Bai 2 1.1 2007/10/10 09:55:03.840 GMT-5 2007/10/10 10:03:31.265 GMT-5 Thang Tien Nguyen hungtran@vef.gov Thang Tien Nguyen hungtran@vef.gov Covalent vs Non-Covalent Bonds

Covalent bonds are the forces that hold atoms together as molecules. For example, the two O-H bonds in water molecules are covalent bonds. The covalent bonds most important in biology (C-C and C-H) have bond energies in the range of 300-400 kJ/mol. Non-covalent interactions (also called noncovalent forces or noncovalent bonds) are weak interactions between ions, molecules, and parts of molecules. They help shape individual molecules and groups of molecules and ions, but are weak enough to be continually broken and re-formed in the dynamic molecular interplay that is life. In fact, Figure 2.1 shows that biologically important noncovalent interactions are 10-100 times weaker than covalent bonds. Figure 2.2: Types of noncovalent interactions.

The different types of noncovalent interactions are summarized in Figure 2.2. All are fundamentally electrostatic in nature; that is, all depend on the forces that electrical charges exert on one another. Note that all but hydrogen bonds decrease with distance. For cells, both covalent bonds and noncovalent interactions are important. DNA, for example, is composed of two intertwined chains of polynucleotides. The forces that hold together the atoms of the nucleotides in each individual chain are covalent. The forces that hold the two chains together, however, are noncovalent hydrogen bonds. The weaker hydrogen bond forces are strong enough to keep the two chains together, but weak enough to enable the cell to pull the chains apart when necessary to perform DNA replication (see here) or other processes that require the chains to be separated. Dielectric Constant

Coulomb's law, F = k*(q1q2)/r2 describes the force between two charges (q1 and q2) separated by a distance r in a vacuum (k is a constant). When q1 and q2 are both positive or both negative, the force is positive and repulsive. When one charge is positive and the other is negative, the force is negative and attractive. In biological systems, charges are separated by water, other molecules, or parts of molecules, not a vacuum. Thus, the cellular medium shields charges from each other. To account for solvent shielding, a dimensionless number called the dielectric constant,

, is inserted into the Coulomb equation, F = k*(q1q2)/(

*r2). The dielectric constant is high for a polar solvent and low for nonpolar organic solvents. Hydrogen Bonds

The hydrogen bond is an interaction between a covalently bonded hydrogen atom on a donor group (i.e., -OH or

) and a pair of nonbonded electrons on an acceptor group (i.e.,

or

) as shown in Figure 2.7. The atom to which hydrogen is covalently bonded is called the hydrogen bond donor and the atom with the nonbonded electron pair is called the hydrogen bond acceptor. Hydrogen bond donors tend to be highly electronegative atoms, such as N and O, because they can withdraw negative charge from the hydrogen atom, thus making it partially positive and, thus, more strongly attracted to the electron pair of the hydrogen bond acceptor. Figure 2.7: The hydrogen bond.

 

Hydrogen bonds have characteristics of covalent and noncovalent bonds. For example, the attraction between the partially positive hydrogen and the negative charge of the electron pair is like a charge-charge interaction (Figure 2.2a). At the same time, there is electron sharing (as in a covalent bond) between the hydrogen atom and the hydrogen bond acceptor. Other features that suggest hydrogen bonds are partially covalent in character include the following: - Hydrogen bond lengths are fixed at about 0.33 nm (See Table 2.3) - Hydrogen bonds are highly directional - the donor H bond tends to point directly at the acceptor electron pair. - The energy of hydrogen bonds is greater than most other noncovalent interactions. Hydrogen bonds provide forces that help stabilize the structures of macromolecules, such as DNA and proteins, give a molecule like water its unusual chemical characteristics for its size (see Table 2.4, Table 2.5), and the hydrogen bonds of water also assist in solubilizing polar compounds.