1.3.3 Hydrogen Bonding

where k = 9.0 x 10 9 nt-meter 2 /coul 2 q = -1.6 x 10 -19 coulombs for an electron. r = distance between the point charges (meters) ε = the dielectric constant of the medium (unit less).

ε reflects the tendency of the medium to shield one charge from another. ε is 1 in a vacuum, around 4 in the interior of a protein and 80 in water. The problem of calculating electrostatic effects in proteins is complex in part because of non-uniformity of the dielectric environment. The dielectric micro-environment is variable, with less shielding of charges in regions of hydrocarbon sidechains and greater shielding in regions of polar sidechains.

The electrostatic energy is given by: ΔE= k a q1 q2 / ε r where a = avogadro's number.

One can crudely estimate the energetics of a charge-charge interaction in a protein. The energy of an amine (charge +1) and a carboxylic acid (charge -1) separated by 4 Å in the interior of protein is given by:

ΔE = -(9.0x109nt-m2/coul2)(6.02x1023)(1.6x10-19coul)2 /4(4x10-10m) = 87 kjoules/mole = 21 kcal/mole

This rough approximation is around 10-fold greater than the values determined experimentally.

Charge-charge interactions fall off slowly with distance (1/r).

1.3.3 Hydrogen Bonding

Hydrogen bonding is an intermolecular or intramolecular attraction that occurs between molecules with hydrogen bond donors and molecules with hydrogen bond acceptors. Hydrogen bond donors are molecules that have a hydrogen attached to an electronegative atom (for example, hydroxyls or amines). Hydrogen bond acceptors are molecules that have a lone pair of electrons located on an electronegative atom (for example, oxygen, nitrogen, or fluorine). Hydrogen bonds are not as strong as covalent and ionic bonds but are stronger than van der Waals interactions. Hydrogen bonding is responsible for the high boiling point of water and is important for the organization of complementary chains of base pairs in DNA and RNA.

 A hydrogen atom attached to a relatively electronegative atom is a hydrogen bond donor. This electronegative atom is usually fluorine, oxygen, or nitrogen.  An electronegative atom such as fluorine, oxygen, or nitrogen is a hydrogen bond acceptor, whether it is bonded to a hydrogen atom or not.  An example of a hydrogen bond donor is ethanol, which has a hydrogen bonded to oxygen; an example of a hydrogen bond acceptor which does not have a hydrogen atom bonded to it is the oxygen atom on diethyl ether.  Examples of hydrogen bond donating (donors) and hydrogen bond accepting groups (acceptors)  Cyclic dimer of acetic acid; dashed green lines represent hydrogen bonds  A hydrogen attached to carbon can also participate in hydrogen bonding when the carbon atom is bound to electronegative atoms, as is the case in chloroform, CHCl3. The electronegative atom attracts the electron cloud from around the hydrogen nucleus and, by decentralizing the cloud, leaves the atom with a positive partial charge.  Because of the small size of hydrogen relative to other atoms and molecules, the resulting charge, though only partial, represents a large charge density. A hydrogen bond results when this strong positive charge density attracts a lone pair of electrons on another heteroatom, which becomes the hydrogen-bond Acceptor.

Figure 1. 26 Interaction between +vely charged Mg+2 and negative DNA molecles

Unit 1  The hydrogen bond is often described as an electrostatic dipole-dipole interaction. It also has some features of covalent bonding: it is directional and strong, produces interatomic distances shorter than sum of van der Waals radii, and usually involves a limited number of interaction partners, which can be interpreted as a type of valence. These covalent features are more substantial when acceptors bind hydrogen from more electronegative donors.  Liquids that display hydrogen bonding are called associated liquids.  Hydrogen bonds can vary in strength from very weak (1–2 kJ mol−1) to extremely strong (161.5 kJ mol−1 in the ion HF−2) Typical enthalpies in vapour include: 1. F−H…: F (161.5 kJ/mol or 38.6 kcal/mol) 2. O−H…: N (29 kJ/mol or 6.9 kcal/mol) 3. O−H…: O (21 kJ/mol or 5.0 kcal/mol) 4. N−H…: N (13 kJ/mol or 3.1 kcal/mol) 5. N−H…: O (8 kJ/mol or 1.9 kcal/mol) 6. HO−H…:OH+3 (18 kJ/mol[10] or 4.3 kcal/mol Hydrogen bond is most obvious in a water molecule  There are two hydrogen atoms and one oxygen atom. Two molecules of water can form a hydrogen bond between them; the simplest case, when only two molecules are present, is called the water dimer and is often used as a model system.  When more molecules are present, as is the case with liquid water, more bonds are possible because the oxygen of one water molecule has two lone pairs of electrons, each of which can form a hydrogen bond with a hydrogen on another water molecule.  This can repeat such that every water molecule is H-bonded with up to four other molecules, as shown in the figure (two through its two lone pairs, and two through its two hydrogen atoms). Hydrogen bonding strongly affects the crystal structure of ice, helping to create an open hexagonal lattice.  The density of ice is less than the density of water at the same temperature; thus, the solid phase of water floats on the liquid, unlike most other substances.  Liquid water's high boiling point is due to the high number of hydrogen bonds each molecule can form, relative to its low molecular mass.  Due to difficulty of breaking these bonds, water has a very high boiling point, melting point, and viscosity compared to otherwise similar liquids not conjoined by hydrogen bonds.  Water is unique because its oxygen atom has two lone pairs and two hydrogen atoms, meaning that the total number of bonds of a water molecule is up to four.  For example, hydrogen fluoride—which has three lone pairs on the F atom but only one H atom— can form only two bonds; (ammonia has the opposite problem: three hydrogen atoms but only one lone pair). H−F…H−F…H−F  Because water forms hydrogen bonds with the donors and acceptors on solutes dissolved within it, it inhibits the formation of a hydrogen bond between two molecules of those solutes or the formation of intramolecular hydrogen bonds within those solutes through competition for their donors and acceptors.  Consequently, hydrogen bonds between or within solute molecules dissolved in water are almost always unfavourable relative to hydrogen bonds between water and the donors and acceptors for hydrogen bonds on those solutes