Chemical Bonding

In order to maintain their existence and carry out the complicated sets of reactions they require, biological systems need to establish and maintain molecular integrity, to maintain higher-order structures, and to bring together multimolecular assemblages of varying specificity and duration.  In order to meet these different requirements, organisms use every available type of interatomic interaction: covalent bonding, and the full range of noncovalent charge interactions.  We will concentrate on five “bond” types:

·         Covalent bonds

·         Ionic bonds

·         Dipole interactions

·         Hydrogen bonds

·         Hydrophobic effects (we’ll discuss in conjunction with water chemistry)


Covalent Bonding

The nature of the covalent bond is covered in General Chemistry; we will touch on it only briefly.  Covalent bonds result from the rearranged orbitals which the bonding electrons assume around the nuclei of the bonded atoms.  The strength of the interaction results from the attraction of the positively charged nuclei to this joined electron “cloud” and is normally quite high.  Covalent bond energies are normally in the range of 25 – 110 kcal/mol  ( 100 – 475 kJ/mol ).


Bond Number*














*the bond numbers illustrated are typical of biological systems and should not be considered set in stone!


The strength of the covalent bond depends on a number of factors; for example, as previously discussed, repulsion between lone electron pairs on adjacent nitrogen atoms results in a relatively weak ( 25 kcal/mol ) N-N bond.  As we’ll discuss soon, the common currency for energy transfer within a cell is the cleavage of a triphosphate to orthophosphate and pyrophosphate.  The electrostatic repulsion that exists between negatively charged oxygen atoms in the phosphate chain destabilizes the bonds between them and makes them readily subject to hydrolysis.  Commonly referred to in textbooks as “high – energy” phosphate bonds, these bonds are actually rather low – energy structures.  If they were really high – energy bonds, they would be difficult to break and would not be key participants in so many reactions!  The “high – energy” misnomer comes from the relatively large amount of free energy recovered from their hydrolysis.

Ionic Bonds:

Ionic bonds, resulting from the attractive force between oppositely charged atomic and / or molecular ions, are just one ( strong ) extreme of a series of electrostatic interactions involving complete or partial charges.  The force of an ionic interaction follows Coulomb’s Law:


The constant  k is related to the permittivity of the vacuum, and is equal to:

q1, q2 are the respective charges in coulombs (for 1 electron, q = 1..6 ·10-19 coulombs ).

r is the distance separating the charges in meters.

e is the dielectric constant of the medium, in dimensionless units.


The dielectric constant for a medium is a measure of it’s own proclivity to engage in charge interactions and thus represents a shielding effect on the charges by the medium itself.  The dielectric constant of the vacuum is defined as 1.  The dielectric constant of water is very high, about 80.  Most organic solvents have dielectric constants much less than water.  The interior of a protein is a very poor dielectric, about 4, and charge interactions in such a locale are quite strong.

We are usually more directly concerned with the energy of the interaction than with the force.  In the case of ionic interactions, the “bond” energy is the energy required to completely bring the two charges together from infinite distance to their current position.  We thus get:



It can thus be seen that ionic interactions are fairly long-range effects, falling off as a function of the distance between the charges.  Ionic interactions are also non-directional.  Charge – charge interactions within and between macromolecules are often referred to as salt links, since the association of a positive charge, say an amino group, with a negative charge, say a carboxyl group, mirrors the association of anions and cations into a salt.


The strength of ionic interactions is quite dependent on the medium.  For example, in a vacuum, the strength of the “bond” between, say, a positive amino group and a negative carboxyl group separated by 4 Å is about 80 kcal/mol, equivalent to a fairly strong covalent bond.  This explains the high melting and boiling points of ionic solids.  In the interior of a protein, however, from which water is excluded, the strength of the same interaction is about 20 kcal/mol.  On the fully exposed and solvated surface of a molecule, the same interaction has a strength of only about 1 kcal/mol.


Dipole Interactions (van der Waal Interactions) :


When electrons are shared unevenly between covalently bonded atoms, the result is a dipole – each atom has a partial negative or positive charge.  Unequal electron distribution results from the different relative abilities of different atom types to attract electrons.  This power to attract electrons is measured by the electronegativity scale.


H = 2.1







Li = 1.0

Be = 1.5

B = 2.0

C = 2.5

N = 3.0

O = 3.5

F = 4.0

Na = 0.9

Mg = 1.2

Al = 1.5

Si = 1.8

P = 2.1

S = 2.5

Cl = 3.0

K = 0.8

Ca = 1.0

Ga = 1.6

Ge = 1.8

As = 2.0

Se = 2.4

Br = 2.8

Rb = 0.8

Sr = 1.0

In = 1.7

Sn = 1.8

Sb = 1.9

Te = 2.1

I = 2.5

Cs = 0.7

Ba = 0.9

Tl = 1.8

Pb = 1.9

Bi = 1.9

Po = 2.0

At = 2.2


The more electronegative an element, the greater it’s power to attract electrons.  Electronegativities increase as one goes up the periodic table and to the right.  Fluorine is the most electronegative element (4.0), cesium the least (0.7)  The greater the difference in electronegativities in two bonded atoms, the greater the inequality of electron distribution between them.  For example, the electronegativity of oxygen is 3.45; that of carbon is 2.55, indicating that oxygen has a higher affinity for electrons than carbon.  The carbon – oxygen bond will therefore form a dipole, with a partial negative charge on O and a partial positive charge on C.  The magnitude of the dipole depends on

            (a)        the magnitude of the charge difference (q), and

            (b)       the separation distance of the partial charges.


That is:                       m = q ´ x


In addition to permanent dipoles, as mentioned above, the presence of a complete or partial charge in close proximity to an otherwise non-polar molecule may alter the electron distribution within the second molecule, producing an induced dipole.  For example, the amino acid glycine normally exists in solution with a negatively charged deprotonated carboxyl group at one end and a positively charged protonated amino group at the other.  If a glycine sits above the plane of a benzene molecule, the p - electron cloud of the benzene ring will be distorted, with the electrons being preferentially drawn to the end of the ring adjacent to the glycine amino group, and preferentially repelled withdrawn from the end adjacent to the glycine carboxyl group.  The resulting induced dipole will interact with the glycine zwitterion.  Induced dipole effects are always attractive.


A special case produces the London dispersion force (frequently referred to by biochemists as “van der Waal forces”, though this term applies to all noncovalent electrostatic interactions.)  Any atom will be found at some given instant of time to be a transient dipole, resulting from the uneven distribution of negative electrons around the positive nucleus at any given moment.  When atoms are in close proximity, there is a tendency for their transient dipoles to flicker in unison.  This produces a weak attractive force between them.  These dispersion forces are individually weak, on the order of 0.1 – 1 kcal/mol.  They are, however, the forces responsible for the liquefaction of noble gases at low temperatures.  Also, in biological macromolecules, the energetic effects of large numbers of such contacts in molecules with complimentary surfaces becomes quite large.

As you might expect, dipoles interact in a manner analogous to ions; the interactions differ both in their strengths ( partial charge interactions are weaker ) and the range over which they operate.  The energy of charge – dipole interactions falls of as 1/r2; the weakest interactions, those between induced dipoles, fall of as 1/r6.


The effective surface of a molecule depends on the relation between the attractive dispersion forces and the repulsive effects of electron cloud overlap which come into play at extremely short distances.  As the distance between charges shrinks, the attractive force becomes stronger.  At very short distances, however, repulsive forces come into play and rapidly make closer approach between the two atoms energetically prohibitive.  Since the repulsive term varies as 1/r12, this “wall of energy” goes very large over very short distances and represents the effective surface of the atom.  The interaction energy at a given radius r is given by the Lennard-Jones equation:



where Uo is the energy minimum – i.e. the strongest possible energy for the interaction - which occurs at the optimum radius ro.  The distance at which energy is minimized is considered the effective radius and is known as the van der Waals radius of the atom.  Van der Waals radii are on the order of 1 to 2 Å

Hydrogen Bonds:


A special type of linkage occurs  when a hydrogen atom, covalently bound to an electronegative atom, makes a partial bond with an electron pair on another nearby electronegative atom.  Such a bond is referred to as a hydrogen bond; hydrogen bonds are very important in aqueous, and especially biological systems.


The group to which the participant hydrogen is covalently bound is the hydrogen bond donor; the group to which the hydrogen bond forms is the hydrogen bond acceptor. 





:N -


:O =


:O -


Some important points regarding hydrogen bonds


·         Hydrogen bonds are partially covalent in character.  The interaction is more than electrostatic, as reflected by the interatomic distance between donor and acceptor atoms, which is considerably less than would be expected from the van der Waals radii, indicating the actual sharing of electrons.


·         Hydrogen bonds are highly directional.  Unlike purely electrostatic interactions, the orientation of the participants is critical to bond strength.  H – bonds are strongest when the donor, the shared hydrogen and the acceptor are collinear.  Bond strengths fall off quickly as the bond is distorted.


·         Hydrogen bond strengths are typically in the range of 2 – 5 kcal/mol.  They are thus of intermediate strength between covalent bonds and van der Waals contacts.


Water – Properties of Biological Import:


Above all, biochemistry is aqueous chemistry.  Most biochemical reactions take place in more – or – less dilute aqueous solution.  As a result, the properties of water are critical for biochemistry.


Water is a tetrahedral molecule, with the oxygen atom central and either a hydrogen or a lone electron pair at each apex of the tetrahedron.   Oxygen is highly electronegative.  As a result, electron density is enhances around the oxygen and attenuated around the hydrogens.  Water is highly polar and, as might be expected, is an excellent hydrogen bond donor and hydrogen bond acceptor.  In the liquid state, each water molecule forms a network of hydrogen bonds with surrounding molecules.  These bonds are constantly breaking and reforming, with half – times on the order of 10-10 seconds.  Time – averaged water molecules are involved in about 3.5 hydrogen bonds. 


This hydrogen bond network is responsible for many of the biochemically and biologically critical characteristics of water.

·         Water has abnormally high melting and boiling points.  Methane, ammonia and even the heavier but structurally similar hydrogen sulfide are all gases at normal earth surface temperatures.  H2S is very similar to water, but the H-S bond length is greater and O has been replaced by the much less electronegative S (2.5 vs. 3.5).  As a result, the molecule is a far weaker dipole, does not form hydrogen bonds, and despite it’s greater molecular weight is a gas at room temperature.


·         Water has a very high specific heat ( 1 cal/g. )  It takes a lot of effort to change the temperature of water.  This buffers organisms against environmental  changes, and on a biosphere level helps moderate the planetary climate.  The oceans provide an enormous heat reservoir, and prevent the huge temperature fluctuations that occur on waterless planets.


·         Water has a very high heat of vaporization ( 540 cal/g, compared to 263 cal/g for MeOH and 59 cal/g for CHCl3. )  This permits the use of evaporative cooling and, once again, moderates planetary climate.


·         Water is an excellent solvent for ionic and polar substances.  The high dielectric constant ( 80 ) does an extraordinarily good job of shielding charges from each other.  Compare water to another hydrogen bonder, methane ( e = 33 ), or with ammonia (e = 15.5.) As a result, the strength of interaction between ions in an ionic solid drops to very low values. 


·         Water has a high surface tension.  This is of minor biochemical importance, but biological and physiological effects are significant.  A number of organisms live at the air – water interface on the surface.


In the frozen state, each molecule forms four hydrogen bonds ( 2 donated, 2 accepted .)  Molecular rotations permit a fair degree of freedom in the spatial arrangement of the molecules, and crystalline ice is rather disordered with a fair amount of empty space in the crystal.  This results in another often overlooked, but critical property of water.


·         The density of water decreases on freezing.  The increase in hydrogen bonding increases the space between individual molecules ( locking the O – H – O distance at 2.76 Å. )  In addition, the loose crystal structure contains empty space.  As a result, there are fewer molecules per unit volume in normal ice than in liquid water.  The upshot of all this is that ice floats.


Were ice more dense than water, bodies of water would freeze from the bottom up, destroying aquatic life.  Instead, water bodies freeze on the surface, leaving an insulating layer between the cold atmosphere and liquid water beneath.


Hydrophobic Interactions:


A corollary of the solvent properties of water for polar compounds is that water s a poor solvent for nonpolar compounds.  This hydrophobic effect is entropically driven.


Water in the liquid state has a high degree of conformational freedom.  When an attempt is made to dissolve a nonpolar compound, the hydrogen bond network is disrupted.  In order to solvate the molecule, water molecules form a ( relatively )highly – ordered cage   In bulk water, the intermolecular forces on an individual water molecule are isotropic – they come equally from all sides; at the interface between a nonpolar molecule and water, the forces are anisotropic.  The water molecules are constrained both translationally and rotationally.  These constraints can be minimized by minimizing the contact between water and the nonpolar substance.  Minimization occurs when the nonpolar material assumes a shape with the minimum surface area.  Depending on the conditions, this will be a sphere or a separate layer.  Biological molecules generally fold in such a way as to minimize contact between nonpolar groups and the aqueous solvent.  The interiors of proteins are generally characterized by close contacts among hydrophobic groups; while the nonpolar interactions among aliphatic chains of fatty acids forms the basic structure of all biological membranes.