NCBI Bookshelf. A service of the National Library of Medicine, National Institutes of Health.

Lodish H, Berk A, Zipursky SL, et al. Molecular Cell Biology. 4th edition. New York: W. H. Freeman; 2000.

  • By agreement with the publisher, this book is accessible by the search feature, but cannot be browsed.
Cover of Molecular Cell Biology

Molecular Cell Biology. 4th edition.

Show details

Section 2.2Noncovalent Bonds

Carbohydrates illustrate the importance of subtle differences in covalent bonds in generating molecules with different biological activities. However, several types of noncovalent bonds are critical in maintaining the three-dimensional structures of large molecules such as proteins and nucleic acids (see Figure 2-1b). Noncovalent bonds also enable one large molecule to bind specifically but transiently to another, making them the basis of many dynamic biological processes.

The energy released in the formation of noncovalent bonds is only 1 – 5 kcal/mol, much less than the bond energies of single covalent bonds (see Table 2-1). Because the average kinetic energy of molecules at room temperature (25 °C) is about 0.6 kcal/mol, many molecules will have enough energy to break noncovalent bonds. Indeed, these weak bonds sometimes are referred to as interactions rather than bonds. Although noncovalent bonds are weak and have a transient existence at physiological temperatures (25 –  37 °C), multiple noncovalent bonds often act together to produce highly stable and specific associations between different parts of a large molecule or between different macromolecules (Figure 2-11). In this section we consider the four main types of noncovalent bonds and discuss their role in stabilizing the structure of biomembranes.

Figure 2-11. Multiple weak bonds stabilize specific associations between large molecules.

Figure 2-11

Multiple weak bonds stabilize specific associations between large molecules. (Left) In this hypothetical complex, seven noncovalent bonds bind the two protein molecules A and B together, forming a stable complex. (Right) Because only four noncovalent (more...)

The Hydrogen Bond Underlies Water’s Chemical and Biological Properties

Hydrogen bonding between water molecules is of crucial importance because all life requires an aqueous environment and water constitutes about 70–80 percent of the weight of most cells. The mutual attraction of its molecules causes water to have melting and boiling points at least 100 °C higher than they would be if water were nonpolar; in the absence of these intermolecular attractions, water on earth would exist primarily as a gas. The exact structure of liquid water is still unknown. It is believed to contain many transient, maximally hydrogen-bonded networks. Most likely, water molecules are in rapid motion, constantly making and breaking hydrogen bonds with adjacent molecules. As the temperature of water increases toward 100 °C, the kinetic energy of its molecules becomes greater than the energy of the hydrogen bonds connecting them, and the gaseous form of water appears.

Properties of Hydrogen Bonds

Normally, a hydrogen atom forms a covalent bond with only one other atom. However, a hydrogen atom covalently bonded to a donor atom, D, may form an additional weak association, the hydrogen bond, with an acceptor atom, A:

Image ch2e12.jpg
In order for a hydrogen bond to form, the donor atom must be electronegative, so that the covalent D—H bond is polar. The acceptor atom also must be electronegative, and its outer shell must have at least one nonbonding pair of electrons that attracts the δ+ charge of the hydrogen atom. In biological systems, both donors and acceptors are usually nitrogen or oxygen atoms, especially those atoms in amino (—NH2) and hydroxyl (—OH) groups. Because all covalent N—H and O—H bonds are polar, their H atoms can participate in hydrogen bonds. By contrast, C—H bonds are nonpolar, so these H atoms are almost never involved in a hydrogen bond.

Water molecules provide a classic example of hydrogen bonding. The hydrogen atom in one water molecule is attracted to a pair of electrons in the outer shell of an oxygen atom in an adjacent molecule. Not only do water molecules hydrogen-bond with one another, they also form hydrogen bonds with other kinds of molecules, as shown in Figure 2-12. The presence of hydroxyl (—OH) or amino (—NH2) groups makes many molecules soluble in water. For instance, the hydroxyl group in methanol (CH3OH) and the amino group in methylamine (CH3NH2) can form several hydrogen bonds with water, enabling the molecules to dissolve in water to high concentrations. In general, molecules with polar bonds that easily form hydrogen bonds with water can dissolve in water and are said to be hydrophilic (Greek, “water-loving”). Besides the hydroxyl and amino groups, peptide and ester bonds are important chemical groups that interact well with water:

Image ch2e13.jpg

Figure 2-12. Water readily forms hydrogen bonds.

Figure 2-12

Water readily forms hydrogen bonds. In liquid water, each water molecule apparently forms transient hydrogen bonds with several others, creating a fluid network of hydrogen-bonded molecules (a). The precise structure of liquid water is still not known (more...)

Most hydrogen bonds are 0.26 – 0.31 nm long, about twice the length of covalent bonds between the same atoms. In particular, the distance between the nuclei of the hydrogen and oxygen atoms of adjacent hydrogen-bonded molecules in water is approximately 0.27 nm, about twice the length of the covalent O—H bonds in water. The hydrogen atom is closer to the donor atom, D, to which it remains covalently bonded, than it is to the acceptor. The length of the covalent D—H bond is a bit longer than it would be if there were no hydrogen bond, because the acceptor “pulls” the hydrogen away from the donor. The strength of a hydrogen bond in water (≈5 kcal/mol) is much weaker than a covalent O—H bond (≈110 kcal/mol).

Hydrogen Bonds as a Stabilizing Force in Macromolecules

An important feature of all hydrogen bonds is directionality. In the strongest hydrogen bonds, the donor atom, the hydrogen atom, and the acceptor atom all lie in a straight line. Nonlinear hydrogen bonds are weaker than linear ones; still, multiple nonlinear hydrogen bonds help to stabilize the three-dimensional structures of many proteins. It is only because of the aggregate strength of multiple hydrogen bonds that they play a central role in the architecture of large biological molecules in aqueous solutions (see Figure 2-11).

The strengths of the hydrogen bonds in proteins and nucleic acids are only 1 to 2 kcal/mol, considerably weaker than the hydrogen bonds between water molecules. The reason for this difference can be seen from Figure 2-13, which depicts the formation of a hydrogen bond between two amino acids in a protein. Initially, both the —OH and —NH2 groups in the protein are hydrogen-bonded to water, and the formation of a hydrogen bond between these groups involves disruption of their hydrogen bonds with water. Thus the net change in energy in forming this —OH···N hydrogen bond will be less than the 5 kcal/mol characteristic of hydrogen bonds between water molecules.

Figure 2-13. In order for a hydrogen bond (red dots) to form between a —OH and an —NH2 group in a protein (right), the hydrogen bonds between these groups and water must be disrupted (left).

Figure 2-13

In order for a hydrogen bond (red dots) to form between a —OH and an —NH2 group in a protein (right), the hydrogen bonds between these groups and water must be disrupted (left).

Ionic Interactions Are Attractions between Oppositely Charged Ions

In some compounds, the bonded atoms are so different in electronegativity that the bonding electrons are never shared: these electrons are always found around the more electronegative atom. In sodium chloride (NaCl), for example, the bonding electron contributed by the sodium atom is completely transferred to the chlorine atom. Even in solid crystals of NaCl, the sodium and chlorine atoms are ionized, so it is more accurate to write the formula for the compound as Na+Cl.

Because the electrons are not shared, the bonds in such compounds cannot be considered covalent. They are, rather, ionic bonds (or interactions) that result from the attraction of a positively charged ion — a cation — for a negatively charged ion — an anion. Unlike covalent or hydrogen bonds, ionic bonds do not have fixed or specific geometric orientations because the electrostatic field around an ion — its attraction for an opposite charge — is uniform in all directions. However, crystals of salts such as Na+Cl do have very regular structures because that is the energetically most favorable way of packing together positive and negative ions. The force that stabilizes ionic crystals is called the lattice energy.

In aqueous solutions, simple ions of biological significance, such as Na+, K+, Ca2+, Mg2+, and Cl, do not exist as free, isolated entities. Instead, each is surrounded by a stable, tightly held shell of water molecules (Figure 2-14). An ionic interaction occurs between the ion and the oppositely charged end of the water dipole, as shown below for the K+ ion:

Image ch2e14.jpg

Figure 2-14. In aqueous solutions, a shell of water molecules surrounds ions.

Figure 2-14

In aqueous solutions, a shell of water molecules surrounds ions. In the case of a magnesium ion (Mg2+), six water molecules are held tightly in place by electrostatic interactions between the two positive charges on the ion and the partial negative charge (more...)

Ions play an important biological role when they pass through narrow, protein-lined pores, or channels, in membranes. For example, ionic movements through membranes are essential for the conduction of nerve impulses and for the stimulation of muscle contraction. As we will see in Chapter 21, ions must lose their shell of water molecules in order to pass through ion channel proteins; channel proteins can then selectively admit only Na+, or K+, or Ca2+ ions, a selectivity essential for nerve function.

Most ionic compounds are quite soluble in water because a large amount of energy is released when ions tightly bind water molecules. This is known as the energy of hydration. Oppositely charged ions are shielded from one another by the water and tend not to recombine. Salts like Na+Cl dissolve in water because the energy of hydration is greater than the lattice energy that stabilizes the crystal structure. In contrast, certain salts, such as Ca3(PO4)2, are virtually insoluble in water; the large charges on the Ca2+ and PO43− ions generate a formidable lattice energy that is greater than the energy of hydration.

Van der Waals Interactions Are Caused by Transient Dipoles

When any two atoms approach each other closely, they create a weak, nonspecific attractive force that produces a van der Waals interaction, named for Dutch physicist Johannes Diderik van der Waals (1837 – 1923), who first described it. These nonspecific interactions result from the momentary random fluctuations in the distribution of the electrons of any atom, which give rise to a transient unequal distribution of electrons, that is, a transient electric dipole. If two noncovalently bonded atoms are close enough together, the transient dipole in one atom will perturb the electron cloud of the other. This perturbation generates a transient dipole in the second atom, and the two dipoles will attract each other weakly. Similarly, a polar covalent bond in one molecule will attract an oppositely oriented dipole in another.

Van der Waals interactions, involving either transient induced or permanent electric dipoles, occur in all types of molecules, both polar and nonpolar. In particular, van der Waals interactions are responsible for the cohesion between molecules of nonpolar liquids and solids, such as heptane, CH3—(CH2)5—CH3, that cannot form hydrogen bonds or ionic interactions with other molecules. When these stronger interactions are present, they override most of the influence of van der Waals interactions. Heptane, however, would be a gas if van der Waals interactions could not form.

The strength of van der Waals interactions decreases rapidly with increasing distance; thus these noncovalent bonds can form only when atoms are quite close to one another. However, if atoms get too close together, they become repelled by the negative charges in their outer electron shells. When the van der Waals attraction between two atoms exactly balances the repulsion between their two electron clouds, the atoms are said to be in van der Waals contact (Figure 2-15). Each type of atom has a van der Waals radius at which it is in van der Waals contact with other atoms. The van der Waals radius of an H atom is 0.1 nm, and the radii of O, N, C, and S atoms are between 0.14 and 0.18 nm. Two covalently bonded atoms are closer together than two atoms that are merely in van der Waals contact. For a van der Waals interaction, the internuclear distance is approximately the sum of the corresponding radii for the two participating atoms. Thus the distance between a C atom and an H atom in van der Waals contact is 0.27 nm, and between two C atoms is 0.34 nm. In general, the van der Waals radius of an atom is about twice as long as its covalent radius. For example, a C—H covalent bond is about 0.107 nm long and a C—C covalent bond is about 0.154 nm long.

Figure 2-15. Two oxygen molecules in van der Waals contact.

Figure 2-15

Two oxygen molecules in van der Waals contact. Transient dipoles in the electron clouds of all atoms give rise to weak attractive forces, called van der Waals interactions. Each type of atom has a characteristic van der Waals radius at which van der Waals (more...)

The energy of the van der Waals interaction is about 1 kcal/mol, only slightly higher than the average thermal energy of molecules at 25 °C. Thus the van der Waals interaction is even weaker than the hydrogen bond, which typically has an energy of 1 – 2 kcal/mol in aqueous solutions. The attraction between two large molecules can be appreciable, however, if they have precisely complementary shapes, so that they make many van der Waals contacts when they come into proximity. Van der Waals interactions, as well as other noncovalent bonds, mediate the binding of many enzymes with their specific substrates (the substances on which an enzyme acts) and of each type of antibody with its specific antigen (Chapter 3).

Hydrophobic Bonds Cause Nonpolar Molecules to Adhere to One Another

Nonpolar molecules do not contain ions, possess a dipole moment, or become hydrated. Because such molecules are insoluble or almost insoluble in water, they are said to be hydrophobic (Greek, “water-fearing”). The covalent bonds between two carbon atoms and between carbon and hydrogen atoms are the most common nonpolar bonds in biological systems. Hydrocarbons — molecules made up only of carbon and hydrogen — are virtually insoluble in water. A large triacylglycerol (or triglyceride) such as tristearin, a component of animal fat, is also insoluble in water, even though its six oxygen atoms participate in some slightly polar bonds with adjacent carbon atoms (Figure 2-16). When shaken in water, tristearin forms a separate phase similar to the separation of oil from the water-based vinegar in an oil-and-vinegar salad dressing.

Figure 2-16. The chemical structure of tristearin, or tristearoyl glycerol, a component of natural fats.

Figure 2-16

The chemical structure of tristearin, or tristearoyl glycerol, a component of natural fats. It contains three molecules of the fatty acid stearic acid, CH3(CH2)16COOH, esterified to one molecule of glycerol, HOCH2CH(OH)CH2OH. One end of the molecule (green) (more...)

The force that causes hydrophobic molecules or nonpolar portions of molecules to aggregate together rather than to dissolve in water is called the hydrophobic bond. This is not a separate bonding force; rather, it is the result of the energy required to insert a nonpolar molecule into water. A nonpolar molecule cannot form hydrogen bonds with water molecules, so it distorts the usual water structure, forcing the water into a rigid cage of hydrogen-bonded molecules around it. Water molecules are normally in constant motion, and the formation of such cages restricts the motion of a number of water molecules; the effect is to increase the structural organization of water. This situation is energetically unfavorable because it decreases the randomness (entropy) of the population of water molecules. The role of entropy in chemical systems is discussed further in a later section.

The opposition of water molecules to having their motion restricted by forming cages around hydrophobic molecules or portions thereof is the major reason molecules such as tristearin and heptane are essentially insoluble in water and interact mainly with other hydrophobic molecules. Nonpolar molecules can also bond together, albeit weakly, through van der Waals interactions. The net result of the hydrophobic and van der Waals interactions is a very powerful tendency for hydrophobic molecules to interact with one another, and not with water.

Small hydrocarbons like butane (CH3—CH2—CH2—CH3) are somewhat soluble in water, because they can dissolve without disrupting the water lattice appreciably. However, 1-butanol (CH3—CH2—CH2—CH2OH) mixes completely with water in all proportions. The replacement of just one hydrogen atom with the polar —OH group allows the molecule to form hydrogen bonds with water and greatly increases its solubility.

Simply put, like dissolves like. Polar molecules dissolve in polar solvents such as water, while nonpolar molecules dissolve in nonpolar solvents such as hexane.

Multiple Noncovalent Bonds Can Confer Binding Specificity

Besides contributing to the stability of large biological molecules, multiple noncovalent bonds can also confer specificity by determining how large molecules will fold or which regions of different molecules will bind together. All types of these weak interactions are effective only over a short range and require close contact between the reacting groups. For noncovalent bonds to form properly, there must be a complementarity between the sites on the two interacting surfaces. Figure 2-17 illustrates how several different weak bonds can bind two protein chains together. Almost any other arrangement of the same groups on the two surfaces would not allow the molecules to bind so tightly. Such multiple, specific interactions allow protein molecules to fold into a unique three-dimensional shape (Chapter 3) and the two chains of DNA to bind together (Chapter 4).

Figure 2-17. The binding of a hypothetical pair of proteins by two ionic bonds, one hydrogen bond, and one large combination of hydrophobic and van der Waals interactions.

Figure 2-17

The binding of a hypothetical pair of proteins by two ionic bonds, one hydrogen bond, and one large combination of hydrophobic and van der Waals interactions. The structural complementarity of the surfaces of the two molecules gives rise to this particular (more...)

Phospholipids Are Amphipathic Molecules

Multiple noncovalent bonds also are critical in stabilizing the structure of biomembranes, whose primary components are phospholipids. Because the essential properties of biomembranes derive from phospholipids, we first examine the chemistry of these compounds and then see how they associate into the sheetlike structures that are the foundation of biomembranes.

All phospholipids contain one or more acyl chains derived from fatty acids, which consist of a hydrocarbon chain attached to a carboxyl group (—COOH). Fatty acids are insoluble in water and salt solutions; they differ in length and in the extent and position of their double bonds. Table 2-2 lists the principal fatty acids found in cells. Most fatty acids have an even number of carbon atoms, usually 16, 18, or 20.

Table 2-2. Some Typical Fatty Acids Found in Cells.

Table 2-2

Some Typical Fatty Acids Found in Cells.

Fatty acids with no double bonds are said to be saturated; those with at least one double bond are unsaturated. Unsaturated fatty acid chains normally have one double bond, but some have two, three, or four. Two stereoisomeric configurations, cis and trans, are possible around each double bond:

Image ch2e15.jpg

A cis double bond introduces a rigid kink in the otherwise flexible straight chain of a fatty acid (Figure 2-18). In general, the fatty acids in biological systems contain only cis double bonds.

Figure 2-18. The effect of a double bond.

Figure 2-18

The effect of a double bond. Shown are space-filling models and chemical structures of the ionized form of palmitic acid, a saturated fatty acid, and oleic acid, an unsaturated one. In saturated fatty acids, the hydrocarbon chain is linear; the cis double (more...)

Phospholipids consist of two long-chain fatty acyl groups linked (usually by an ester bond) to small, highly hydrophilic groups. Consequently, unlike tristearin, phospholipids do not clump together in droplets but orient themselves in sheets, exposing their hydrophilic ends to the aqueous environment. Molecules in which one end (the “head”) interacts with water and the other end (the “tail”) is hydrophobic are said to be amphipathic (Greek, “tolerant of both”). The tendency of amphipathic molecules to form organized structures spontaneously in water is the key to the structure of cell membranes.

In phosphoglycerides, a principal class of phospholipids, fatty acyl side chains are esterified to two of the three hydroxyl groups in glycerol

Image ch2e16.jpg

but the third hydroxyl group is esterified to phosphate. The simplest phospholipid, phosphatidic acid, contains only these components:

Image ch2e17.jpg
where R1 and R2 are fatty acyl groups.

In most phospholipids, however, the phosphate group is also esterified to a hydroxyl group on another hydrophilic compound. In phosphatidylcholine, for example, choline is attached to the phosphate (Figure 2-19). In other phosphoglycerides, the phosphate group is linked to other molecules, such as ethanolamine, the amino acid serine, or the sugar inositol. The negative charge on the phosphate as well as the charged groups or hydroxyl groups on the alcohol esterified to it interact strongly with water.

Figure 2-19. Phosphatidylcholine, a typical phosphoglyceride, has a hydrophobic tail and a hydrophilic head in which choline is linked to glycerol by phosphate.

Figure 2-19

Phosphatidylcholine, a typical phosphoglyceride, has a hydrophobic tail and a hydrophilic head in which choline is linked to glycerol by phosphate. Either or both of the fatty acyl side chains in a phosphoglyceride may be saturated or unsaturated.

The Phospholipid Bilayer Forms the Basic Structure of All Biomembranes

When a suspension of phospholipids is mechanically dispersed in aqueous solution, they can assume three different forms: micelles, bilayer sheets, and liposomes (Figure 2-20). The type of structure formed by a pure phospholipid or a mixture of phospholipids depends on the length of the fatty acyl chains and their degree of saturation, on the temperature, on the ionic composition of the aqueous medium, and on the mode of dispersal of the phospholipids in the solution. In all three forms, hydrophobic interactions cause the fatty acyl chains to aggregate and exclude water molecules from the “core.” Micelles are rarely formed from natural phosphoglycerides, whose fatty acyl chains generally are too bulky to fit into the interior of a micelle.

Figure 2-20. Cross-sectional views of the three structures that can be formed by mechanically dispersing a suspension of phospholipids in aqueous solutions.

Figure 2-20

Cross-sectional views of the three structures that can be formed by mechanically dispersing a suspension of phospholipids in aqueous solutions. Shown are a spherical micelle with a hydrophobic interior composed entirely of fatty acyl chains; a spherically (more...)

Under suitable conditions, phospholipids of the composition present in cells spontaneously form symmetric sheetlike structures, called phospholipid bilayers, that are two molecules thick. Each phospholipid layer in this lamellar structure is called a leaflet. The hydrocarbon side chains in each leaflet minimize contact with water by aligning themselves tightly together in the center of the bilayer, forming a hydrophobic core that is about 3 nm thick. The close packing of these hydrocarbon side chains is stabilized by van der Waals interactions between them. Ionic and hydrogen bonds stabilize the interaction of the phospholipid polar head groups with each other and with water. At neutral pH, the polar head groups in some phospholipids (e.g., phosphatidylcholine) have no net electric charge, whereas the head groups in others have a net negative charge. Nonetheless, all phospholipids can pack together into the characteristic bilayer structure.

A phospholipid bilayer can be of almost unlimited size — from micrometers (µ) to millimeters (mm) in length or width — and can contain tens of millions of phospholipid molecules. Because of their hydrophobic core, bilayers are impermeable to salts, sugars, and most other small hydrophilic molecules. Like a phospholipid bilayer, all biological membranes have a hydrophobic core, and they all separate two aqueous solutions. The plasma membrane, for example, separates the interior of the cell from its surroundings. Similarly, the membranes that surround the organelles of eukaryotic cells separate one aqueous phase — the cell cytosol — from another — the interior of the organelle. Several types of evidence indicate that the phospholipid bilayer is the basic structural unit of nearly all biomembranes (Chapter 5). Associated with membrane phospholipids are various proteins that help confer unique properties on each type of membrane. We describe the general structure of membrane proteins and their association with the phospholipid bilayer in Chapter 3.


  •  Noncovalent bonds determine the shape of many large biological molecules and stabilize complexes composed of two or more different molecules.
  •  There are four main types of noncovalent bonds in biological systems: hydrogen bonds, ionic bonds, van der Waals interactions, and hydrophobic bonds. The bond energies for these interactions range from about 1 to 5 kcal/mol.
  •  In a hydrogen bond, a hydrogen atom covalently bonded to an electronegative donor atom associates with an acceptor atom whose nonbonding electrons attract the hydrogen (see Figure 2-12). Hydrogen bonds among water molecules are largely responsible for the properties of both liquid water and the crystalline solid form (ice).
  •  Ionic bonds result from the electrostatic attraction between the positive and negative charges of ions. In aqueous solutions, all cations and anions are surrounded by a tightly bound shell of water molecules.
  •  The weak and relatively nonspecific van der Waals interactions are created whenever any two atoms approach each other closely (see Figure 2-15). They result from the attraction between transient dipoles associated with all molecules.
  •  Hydrophobic bonds occur between nonpolar molecules, such as hydrocarbons, in an aqueous environment. Hydrophobic bonds result mainly because aggregation of the hydrophobic molecules necessitates less organization of water into “cages” (and, hence, less reduction in entropy) than if many cages of water molecules had to surround individual hydrophobic molecules.
  •  Although any single noncovalent bond is quite weak, several such bonds between molecules or between the parts of one molecule can stabilize the threedimensional structures of proteins and nucleic acids and mediate specific binding interactions.
  •  Phospholipids, the main components of biomembranes, are amphipathic molecules (see Figure 2-19). Noncovalent bonds are responsible for organizing and stabilizing phospholipids into one of three structures in aqueous solution (see Figure 2-20).
  •  The basic structure of biomembranes consists of a phospholipid bilayer in which the long hydrocarbon fatty acyl side chains in each leaflet are oriented toward one another, forming a hydrophobic core, and the polar head groups line both surfaces. Natural biomembranes also contain proteins, cholesterol, and other components inserted into the phospholipid bilayer.
Image ch2f1a

By agreement with the publisher, this book is accessible by the search feature, but cannot be browsed.

Copyright © 2000, W. H. Freeman and Company.
Bookshelf ID: NBK21726