HYDROGEN BONDING

Introduction:

An atom of hydrogen linked covalently to a strongly electronegative atom can establish an extra weak attachment to another electronegative atom in the same or different molecules. This attachment is called a hydrogen bond. To distinguish from a normal covalent bond, a hydrogen bond is represented by a broken line eg X – H^…Y where X & Y are two electronegative atoms. The strength of hydrogen bond is quite low about 2-10 kcal mol^–1 or 8.4–42 kJ mol^–1 as compared to a covalent bond strength 50–100 kcal mol^–1 or 209 –419 kJ mol^–1

Conditions for Hydrogen Bonding:

$\bullet$Hydrogen should be linked to a highly electronegative element.

$\bullet$The size of the electronegative element must be small.

These two criteria are fulfilled by F, O, and N in the periodic table. Greater the electronegativity and smaller the size, the stronger is the hydrogen bond which is evident from the relative order of energies of hydrogen bonds.

Types of Hydrogen Bonding:

$\bullet$Intermolecular hydrogen bonding: This type of bonding takes place between two molecules of the same or different types. For example,

Intermolecular hydrogen bonding leads to molecular association in liquids like water etc. Thus in water only a few percent of the water molecules appear not to be hydrogen bonded even at 90°C. Breaking of those hydrogen bonds throughout the entire liquid requires appreciable heat energy. This is indicated in the relatively higher boiling points of hydrogen bonded liquids. Crystalline hydrogen fluoride consists of the polymer (HF)_n. This has a zig-zag chain structure involving

H-bond.

$\bullet$Intramolecular hydrogen bonding: This type of bonding occurs between atoms of the same molecule present on different sites. Intramolecular hydrogen bonding gives rise to a closed ring structure for which the term chelation is sometimes used. Examples are

o-nitrophenol, salicylaldehyde.

Effect of Hydrogen Bonding

Hydrogen bonding has got a very pronounced effects on certain properties of the molecules. They have got effects on

$\bullet$State of the substance

$\bullet$Solubility of the substance

$\bullet$Boiling point

$\bullet$Acidity of different isomers

These can be evident from the following examples.

Example 1. H₂O is a liquid at ordinary temperature while H₂S is a gas although both O and S belong to the same group of the periodic table.

Solution: H₂O is capable of forming intermolecular hydrogen bonds. This is possible due to high electronegativity and small size of oxygen. Due to intermolecular H-bonding, molecular association takes place. As a result the effective molecular weight increases and hence the boiling point increases. So H₂O is a liquid. But in H₂S no hydrogen bonding is possible due to large size and less electronegativity of S. So it's boiling point is equal to that of an isolated H₂S molecule and therefore it is a gas.

Importance of Hydrogen Bonding in Biological Systems:

Hydrogen bonding plays a vital role in physiological systems. Proteins contain chains of amino acids. The amino acid units are arranged in a spiral form somewhat like a stretched coil spring (forming a helix). The N-H group of each amino acid unit and the fourth C=O group following it along the chain, establishes the N–H---O hydrogen bonds. These bonds are partly responsible for the stability of the spiral structure. Double helix structure of DNA also consists of two strands forming a double helix and are joined to each other through hydrogen bond.

INTERMOLECULAR FORCES

We have enough reasons to believe that a net attractive force operates between molecules of a gas. Though weak in nature, this force is ultimately responsible for liquifaction and solidification of gases. But we cannot explain the nature of this force from the ideas of ionic and covalent bond developed so far, particularly when we think of saturated molecules like H₂, CH₄, He etc. The existence of intermolecular attraction in gases was first recognised by Vanderwaal's and accordingly intermolecular forces have been termed as Vanderwaal's forces. It has been established that such forces are also present in the solid and liquid states of many substances. Collectively they have also been termed London forces since their nature was first explained by London using wave mechanics.

Nature of Vanderwaal's Forces:

The Vanderwaal's forces are very weak in comparison to other chemical forces. In solid NH₃ it amount to about 39 KJ mol^–1 (bond energy N-H bond = 389 KJ mol^–1). The forces are non directional. The strength of Vanderwaal's force increases as the size of the units linked increases. When other factors (like H-bonding is absent), this can be appreciated by comparison of the melting or boiling points of similar compounds in a group.

Origin of Intermolecular Forces:

Intermolecular forces may have a wide variety of origin.

$\bullet$Dipole-dipole interaction: This force would exist in any molecule having a permanent dipole e.g. HF, HCl, H₂O etc.

$\bullet$ Ion-dipole interaction: These interactions are operative in solvation and dissolution of ionic compounds in polar solvents.

$\bullet$ Induced dipole interaction: These generate from the polarisation of a neutral molecule by a charge or ion.

$\bullet$ Instantaneous dipole-induced dipole interaction: In non polar molecules dipoles may generate due to temporary fluctuations in electron density. These transient dipole can now induce dipole in neighbouring molecules producing a weak temporary interaction.

METALLIC BONDING

Metals are characterised by bright, lustre, high electrical and thermal conductivity, malleability, ductility and high tensile strength. A metallic crystal consists of very large number of atoms arranged in a regular pattern. Different model have been proposed to explain the nature of metallic bonding two most important modules are as follows

The electron sea Model

In this model a metal is assumed to consist of a lattice of positive ion (or kernels) immersed in a sea of mobile valence electrons, which move freely within the boundaries of a crystal. A positive kernel consists of the nucleus of the atom together with its core on a kernel is, therefore, equal in magnitude to the total valence electronic charge per atom. The free electrons shield the positively charged ion cores from mutual electrostatic repulsive forces which they would otherwise exert upon one another. In a way these free electrons act as 'glue' to hold the ion cores together.

The forces that hold the atoms together in a metal as a result of the attraction between positive ions and surrounding freely mobile electrons are known as metallic bonds.

Through the electron sea predated quantum mechanics it still satisfactorily explains certain properties of the metals. The electrical and thermal conductivity of metals for example, can be explained by the presence of mobile electrons in metals. On applying an electron field, these mobile electrons conduct electricity throughout the metals from one end to other. Similarly, if one part of metal is heated, the mobile electrons in the part of the metals acquire a large amount of kinetic energy. Being free and mobile, these electrons move rapidly throughout the metal and conduct heat to the other part of the metal.

On the whole this model is not satisfactory.

Example 2. Sodium metal conducts electricity because it

(A) is a soft metal

(B) contains only one valence electron

(C) has mobile electron

(D) reacts with water to form H₂ gas

Solution: (C)