2-10). For example, the covalent bond between oxygen and hydrogen, -O-H, or between nitrogen and hydrogen, -N-H, is polar, whereas that between carbon and hydrogen, -C-H, has the electrons attracted much more equally by both atoms and is relatively nonpolar.

Figure 2-10

Polar and nonpolar covalent bonds. The electron distributions in the polar water molecule (H2O) and the nonpolar oxygen molecule (O2) are compared (δ+, partial positive charge; δ-, partial negative charge).

Polar covalent bonds are extremely important in biology because they create permanent dipoles that allow molecules to interact through electrical forces. Any large molecule with many polar groups will have a pattern of partial positive and negative charges on its surface. When such a molecule encounters a second molecule with a complementary set of charges, the two molecules will be attracted to each other by permanent dipole interactions that resemble (but are weaker than) the ionic bonds discussed previously for NaCl.

An Atom Often Behaves as if It Has a Fixed Radius

When a covalent bond forms between two atoms, the sharing of electrons brings the nuclei of these atoms unusually close together. But most of the atoms that are rapidly jostling each other in cells are located in separate molecules. What happens when two such atoms touch?

For simplicity and clarity, atoms and molecules are usually represented in a highly schematic way—either as a line drawing of the structural formula or as a ball and stick model. However, a more accurate representation can be obtained through the use of so-called space-filling models. Here a solid envelope is used to represent the radius of the electron cloud at which strong repulsive forces prevent a closer approach of any second, non-bonded atom—the so-called van der Waals radius for an atom. This is possible because the amount of repulsion increases very steeply as two such atoms approach each other closely. At slightly greater distances, any two atoms will experience a weak attractive force, known as a van der Waals attraction. As a result, there is a distance at which repulsive and attractive forces precisely balance to produce an energy minimum in each atom's interaction with an atom of a second, non-bonded element (Figure 2-11).

Figure 2-11

The balance of van der Waals forces between two atoms. As the nuclei of two atoms approach each other, they initially show a weak bonding interaction due to their fluctuating electric charges. However, the same atoms will strongly repel each other if (more...)

Depending on the intended purpose, we shall represent small molecules either as line drawings, ball and stick models, or space filling models throughout this book. For comparison, the water molecule is represented in all three ways in Figure 2-12. When dealing with very large molecules, such as proteins, we shall often need to further simplify the representation used (see, for example, Panel 3-2, pp. 138–139).

Figure 2-12

Three representations of a water molecule. (A) The usual line drawing of the structural formula, in which each atom is indicated by its standard symbol, and each line represents a covalent bond joining two atoms. (B) A ball and stick model, in which atoms (more...)

Water Is the Most Abundant Substance in Cells

Water accounts for about 70% of a cell's weight, and most intracellular reactions occur in an aqueous environment. Life on Earth began in the ocean, and the conditions in that primeval environment put a permanent stamp on the chemistry of living things. Life therefore hinges on the properties of water.

In each water molecule (H2O) the two H atoms are linked to the O atom by covalent bonds (see Figure 2-12). The two bonds are highly polar because the O is strongly attractive for electrons, whereas the H is only weakly attractive. Consequently, there is an unequal distribution of electrons in a water molecule, with a preponderance of positive charge on the two H atoms and of negative charge on the O (see Figure 2-10). When a positively charged region of one water molecule (that is, one of its H atoms) comes close to a negatively charged region (that is, the O) of a second water molecule, the electrical attraction between them can result in a weak bond called a hydrogen bond. These bonds are much weaker than covalent bonds and are easily broken by the random thermal motions due to the heat energy of the molecules, so each bond lasts only an exceedingly short time. But the combined effect of many weak bonds is far from trivial. Each water molecule can form hydrogen bonds through its two H atoms to two other water molecules, producing a network in which hydrogen bonds are being continually broken and formed (Panel 2-2, pp. 112–113). It is only because of the hydrogen bonds that link water molecules together that water is a liquid at room temperature, with a high boiling point and high surface tension—rather than a gas.

Panel 2-2

Water and Its Influence on the Behavior of Biological Molecules.

Molecules, such as alcohols, that contain polar bonds and that can form hydrogen bonds with water dissolve readily in water. As mentioned previously, molecules carrying plus or minus charges (ions) likewise interact favorably with water. Such molecules are termed hydrophilic, meaning that they are water-loving. A large proportion of the molecules in the aqueous environment of a cell necessarily fall into this category, including sugars, DNA, RNA, and a majority of proteins. Hydrophobic (water-hating) molecules, by contrast, are uncharged and form few or no hydrogen bonds, and so do not dissolve in water. Hydrocarbons are an important example (see Panel 2-1, pp. 110–111). In these molecules the H atoms are covalently linked to C atoms by a largely nonpolar bond. Because the H atoms have almost no net positive charge, they cannot form effective hydrogen bonds to other molecules. This makes the hydrocarbon as a whole hydrophobic—a property that is exploited in cells, whose membranes are constructed from molecules that have long hydrocarbon tails, as we shall see in Chapter 10.

Some Polar Molecules Form Acids and Bases in Water

One of the simplest kinds of chemical reaction, and one that has profound significance in cells, takes place when a molecule possessing a highly polarcovalent bond between a hydrogen and a second atom dissolves in water. The hydrogen atom in such a molecule has largely given up its electron to the companion atom and so exists as an almost naked positively charged hydrogen nucleus—in other words, a proton (H+). When the polar molecule becomes surrounded by water molecules, the proton is attracted to the partial negative charge on the O atom of an adjacent water molecule and can dissociate from its original partner to associate instead with the oxygen atoms of the water molecule to generate a hydronium ion (H3O+) (Figure 2-13A). The reverse reaction also takes place very readily, so one has to imagine an equilibrium state in which billions of protons are constantly flitting to and fro from one molecule in the solution to another.

Figure 2-13

Acids in water. (A) The reaction that takes place when a molecule of acetic acid dissolves in water. (B) Water molecules are continuously exchanging protons with each other to form hydronium and hydroxyl ions. These ions in turn rapidly recombine to form (more...)

Substances that release protons to form H3O+ when they dissolve in water are termed acids. The higher the concentration of H3O+, the more acidic the solution. H3O+ is present even in pure water, at a concentration of 10-7 M, as a result of the movement of protons from one water molecule to another (Figure 2-13B). By tradition, the H3O+ concentration is usually referred to as the H+ concentration, even though most H+ in an aqueous solution is present as H3O+. To avoid the use of unwieldy numbers, the concentration of H+ is expressed using a logarithmic scale called the pH scale, as illustrated in Panel 2-2 (pp. 112–113). Pure water has a pH of 7.0.

Because the proton of a hydronium ion can be passed readily to many types of molecules in cells, altering their character, the concentration of H3O+ inside a cell (the acidity) must be closely regulated. Molecules that can give up protons will do so more readily if the concentration of H3O+ in solution is low and will tend to receive them back if the concentration in solution is high.

The opposite of an acid is a base. Just as the defining property of an acid is that it donates protons to a water molecule so as to raise the concentration of H3O+ ions, the defining property of a base is that it raises the concentration of hydroxyl (OH-) ions—which are formed by removal of a proton from a water molecule. Thus sodium hydroxide (NaOH) is basic (the term alkaline is also used) because it dissociates in aqueous solution to form Na+ ions and OH- ions. Another class of bases, especially important in living cells, are those that contain NH2 groups. These groups can generate OH- by taking a proton from water: -NH2 + H2O → -NH3+ + OH-.

Because an OH-ion combines with a H3O+ ion to form two water molecules, an increase in the OH- concentration forces a decrease in the concentration of H3O+, and vice versa. A pure solution of water contains an equally low concentration (10-7 M) of both ions; it is neither acidic nor basic and is therefore said to be neutral with a pH of 7.0. The inside of cells is kept close to neutrality.

Four Types of Noncovalent Interactions Help Bring Molecules Together in Cells

In aqueous solutions, covalent bonds are 10 to 100 times stronger than the other attractive forces between atoms, allowing their connections to define the boundaries of one molecule from another. But much of biology depends on the specific binding of different molecules to each other. This binding is mediated by a group of noncovalent attractions that are individually quite weak, but whose bond energies can sum to create an effective force between two separate molecules. We have already introduced three of these noncovalent forces: ionic bonds, hydrogen bonds and van der Waals attractions. In Table 2-2, the strengths of these three types of bonds are compared to that of a typical covalent bond, both in the presence and the absence of water. Because of their fundamental importance in all biological systems, we shall summarize their properties here.

Table 2-2

Covalent and Noncovalent Chemical Bonds.

  • Ionic bonds. These are purely electrostatic attractions between oppositely charged atoms. As we saw for NaCl, these forces are quite strong in the absence of water. However, the polar water molecules cluster around both fully charged ions and polar molecules that contain permanent dipoles (Figure 2-14). This greatly reduces the potential attractiveness of these charged species for each other (see Table 2-2).

  • Hydrogen bonds. The structure of a typical hydrogen bond is illustrated in Figure 2-15. This bond represents a special form of polar interaction in which an electropositive hydrogen atom is partially shared by two electronegative atoms. Its hydrogen can be viewed as a proton that has partially dissociated from a donor atom, allowing it to be shared by a second acceptor atom. Unlike a typical electrostatic interaction, this bond is highly directional—being strongest when a straight line can be drawn between all three of the involved atoms. As already discussed, water weakens these bonds by forming competing hydrogen-bond interactions with the involved molecules (see Table 2-2).

  • van der Waals attractions. The electron cloud around any nonpolar atom will fluctuate, producing a flickering dipole. Such dipoles will transiently induce an oppositely polarized flickering dipole in a nearby atom. This interaction generates an attraction between atoms that is very weak. But since many atoms can be simultaneously in contact when two surfaces fit closely, the net result is often significant. These so-called van der Waals attractions are not weakened by water (see Table 2-2).

Figure 2-14

How the dipoles on water molecules orient to reduce the affinity of oppositely charged ions or polar groups for each other.

Figure 2-15

Hydrogen bonds. (A) Ball- and-stick model of a typical hydrogen bond. The distance between the hydrogen and the oxygen atom here is less than the sum of their van der Waals radii, indicating a partial sharing of electrons. (B) The most common hydrogen (more...)

The fourth effect that can play an important part in bringing molecules together in water is a hydrophobic force. This force is caused by a pushing of nonpolar surfaces out of the hydrogen-bonded water network, where they would physically interfere with the highly favorable interactions between water molecules. Because bringing two nonpolar surfaces together reduces their contact with water, the force is a rather nonspecific one. Nevertheless, we shall see in Chapter 3 that hydrophobic forces are central to the proper folding of protein molecules.

Panel 2-3 provides an overview of the four types of interactions just described. And Figure 2-16 illustrates, in a schematic way, how many such interactions can sum to hold together the matching surfaces of two macromolecules, even though each interaction by itself would be much too weak to be effective.

Panel 2-3

The Principal Types of Weak Noncovalent Bonds that Hold Macromolecules Together.

Figure 2-16

How two macro-molecules with complementary surfaces can bind tightly to one another through noncovalent interactions. In this schematic illustration, plus and minus are used to mark chemical groups that can form attractive interactions when paired.

A Cell Is Formed from Carbon Compounds

Having looked at the ways atoms combine into small molecules and how these molecules behave in an aqueous environment, we now examine the main classes of small molecules found in cells and their biological roles. We shall see that a few basic categories of molecules, formed from a handful of different elements, give rise to all the extraordinary richness of form and behavior shown by living things.

If we disregard water, nearly all the molecules in a cell are based on carbon. Carbon is outstanding among all the elements in its ability to form large molecules; silicon is a poor second. Because it is small and has four electrons and four vacancies in its outermost shell, a carbon atom can form four covalent bonds with other atoms. Most important, one carbon atom can join to other carbon atoms through highly stable covalent C-C bonds to form chains and rings and hence generate large and complex molecules with no obvious upper limit to their size (see Panel 2-1, pp. 110–111). The small and large carbon compounds made by cells are called organic molecules.

Certain combinations of atoms, such as the methyl (-CH3), hydroxyl (-OH), carboxyl (-COOH), carbonyl (-C=O), phosphate (-PO32-), and amino (-NH2) groups, occur repeatedly in organic molecules. Each such chemical group has distinct chemical and physical properties that influence the behavior of the molecule in which the group occurs. The most common chemical groups and some of their properties are summarized in Panel 2-1, pp. 110–111.

Cells Contain Four Major Families of Small Organic Molecules

The small organic molecules of the cell are carbon-based compounds that have molecular weights in the range 100 to 1000 and contain up to 30 or so carbon atoms. They are usually found free in solution and have many different fates. Some are used as monomer subunits to construct the giant polymeric macromolecules—the proteins, nucleic acids, and large polysaccharides—of the cell. Others act as energy sources and are broken down and transformed into other small molecules in a maze of intracellular metabolic pathways. Many small molecules have more than one role in the cell—for example, acting both as a potential subunit for a macromolecule and as an energy source. Small organic molecules are much less abundant than the organic macromolecules, accounting for only about one-tenth of the total mass of organic matter in a cell (Table 2-3). As a rough guess, there may be a thousand different kinds of these small molecules in a typical cell.

Table 2-3

The Approximate Chemical Composition of a Bacterial Cell.

All organic molecules are synthesized from and are broken down into the same set of simple compounds. Both their synthesis and their breakdown occur through sequences of chemical changes that are limited in scope and follow definite rules. As a consequence, the compounds in a cell are chemically related and most can be classified into a small number of distinct families. Broadly speaking, cells contain four major families of small organic molecules: the sugars, the fatty acids, the amino acids, and the nucleotides (Figure 2-17). Although many compounds present in cells do not fit into these categories, these four families of small organic molecules, together with the macromolecules made by linking them into long chains, account for a large fraction of cell mass (see Table 2-3).

Figure 2-17

The four main families of small organic molecules in cells. These small molecules form the monomeric building blocks, or subunits, for most of the macromolecules and other assemblies of the cell. Some, like the sugars and the fatty acids, are also energy (more...)

Sugars Provide an Energy Source for Cells and Are the Subunits of Polysaccharides

The simplest sugars—the monosaccharides—are compounds with the general formula (CH2O)n, where n is usually 3, 4, 5, 6, 7, or 8. Sugars, and the molecules made from them, are also called carbohydrates because of this simple formula. Glucose, for example, has the formula C6H12O6 (Figure 2-18). The formula, however, does not fully define the molecule: the same set of carbons, hydrogens, and oxygens can be joined together by covalent bonds in a variety of ways, creating structures with different shapes. As shown in Panel 2-4 (pp. 116–117), for example, glucose can be converted into a different sugar—mannose or galactose—simply by switching the orientations of specific OH groups relative to the rest of the molecule. Each of these sugars, moreover, can exist in either of two forms, called the d-form and the l-form, which are mirror images of each other. Sets of molecules with the same chemical formula but different structures are called isomers, and the subset of such molecules that are mirror-image pairs are called optical isomers. Isomers are widespread among organic molecules in general, and they play a major part in generating the enormous variety of sugars.

Figure 2-18

The structure of glucose, a simple sugar. As illustrated previously for water (see Figure 2-12), any molecule can be represented in several ways. In the structural formulas shown in (A), (B) and (E), the atoms are shown as chemical symbols linked together (more...)

Panel 2-4

An Outline of Some of the Types of Sugars Commonly Found in Cells.

An outline of sugar structures and chemistry is given in Panel 2-4. Sugars can exist in either a ring or an open-chain form. In their open-chain form, sugars contain a number of hydroxyl groups and either one aldehyde (H>C=O) or one ketone ( C=O) group. The aldehyde or ketone group plays a special role. First, it can react with a hydroxyl group in the same molecule to convert the molecule into a ring; in the ring form the carbon of the original aldehyde or ketone group can be recognized as the only one that is bonded to two oxygens. Second, once the ring is formed, this same carbon can become further linked to one of the carbons bearing a hydroxyl group on another sugar molecule, creating a disaccharide; such as sucrose, which is composed of a glucose and a fructose unit. Larger sugar polymers range from the oligosaccharides (trisaccharides, tetrasaccharides, and so on) up to giant polysaccharides, which can contain thousands of monosaccharide units.

The way that sugars are linked together to form polymers illustrates some common features of biochemical bond formation. A bond is formed between an -OH group on one sugar and an -OH group on another by a condensation reaction, in which a molecule of water is expelled as the bond is formed (Figure 2-19). Subunits in other biological polymers, such as nucleic acids and proteins, are also linked by condensation reactions in which water is expelled. The bonds created by all of these condensation reactions can be broken by the reverse process of hydrolysis, in which a molecule of water is consumed (see Figure 2-19).

Figure 2-19

The reaction of two monosaccharides to form a disaccharide. This reaction belongs to a general category of reactions termed condensation reactions, in which two molecules join together as a result of the loss of a water molecule. The reverse reaction (more...)

Because each monosaccharide has several free hydroxyl groups that can form a link to another monosaccharide (or to some other compound), sugar polymers can be branched, and the number of possible polysaccharide structures is extremely large. Even a simple disaccharide consisting of two glucose residues can exist in eleven different varieties (Figure 2-20), while three different hexoses (C6H12O6) can join together to make several thousand trisaccharides. For this reason it is a much more complex task to determine the arrangement of sugars in a polysaccharide than to determine the nucleotide sequence of a DNAmolecule, where each unit is joined to the next in exactly the same way.

Figure 2-20

Eleven disaccharides consisting of two D-glucose units. Although these differ only in the type of linkage between the two glucose units, they are chemically distinct. Since the oligosaccharides associated with proteins and lipids may have six or more (more...)

The monosaccharideglucose has a central role as an energy source for cells. In a series of reactions, it is broken down to smaller molecules, releasing energy that the cell can harness to do useful work, as we shall explain later. Cells use simple polysaccharides composed only of glucose units—principally glycogen in animals and starch in plants—as long-term stores of energy.

Sugars do not function only in the production and storage of energy. They also can be used, for example, to make mechanical supports. Thus, the most abundant organic chemical on Earth—the cellulose of plant cell walls—is a polysaccharide of glucose. Another extraordinarily abundant organic substance, the chitin of insect exoskeletons and fungal cell walls, is also a polysaccharide—in this case a linear polymer of a sugar derivative called N-acetylglucosamine. Polysaccharides of various other sorts are the main components of slime, mucus, and gristle.

Smaller oligosaccharides can be covalently linked to proteins to form glycoproteins and to lipids to form glycolipids, which are found in cell membranes. As described in Chapter 10, the surfaces of most cells are clothed and decorated with sugar polymers belonging to glycoproteins and glycolipids in the cell membrane. These sugar side chains are often recognized selectively by other cells. And differences between people in the details of their cell-surface sugars are the molecular basis for the major different human blood groups.

Fatty Acids Are Components of Cell Membranes

A fatty acidmolecule, such as palmitic acid, has two chemically distinct regions (Figure 2-21). One is a long hydrocarbon chain, which is hydrophobic and not very reactive chemically. The other is a carboxyl (-COOH) group, which behaves as an acid (carboxylic acid): it is ionized in solution (-COO-), extremely hydrophilic, and chemically reactive. Almost all the fatty acid molecules in a cell are covalently linked to other molecules by their carboxylic acid group.

Figure 2-21

A fatty acid. A fatty acid is composed of a hydrophobic hydrocarbon chain to which is attached a hydrophilic carboxylic acid group. Palmitic acid is shown here. Different fatty acids have different hydrocarbon tails. (A) Structural formula. The carboxylic (more...)

The hydrocarbon tail of palmitic acid is saturated: it has no double bonds between carbon atoms and contains the maximum possible number of hydrogens. Stearic acid, another one of the common fatty acids in animal fat, is also saturated. Some other fatty acids, such as oleic acid, have unsaturated tails, with one or more double bonds along their length. The double bonds create kinks in the molecules, interfering with their ability to pack together in a solid mass. It is this that accounts for the difference between hard (saturated) and soft (polyunsaturated) margarine. The many different fatty acids found in cells differ only in the length of their hydrocarbon chains and the number and position of the carbon-carbon double bonds (see Panel 2-5, pp. 118–119).

Fatty acids serve as a concentrated food reserve in cells, because they can be broken down to produce about six times as much usable energy, weight for weight, as glucose. They are stored in the cytoplasm of many cells in the form of droplets of triacylglycerol molecules, which consist of three fatty acid chains joined to a glycerolmolecule (see Panel 2-5); these molecules are the animal fats found in meat, butter, and cream, and the plant oils like corn oil and olive oil. When required to provide energy, the fatty acid chains are released from triacylglycerols and broken down into two-carbon units. These two-carbon units are identical to those derived from the breakdown of glucose and they enter the same energy-yielding reaction pathways, as will be described later in this chapter.

Fatty acids and their derivatives such as triacylglycerols are examples of lipids. Lipids comprise a loosely defined collection of biological molecules with the common feature that they are insoluble in water, while being soluble in fat and organic solvents such as benzene. They typically contain either long hydrocarbon chains, as in the fatty acids and isoprenes, or multiple linked aromatic rings, as in the steroids.

The most important function of fatty acids in cells is in the construction of cell membranes. These thin sheets enclose all cells and surround their internal organelles. They are composed largely of phospholipids, which are small molecules that, like triacylglycerols, are constructed mainly from fatty acids and glycerol. In phospholipids the glycerol is joined to two fatty acid chains, however, rather than to three as in triacylglycerols. The “third” site on the glycerol is linked to a hydrophilic phosphate group, which is in turn attached to a small hydrophilic compound such as choline (see Panel 2-5). Each phospholipidmolecule, therefore, has a hydrophobic tail composed of the two fatty acid chains and a hydrophilic head, where the phosphate is located. This gives them different physical and chemical properties from triacylglycerols, which are predominantly hydrophobic. Molecules like phospholipids, with both hydrophobic and hydrophilic regions, are termed amphipathic.

The membrane-forming property of phospholipids results from their amphipathic nature. Phospholipids will spread over the surface of water to form a monolayer of phospholipid molecules, with the hydrophobic tails facing the air and the hydrophilic heads in contact with the water. Two such molecular layers can readily combine tail-to-tail in water to make a phospholipid sandwich, or lipid bilayer. This bilayer is the structural basis of all cell membranes (Figure 2-22).

Figure 2-22

Phospholipid structure and the orientation of phospholipids in membranes. In an aqueous environment, the hydrophobic tails of phospholipids pack together to exclude water. Here they have formed a bilayer with the hydrophilic head of each phospholipid (more...)

Amino Acids Are the Subunits of Proteins

Amino acids are a varied class of molecules with one defining property: they all possess a carboxylic acid group and an amino group, both linked to a single carbon atom called the α-carbon (Figure 2-23). Their chemical variety comes from the side chain that is also attached to the α-carbon. The importance of amino acids to the cell comes from their role in making proteins, which are polymers of amino acids joined head-to-tail in a long chain that is then folded into a three-dimensional structure unique to each type of protein. The covalent linkage between two adjacent amino acids in a protein chain is called a peptide bond; the chain of amino acids is also known as a polypeptide (Figure 2-24). Regardless of the specific amino acids from which it is made, the polypeptide has an amino (NH2) group at one end (its N-terminus) and a carboxyl (COOH) group at its other end (its C-terminus). This gives it a definite directionality—a structural (as opposed to an electrical) polarity.

Figure 2-23

The amino acid alanine. (A) In the cell, where the pH is close to 7, the free amino acid exists in its ionized form; but when it is incorporated into a polypeptide chain, the charges on the amino and carboxyl groups disappear. (B) A ball-and-stick model (more...)

Sir William Herschel was the first to recognize the existence of infrared in 1800. Interest in IR was not explored further for 80 years. During 1882-1900 several investigations were made into the IR region. Abney and Festing photographed absorption spectra for 52 compounds and correlated absorption bands with the presence of certain organic groups in the molecule (Smith).

W. W. Coblentz laid the real groundwork for IR spectroscopy. Starting in 1903 he investigated the spectra of hundreds of substances, both organic and inorganic. His work in the rock salt region, from 0.7 to 18 (m, was so thorough and accurate that many of his spectra are still usable. The experimental difficulties of the early researchers were many. They not only had to design and build their own instruments but all the components too. Obtaining a spectrum was a tedious job requiring 3-4 hours or more since each point in the spectrum had to be measured separately and at least 10 points per micrometer were measured. After World War II advances in electronics made it possible to obtain a spectrum in 1-2 hours (Smith).

The end result of this early work was the recognition that each compound had a unique IR spectra and that certain groups, even when they were in different molecules, gave absorption bands that were found at approximately the same wavelength.

The IR absorption spectrum of a compound is its most unique physical property. The samples can be liquids, solids, or gases. They can be organic or inorganic. The only molecules transparent to IR radiation under ordinary conditions are monatomic and homonuclear molecules such as Ne, He, O2, N2, and H2. One limitation of IR spectroscopy is that the solvent water is a very strong absorber and attacks NaCl sample cells.

In terms of a comparison of physical properties, a melting point, refractive index, or specific gravity gives only a single point of comparison with other substances. An IR spectrum, in contrast, gives a multitude of such points. Not only can the position of bands be compared but their intensity as well since the intensity is indicative of the number of a particular group contributing to an absorption. IR is usually preferred when a combination of qualitative and quantitative analysis is required. It is often used to follow the course of organic reactions allowing the researcher to characterize the products as the reaction proceeds.

A term often encountered in IR spectroscopy is wavenumber ((), whose relationship to wavelength ( is ( (cm-1) = (104/() where ( is measured in micrometers. The wavenumber may be visualized as the number of wavelengths per centimeter.

All IR spectrometers have the following elements in common, the source, optical system, detector, and amplifier. In the region 100-4000 cm-1 the most popular sources are the Globar and the Nernst glower, which are heated electrically to about 1500 C. The purpose of the optical system is to channel the radiation along the proper path. Mirrors are used rather than lenses because lenses are subject to chromatic aberration.

Because infrared radiations are essentially radiant heat, thermal detectors are used to detect changes in the radiations. Thermal detectors are made as small as possible to reduce their heat capacity so that for a given amount of energy there will be a large temperature rise. In order to make the detector rapid in response it must be able to dissipate the heat very rapidly. There are three main types of detector.

The thermocouple uses the principle that the change in temperature of a junction of two dissimilar metals creates an electromotive force (emf) which may be measured. The bolometer uses the principle that the electrical resistance of a pure metal or semi-conductor is temperature-sensitive. If a constant potential is applied to such a detector the variation of the resistance with temperature may be measured by the variations in the current flowing in the circuit. In the Golay cell the detector is a small metal cylinder enclosed by a blackened metal plate at one end and a flexible metallised diaphragm at the other. The cylinder is filled with a gas and sealed. As IR radiation falls on the blackened plate the gas in the cylinder expands deforming the diaphragm. Light from a lamp inside the detector is focused on the diaphragm. The light is reflected from the metallised diaphragm and falls onto a photocell. Movement of the diaphragm moves the light beam across the photocell. The output of the photocell is proportional to the expansion of the gas.

Early IR spectrometers used optical amplification of the detector signal obtained from a galvanometer or radiometer. Current instrumentation uses chopped radiation with electronic amplification.

An infrared spectrometer may use either a single beam or double beam design. In the single beam design light from the radiation is focused and passed through a sample contained in a special cell. After passing through the sample the emergent light beam is dispersed by a monochromator, either a prism or a diffraction grating, into its component wavelengths. The spectrum is scanned by slowly rotating the prism or grating. The main difference in a double beam spectrometer is that the original light is split into two beams, one of which passes through the sample and the other through a reference cell. The instrument records the difference in intensity of these two beams. The double beam spectrometer is especially useful if the spectrum is to be measured in solution. In this case the reference cell would contain pure solvent and any absorption due to the solvent would be canceled out.

The prism and cells used in IR spectrometers cannot be made of glass because glass absorbs strongly in the infrared region of interest. The prism and sample cell walls are usually made from large NaCl or KBr crystals.

If we build an instrument in such a way as to allow a sample of unknown material to be held in position while various wavelengths of the IR region shine on it in turn, we can find which wavelengths are absorbed and which are not. This scan can be plotted on a graph and is called the infrared spectrum of the material. A spectrum is a plot of absorbance (or transmittance) versus wavelength, frequency, or wavenumber. For IR spectra, we usually plot wavenumber (in units of cm-1) which is the reciprocal of wavelength calculated as follows: cm-1= 1/(m x 104 .

Interpretation of Spectra

Certain atomic movements give rise to bands that occur in approximately the same position in a large variety of compounds and are only slightly affected by the rest of the molecule. These vibrations are assigned to groups of atoms termed functional groups. Although absorption bands are characteristic of the molecule as a whole, it is a useful approximation to consider that molecular vibrations are localized in particular functional groups. This allows one to relate absorption band position with a particular functional group and to tabulate these relationships. The tables showing the positions where the functional group absorptions occur are called correlation tables. The intensity of the absorption bands is also shown on good correlation tables.

Most of the spectral features that allow us to readily identify functional groups are found in the left part of the spectrum. The right hand portion of the spectrum is more complex, and each peak is not readily identified with a particular part of the molecule. The entire spectral pattern is unique for a given compound.

The steps used by a chemist to find information about molecular structure from the IR spectrum are as follows:

1. Obtain a spectrum of the material on an IR spectrophotometer.
2. Using information from correlation tables and absorbances from the functional group region of the spectrum, identify the functional groups that are present or sometimes more importantly absent.
3. Compare this spectrum with those of known compounds or obtain a known sample of a suspected material and run its spectrum for comparison.
Many of the group frequencies vary over a wide range because the bands arise from complex interacting vibrations within the molecule. Absorption bands may, however, represent predominantly a single vibrational mode. Certain absorption bands, for example, those arising from C-H, O-H, and C=O stretching modes, remain within fairly narrow regions of the spectrum.

The two important areas for a preliminary examination of a spectrum are the region 4000-1300 cm-1 and the 909-650 cm-1 region. The high frequency portion of the spectrum is called the functional group region. The characteristic stretching frequencies for important functional groups such as OH, NH, and C=O occur in this portion of the spectrum. The absence of absorption in the assigned ranges for the various functional groups can usually be used as evidence for the absence of such groups from the molecule. The absence of absorption in the 1850-1540 cm-1 region excludes a structure containing a carbonyl group. Strong skeletal bands for aromatics and heteroaromatics fall in the 1600-1300 cm-1 region of the spectrum. These skeletal bands arise from the stretching of the carbon-carbon bonds in the ring structure.

The lack of strong absorption bands in the 909-650 cm-1 region generally indicates a nonaromatic structure. Aromatic and heteroaromatic compounds display strong out-of-plane C-H bending and ring bending absorption bands in this region.

The intermediate portion of the spectrum, 1300-909 cm-1 is usually referred to as the fingerprint region. The absorption pattern in this region is complex, with bands originating in interacting vibrational modes. Absorption in this intermediate region is probably unique for every molecular species.

Conclusions reached after examination of a particular band should be confirmed by examination of other portions of the spectrum if possible. For example the assignment of a carbonyl band to the presence of an ester should be confirmed by observation of a strong band in the C-O stretching region, 1300-1100 cm-1.

Characteristic Group Frequencies of Organic Molecules Table 1 Characteristic Infrared Group Frequencies

Class Group Wavenumber (cm-1) AlkaneC-H2850-3000
C-H2700 &2900
Carboxylic AcidsC=O1700-1725
C-O1000-1300 (2 bands)


Hydrocarbons are classified as saturated or unsaturated based on the absence or presence of multiple bonds. The presence of multiple bonds decreases the number of hydrogens from the number in a saturated compound of formula CnH2n+2 . The decrease in number of hydrogens alone does not confirm the presence of a multiple bond. For example, 1-octene and cyclooctane have the same molecular formula, C8H16. What are the structural features that are present in 1-octene that are absent in cyclooctane? The carbon-carbon double bond and sp2-hybridized C-H bonds distinguish 1-octene from cyclooctane. It is characteristic group absorbances of these structures that will be present in the spectrum of 1-octene and absent in the spectrum of cyclooctane.

The energy of the infrared light absorbed by a C-H bond depends on the hybridization of the hybrid orbital. The bond strengths of carbon-hydrogen bonds are in the order of sp3>sp2>sp, because the increased s character of the hybrid gives better overlap with the hydrogen s-orbital. The sp3-hybridized C-H bonds in saturated hydrocarbons like octane absorb in the 2850-3000 cm-1 region. The sp2-hybridized C-H bonds in alkenes such as 1-octene absorbs at 3080 cm-1. A sp-hybridized C-H bond in a molecule such as 1-octyne absorbs infrared at 3320 cm-1.

Hydrocarbons can also be classified based on absorptions due to the carbon-carbon bond. Carbon-carbon bond strength increases in the order of singledoubletriple. Therefore, the wavenumber position of the absorption corresponding to the stretching of these bonds increases in the same order. Saturated hydrocarbons all contain carbon-carbon single bonds that absorb in the 800-1000 cm-1 region. Unsaturated hydrocarbons also contain carbon-carbon single bonds that absorb in this same region. This is not a very diagnostic region because we already know that most organic compounds have carbon-carbon single bonds.

Alkenes are identified by the absorption of the carbon-carbon double bond, which occurs in the 1630-1670 cm-1 region. Terminal alkenes have the most intense absorptions as the absorption decreases with increased substitution. Alkyne C(C stretches occur in a region of the IR spectrum where very little else appears. The alkyne carbon-carbon stretch occurs in the range of 2100-2260 cm-1. The intensity can very dramatically since the dipole moment change depends entirely upon what is attached to each carbon. Terminal alkynes, alkynes which have an H attached to one of the alkyne carbons, generally display greater band intensities as well as characteristic signals near 3300 cm-1.

Activity Four

Provide students with unlabelled spectra for octane, 1-octene, and 1-octyne (or any other straight chain hydrocarbons with the same number of carbon atoms). Ask students to assign the hydrocarbon to its spectra and justify their selections.

Oxygen-Containing Compounds

Many functional groups contain oxygen. These functional groups have the characteristic infrared absorptions given in Table 1. The characteristic group frequencies of aldehydes and ketones are from 1700-1780 cm-1. The carbon-oxygen double bond of carbonyl compounds requires more energy to stretch than does the carbon-oxygen single bond of ethers and alcohols. Therefore, aldehydes and ketones absorb infrared at higher wavenumber positions than alcohols and ethers. Since the carbonyl is highly polar, stretching of this bond results in a relatively large change in dipole moment producing an intense band. The carbonyl region is also free of conflicting absorptions making the recognition of the carbonyl band easy. Carefully examining the precise wavenumber of the C=O stretch, as well as the presence or absence of other signals, will usually allow one to distinguish among the many possible C=O containing compounds.

The position of the carbonyl group absorption of acyl derivatives depends on the inductive and resonance effects of atoms bonded to the carbonyl carbon atom. We can represent a carbonyl group by two contributing resonance structures. (See Figure 3) Since less energy is required to stretch a single bond than a double bond, any structural feature that stabilizes the contributing polar resonance form with a carbon-oxygen single bond will cause the infrared absorption to occur at lower wavenumber position. In other words, any group that donates electrons by resonance causes a shift in the absorption to lower wavenumbers. For example, the nitrogen atom of amides is very effective in donation of electrons to the carbonyl carbon atom. (See Figure 4) Therefore, the double bond character of the carbonyl decreases. As a result, in amides the carbonyl group absorbs in the 1650-1690 cm-1 region, which is at a lower wavenumber than for aldehydes or ketones.

When carbonyls (or other multiple bonds) are in conjugation with another double or triple bond a resonance form can be drawn in which the carbonyl oxygen bears a negative charge. The contribution of this resonance form reduces the double bond character of the carbonyl shifting the absorption to a lower frequency. (See Figure 5)

The characteristic bands observed for alcohols result from O-H stretching in addition to C-O stretching. The carbon-oxygen stretching vibration of alcohols appears in a region complicated by many other absorptions, the fingerprint region. The presence of a hydroxyl is better established by the O-H stretching. The shape and frequency of an O-H band depends on hydrogen bonding. As hydrogen bonding becomes stronger, O-H stretches appear at lower frequencies. In the vapor phase or in dilute solution in nonpolar solvents �free� hydroxyl group of alcohols absorbs strongly around 3600 cm-1. As the concentration of the solution increases intermolecular hydrogen bonding increases and we see additional bands start to appear at lower frequencies, 3550-3200 cm-1 and the �free� hydroxyl band decreases. When the hydrogen of a hydroxyl group is involved in a hydrogen bond a resonance form can be drawn in which the oxygen bears a negative charge. The contribution of this resonance form reduces the single bond character of the hydroxyl bond shifting the absorption to a lower frequency. (See Figure 6)

Carboxylic acids tend to form strongly hydrogen bonded dimers which shift the O-H stretch to frequencies lower than 3000, however, carboxylic O-H stretches can occur anywhere between 2500-3300 depending upon the strength of hydrogen bonding.

A process of elimination can identify ethers. If a compound contains oxygen and the infrared spectrum lacks absorptions characteristic of a carbonyl group or a hydroxyl group, we may conclude that the compound is an ether.

Activity Five

As we talk about the characteristic bands for functional groups students will be given actual IR spectra that illustrate each of the bands being discussed. At the culmination of the activity they will be given the spectra of carefully selected unknown compounds and will be asked to determine which functional groups are present and/or absent.

Activity Six

Students will be given the IR spectrum and other data about a particular compound, such as elemental analysis and molecular mass. They will be asked to draw a structural formula for the compound consistent with the information given and be required to justify their choice of structure.

Investigation 1: What factors affect the frequency of oscillation of a swing?

Work with a partner to design and conduct an experiment to determine what factors, if any, affect the frequency of oscillation of a swing. Frequency is defined as the number of complete oscillations per unit of time. For the purposes of comparison report your frequency in oscillations/minute. As you plan keep in mind the variables you want to control and the ones you want to test. The experiment will be conducted across the street at College Woods.

After gathering your data address the following questions in the results and conclusions section of your lab report.

1.From your data, which factor(s) affect the frequency of oscillation of a swing?
2.For each factor that you found affected the frequency, describe the relationship between frequency and the factor. In other words, how does varying the factor change the frequency?
Investigation 2: Masses on a spring

The bonds between atoms may be modeled using masses connected by springs. If one mass is held stationary and the other is allowed to move, the stretching vibration of a bond between the two atoms may be modeled by stretching the spring and releasing it.

Using the model above investigate the effect of mass on the frequency of vibration and the effect of bond strength on the frequency of vibration.

After gathering your data, address the following questions in the results and conclusions section of your lab report.

1.Did the mass affect the frequency of vibration? If so, quantitatively describe the relationship between mass and frequency?
2. Did the number of springs or spring tightness affect the frequency of vibration? If so, describe the relationship between spring tightness and the frequency.
3. Relate what you discovered about the effects of mass and spring strength to atoms joined by a chemical bond. Make a statement about the relationship between atomic mass and frequency and bond strength and frequency.
4. Compare your results from this investigation to the results for the swing in Investigation 1.
Worksheet 1: Predicting Vibrational Frequencies

Consider the following hypothetical diatomic molecules C-H, C-C, and C-I. The carbon atom is common to the molecules. If we were to hold the carbon atom in a clamp and represent the bond by a spring we could attach masses representing the H, C, and I atoms one by one and measure the frequency of vibration as we did in investigation 2. In your investigation you found that the frequency of vibration was inversely related to the mass (as mass increased the frequency decreased). In the infrared spectrum the infrared vibrations of these atoms will occur at:

C-H, 3000 cm-1 C-C, 1000 cm-1 C-I, 500 cm-1

This is consistent with our observation that as we increased the mass the frequency decreased.

If the diatomic grouping consisted of C-O, near to which of the three frequencies would you expect to find the absorption band?

Would the frequency of the absorption band be higher or lower than this frequency? Why?

Suppose now we connect two of these diatomic groupings having sufficiently different frequencies to make a hypothetical triatomic grouping. We now have two bond springs and there will be two ways to stretch each grouping.

( ( ( (
H - C - CC - C - I
( ( ( ( ( (: direction of motion of the atoms when the two springs are vibrating.

The bonds in the H-C-C grouping are found to vibrate practically independently, when the C-H bond vibrates the C-C hardly changes, and the spectrum of this compound would show two absorption bands at 3000 cm-1 and 1000 cm-1. Similarly in the C-C-I grouping the two vibrations are shown to be almost independent of each other and to occur very close to their diatomic positions.

Where would you expect to find the absorption bands for the C-C-I grouping?

Worksheet 1

Where would you expect to find the absorption bands for the grouping H-C-I ?

If we were to link two identical bonds together to make a C-C-C grouping we would still get two absorption bands but these would involve both of the bonds vibrating. The two frequencies would arise from an in-phase and an out-of-phase vibrational motion.


C- C- CC- C- C
( ( ( (
The two vibrations are termed symmetric and asymmetric. The two bonds do not move independently of each other, therefore, we say there is interaction between the groups and the frequencies of vibration will be displaced from the diatomic frequency.

In which of these groupings will there be symmetric and asymmetric vibrations?

H-C-I C- C- I O- C- O

If each bond is represented by a spring then the C-N will have one spring, the C=N will have two springs and the C(N will have three. This means the strength of the bond has been increased as you go from C-N to C(N. If the strength of the spring has been increased and the atoms remain constant the frequency of the vibration will also be increased. The frequencies of these groups are approximately:

C-N, 1070 cm-1 C=N, 1650 cm-1 C(N, 2250 cm-1

Consider the following triatomic groupings, which will have symmetric and asymmetric vibrations? Why?


Worksheet 2: Deformations

Vibrations other than the stretching of a bond also occur. These vibrations are called deformations and they refer to the bond angle changing between the atoms of a molecule. These deformations occur at frequencies lower than those of the stretching vibrations. If we consider groups of atoms of the type XY2 there are some general descriptions of the deformational vibrations we can apply to these groups.

The terms used are scissors, rock, wag, and twist. These four motions may be further divided into in-plane and out-of-plane motions of the atoms.

H(In-plane symmetric deformation
If you think of the C as the pivot and the H as the points of a pair of scissors then there is a plane through all of the atoms and the motion is in-plane.
CIn-plane asymmetric deformation

Consider the hydrogens to be the tips of a rocker on a rocking chair and the carbon is sitting on the chair. There is a plane through all of the atoms. This is a poor group frequency because all of the atoms move.

H +Out-of-plane symmetric deformation
H +
Consider the C to be the body of a horse and the H to be the tail. When the horse wags its tail the motion is from side to side and out of the plane of the horse (in the diagram the plane of the paper). The �+� indicates motion perpendicular to the plane of the paper. This is also a poor group frequency.
+ H H -Out-of-plane asymmetric deformation
Look at the diagram of the CH2ClBr molecule and consider the CH2 part only. What is the total number of vibrations that may occur?

Worksheet 2

The infrared spectrum has in its composition three types of absorption bands, fundamental, combination, and overtone. When we stretched the masses on the springs we observed that they vibrated at a certain frequency, this is the fundamental frequency for that system. The masses were vibrating with simple harmonic motion.

If a molecule with a fundamental vibration occurring at a frequency (1 is subjected to radiation at a frequency 2(1 , an absorption may also be observed at this frequency and is called an overtone absorption band. This does not mean that the molecule itself is vibrating at 2(1 , but the fundamental is being excited by the radiation at twice its frequency and such bands are generally much weaker than the fundamental. If a molecule has two different fundamental absorptions at (a and (b it is possible that an absorption may be observed at frequencies corresponding to (a + (b and (a -(b and these are called combination bands. Combination bands will usually be weaker than the fundamentals involved.

What would be the frequency of the fundamental absorption if its first overtone absorption was observed at 2000 cm-1?

Teacher Bibliography

Student Reading List

Classroom Materials

The following materials should be available for student use for the investigations: stopwatches, swings of various lengths, springs with different force constants or many springs with the same force constant, masses of different magnitude (washers may be used if commercial masses are not available), measuring tape, triple beam balances. If computers are available MacSpartan for the Macintosh or PC Spartan for Windows brings the vibrational motion of atoms within specific molecules to life.

Answer: 1000 cm-1 and 500 cm-1

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