Hydrogen Bonds

 

The electron configuration of H atom is 1s¹. It precedes the nearest inert gas, He (1s²). Therefore, it normally shows monovalency, either electrovalency or covalency. There are many well-known examples in which a hydrogen atom forms two bonds with two highly electronegative atoms from the group F, O, N etc. One of the bonds is covalent and the other is thought to be mainly electrostatic; these two atoms maybe identical or different. Actually, a combination of intermolecular forces such as electrostatic, covanlency, London force etc. gives rise to a hydrogen bond.

Thus in a molecule of hydrogen fluoride, the hydrogen atoms forms a covalent bond with fluorine atom. This H atom also links another F atom of the adjacent HF molecule by a comparatively weak electrostatic force and is represented by a dotted line.

 In this case the H atom serves as a bridge between the two electronegative atoms. For this reason the weak electrostatic bond between the H atom of one molecule and the F atom of the other molecule is called hydrogen bond or H-bond.

Thus H-Bond may be defined as a week electrostatic chemical bond which is found to act between one covalently bonded H atom of a species and a highly electronegative atom like F, O, N, Cl, S of another species.

When a hydrogen atom is covalently linked to a strong electronegative element, the molecule formed is a dipole. Thus H-F is a dipole. In an aggregate of H-F molecules, dipole-dipole interactions take place when they come very close to each other. The obvious result is the orientation of poles and a weak bond is established between two unlike poles. So, H-bond is a result of dipole-dipole interactions.

A hydrogen atom, covalently bonded to a strong electronegative atom, has relatively high positive charge density but minimal coordination number (only two) because it has a very small volume which makes H-bonding possible; while other electropositive atoms have neither high positive charge density norm minimal coordination number because of their large volume. The hydrogen bond dissociation energy varies from 21 to 42 KJ/mole. Hence it is stronger than usual dipole-dipole interactions. The strength of a bond depends upon the electronegativity of the atom bonded to the H atom; it decreases with the decreasing electronegativity within the series F, O, N, Cl, S.

The hydrogen bond length is found to be within the limit 0.12 nm to 0.34 nm; this is always less than the sum of van der Waals radii of the atoms involved, i.e., H atom and another electronegative atom. The hydrogen bond length is the longest in H-S bond and it is the shortest in H-F bond.

Hydrogen bonds are of two types

1) Intermolecular hydrogen bond

2) intramolecular or internal hydrogen bond

When hydrogen bonding happens to occur between two or more molecules, it is known as intermolecular hydrogen bond. It gives rise to association of molecules and species.

When hydrogen bond is formed between two atoms within a molecule, it is called and internal or intramolecular hydrogen bond. The result is the formation of a ring known as chelation.

  Intramolecular hydrogen bond can be found in salicyldehyde, ortho-nitro-phenol, etc. Intermolecular hydrogen bonding can also form ring. In this case, we get a dimeric molecule instead of a polymer, e.g., two molecules of acidic acid associate with each other to form a ring.

Effect of hydrogen bonding force

Effect on melting and boiling points: Owning to the association of identical molecules by intermolecular hydrogen bond, the intermolecular forces amongst the molecules increase to a relatively greater extent. This has a mark effect on melting and boiling points of the compounds. An extra amount of heat energy is required to break the hydrogen bond. Hence, compounds with intermolecular hydrogen bond have greater melting and boiling points than those of compounds without intermolecular hydrogen bonds, provided they have comparable molecular weights.

CH₄, HF and H₂O molecules have compatible molecular weights. CH₄ cannot form hydrogen bond, because carbon atom has low electronegativity; HF can utilize its loan hydrogen atom to form hydrogen bond but a water molecule forms hydrogen bond with both the hydrogen atoms. Consequently, the melting point and boiling point orders of the state compounds are as follow

M. P. order    H₂O = 0^oC          HF = -110^oC         CH₄ = -190^oC

B.P. order      H₂O = 100^oC     HF = 19.5^oC         CH₄ = -161.5^oC

On the other hand, intramolecular hydrogen bonding in a compound prevents intermolecular hydrogen bonding and thus prevents association which would raise melting point and boiling point. Therefore, two compounds of comparable molecular weights, one with intermolecular hydrogen bond and the other with intramolecular hydrogen bond, should have different melting and boiling points. The former compound will possess higher melting and boiling point then those of latter compound. This difference in boiling point often helps separation of two compounds from their mixture by steam distillation. The compound possessing intramolecular hydrogen bonding being of lower boiling point, this distils out with steam when subjected to steam distillation but the other remains in the distillation flask.

Effect on solubility: Hydrogen bond in a compound makes its soluble in water and in other solvent possessing hydrogen bonds. The hydrogen bond in solvent molecules and that in solute molecules break on dissolution. New hydrogen bonding occurs between the solvent and the solute molecules because the energy gained by this new bond formation is more than that lost by the bond breaking processes.

Effect on stability of molecules: Intramolecular hydrogen bonding brings stability to a compound. Acetylacetone molecules in liquid state remain in two tautomeric forms keto and enol.

Owing to the intramolecular hydrogen bond and conjugation in the annual form, it is more stable than the keto form in the liquid acetylacetone and in an n-hexane solution and their percentage is 80% and 20% respectively. In the aqueous medium, however, in all form is 16% and keto form is 84%.

Effect on acidity and basicity: Hydrogen bonding not only brings stability to molecules but also it does so to ions. Therefore, stability of an acid and its conjugate base or that of a base and its conjugate acid can be compared if hydrogen bonding places any role in the species concerned. Again, we know that if a conjugate base is more stable than the parent acid, the base will be relatively weaker than the acid. If the order of their relative stability is a reverse one, then the order of their relative strength will change accordingly. The relation between stability and strength of a base and its conjugate acid follows the same principle as stated already.

Thus dimethylamine is a stronger base than trimethylamine in aqueous medium as the conjugate acid of the former base is more stable than that of the latter compound. The stability order is so because of the greater hydration effect on dimethylammonium ion then that on trimethylammonium ion. This is revealed from the structures of hydrated dimethylammonium ion and that of trimethylammonium ion.

Similarly, the conjugate base of salicylic acid get stabilized by the intra ionic hydrogen bonding, but p- hydroxybenzoate ion, the conjugate base of the p-hydroxybenzoic acid, does not stabilize by such a hydrogen bonding. Hence, we can easily conclude that salicylate ion is a weaker base than p- hydroxybenzoate ion. This means: salicylic acid is a stronger asset than p-hydroxybenzoic acid.

Ion assisted hydrogen bonding: Ion-assisted hydrogen bonding refers to a situation where the presence of ions (charged atoms or molecules) in a solution or material enhances the strength or stability of hydrogen bonds between other molecules. In this context, the ions act as mediators or facilitators for the formation or stabilization of hydrogen bonds. The mechanism of ion-assisted hydrogen bonding can vary depending on the specific system under consideration. Generally, ion-assisted hydrogen bonding can exhibit a range of energies similar to conventional hydrogen bonding, which typically ranges from about 10-40 kJ/mol. However, there are a few general ways in which ions can influence hydrogen bonding:

Electrostatic Interaction: Ions in solution can interact electrostatically with polar molecules, such as those containing hydrogen bond donors and acceptors (e.g., water molecules, alcohol molecules, or biomolecules). The presence of ions can alter the distribution of charges in these molecules, leading to enhanced hydrogen bonding interactions.

Polarization Effects: Ions can induce polarization in nearby molecules, thereby increasing the strength of hydrogen bonds. For example, positively charged ions (cations) can attract the electron clouds of neighbouring molecules, effectively strengthening hydrogen bonds by enhancing the dipole-dipole interactions.

Hydration Shell Formation: In aqueous solutions, ions are typically surrounded by a hydration shell consisting of water molecules. These water molecules can participate in hydrogen bonding interactions with other molecules in the solution. The presence of hydrated ions can increase the overall density of hydrogen bond donors and acceptors in the solution, leading to enhanced hydrogen bonding.

Specific Ion Effects: Some ions exhibit specific interactions with certain functional groups, favouring the formation of hydrogen bonds. For example, ions with hydrogen bond acceptor or donor groups in their structures can directly participate in hydrogen bonding interactions with other molecules.

Hydrogen bonding in biological system: Both inter and intramolecular hydrogen bonding play important role in biological system. The long helical structure of protein molecules and double helical structure of DNA molecules are stabilized by the existence of hydrogen bonding, e.g., a protein molecule contains a large number of -C=O and N-H groups, and hydrogen bonding occurs between these two groups. For example, hydrogen bonding between amino acid residues stabilizes the secondary and tertiary structures of proteins, while hydrogen bonding between complementary base pairs (A-T and G-C) holds the double helix structure of DNA together. A single protein contains large number of this hydrogen bond. All the each of them are very weak (the dissociation energy around 21 KJ/mole), yet the overall effect of the hydrogen bonds is highly significant and bring stability to the molecule of protein. The hydrogen bonding force in protein and cellulose make them able to absorb dyes.

Hydrogen bonding in water: Water molecules consist of one oxygen atom covalently bonded to two hydrogen atoms. Because oxygen is more electronegative, it attracts electrons more strongly, leading to a partial negative charge on the oxygen and partial positive charges on the hydrogen atoms. These partial charges allow water molecules to form hydrogen bonds with neighbouring water molecules. These bonds are electrostatic attractions between the positively charged hydrogen atoms of one molecule and the negatively charged oxygen atoms of another.

Here are some important aspects of hydrogen bonds in water:

Directionality: Hydrogen bonds have a specific orientation, aligning with the angle of the water molecule.

Strength: While not as strong as covalent bonds, hydrogen bonds significantly influence the physical properties of water, being stronger than van der Waals interactions.

Cohesion and Surface Tension: Hydrogen bonds create cohesion among water molecules, contributing to surface tension.

Adhesion: Water molecules adhere to other polar substances due to hydrogen bonding.

High Boiling and Melting Points: Hydrogen bonds require more energy to break, leading to higher boiling and melting points compared to similar-sized molecules without hydrogen bonding.

Density Anomalies: Hydrogen bonding affects water density, with maximum density occurring at around 4 degrees Celsius, unlike most substances.

In summary, hydrogen bonds in water are vital for its cohesive and adhesive properties, high surface tension, and unusual density and temperature behaviours.