Covalent & Noncovalent Bonds Stabilize Biologic Molecules

The covalent bond is the strongest force that holds molecules together (Table1). Noncovalent forces, while of lesser magnitude, predominate in stabilizing the folding of the polypeptides and other macromolecules into the complex three-dimensional conformations essential to their functional competence as well as the association of bio molecules into multicomponent complexes. Examples of the latter include the coalescence of the polypeptide subunits that form the hemoglobin tetramer; the association of the two polynucleotide strands that comprise a DNA double helix; and the coalescence of billions of phospholipid, glycosphingolipid, cholesterol, and other molecules into the bilayer that constitutes the foundation of the plasma membrane of an animal cell. These forces, which can be either attractive or repulsive, involve interactions both within the biomolecule and, most importantly, between it and the water that forms the principal component of the sur rounding environment.

Table1. Bond Energies for Atoms of Biologic Significance

In Water, Biomolecules Fold to Position Hydrophobic Groups Within Their Interior

Most biomolecules are amphipathic; that is, they possess regions rich in charged or polar functional groups as well as regions with hydrophobic character. Proteins tend to fold with the R-groups of amino acids with hydrophobic side chains in the interior. Amino acids with charged or polar amino acid side chains (eg, arginine, glutamate, serine) generally are present on the surface in contact with water. A similar pattern prevails in a phospholipid bilayer where the charged “head groups” of phosphatidylserine or phosphatidyl ethanolamine contact water while their hydrophobic fatty acyl side chains cluster together, excluding water. This pattern minimizes energetically unfavorable contacts between water and hydrophobic groups. It also maximizes the opportunities for the formation of energetically favorable charge-dipole, dipole-dipole, and hydrogen bonding interactions between polar groups on the biomolecule and water.

Hydrophobic Interactions

Hydrophobic interaction refers to the tendency of nonpolar compounds to self-associate in an aqueous environment. This self-association is driven neither by mutual attraction nor by what are sometimes incorrectly referred to as “hydrophobic bonds.” Self-association minimizes the disruption of energetically favorable interactions between and is therefore driven by the surrounding water molecules.

While the hydrogen atoms of nonpolar groups such as the methylene groups of hydrocarbons do not form hydrogen bonds, they do affect the structure of the water with which they are in contact. Water molecules adjacent to a hydrophobic group are restricted in the number of orientations (degrees of freedom) that permit them to participate in the maximum number of energetically favorable hydrogen bonds. Maximal formation of multiple hydrogen bonds, which maximizes enthalpy, can be maintained only by increasing the order of the adjacent water molecules, with an accompanying decrease in entropy.

It follows from the second law of thermodynamics that the optimal free energy of a hydrocarbon-water mixture is a function of both maximal enthalpy (from hydrogen bonding) and highest entropy (maximum degrees of freedom). Thus, nonpolar molecules tend to form droplets that minimize exposed surface area and reduce the number of water molecules whose motional freedom becomes restricted (Figure1). Similarly, in the aqueous environment of the living cell the hydrophobic portions of amphipathic biopolymers tend to be buried inside the structure of the molecule, or within a lipid bilayer, minimizing contact with water.

Fig1. Hydrophobic interactions are driven by the surrounding water molecules. Water molecules are represented by one red (oxygen) and two blue (hydrogen) circles. The hydrophobic surfaces of solute molecules are colored gray and, where present, hydrophilic ones are colored green. A. When the six hydrophobic cubes shown are dispersed in water (left), the surrounding water molecules (red oxygens and blue hydrogens) are forced to engage in entropically unfavorable interactions with all 36 faces of the cubes. However, when the six hydrophobic cubes aggregate together (right), the number of exposed faces is reduced to 22. The aggregate forms and its stability is maintained, not by some attractive force, but because aggregation reduces the number of water molecules that are unfavorably affected by nearly 40%. B. Amphipathic molecules associate together for the same reason. However, the structure of the resulting complex (eg, micelle or bilayer) is determined by the geometries of the hydrophobic (gray) and hydrophilic (green) regions.

Electrostatic Interactions

 Electrostatic interactions between oppositely charged groups within or between biomolecules are termed salt bridges. Salt bridges are comparable in strength to hydrogen bonds but act over larger distances. They therefore often facilitate the binding of charged molecules and ions to proteins and nucleic acids.

van der Waals Forces

van der Waals forces arise from attractions between transient dipoles generated by the rapid movement of electrons in all neutral atoms. Significantly weaker than hydrogen bonds but potentially extremely numerous, van der Waals forces decrease as the sixth power of the distance separating atoms (Figure 2). Thus, they act over very short distances, typically 2 to 4 Å.

Fig2. The strength of van der Waals interactions varies with the distance,R, between interacting species. The force of interaction between interacting species increases with decreasing distance between them until they are separated by the van der Waals contact distance (see arrow marked A). Repulsion due to interaction between the electron clouds of each atom or molecule then supervenes. While individual van der Waals interactions are extremely weak, their cumulative effect is nevertheless substantial for macro molecules such as DNA and proteins which have many atoms in close contact.

Multiple Forces Stabilize Biomolecules

The DNA double helix illustrates the contribution of multiple forces to the structure of biomolecules. While each individual DNA strand is held together by covalent bonds, the two strands of the helix are held together exclusively by noncovalent interactions such as hydrogen bonds between nucleotide bases (Watson-Crick base pairing) and van der Waals interactions between the stacked purine and pyrimidine bases. The double helix presents the charged phosphate groups and polar hydroxyl groups from the ribose sugars of the DNA backbone to water while burying the relatively hydrophobic nucleotide bases inside. The extended backbone maximizes the distance between negatively charged phosphates, minimizing unfavorable electrostatic interactions.

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