A very significant aspect of the chemistry of hydrogen is its formation of hydrogen bonds. Whenever hydrogen is covalently bonded to a extremely electronegative element like F, O, and N, the electronegative element attracts the electron pair towards itself giving increase to an induced positive charge ('δ+') on the hydrogen atom or negative charge (δ-) on the electronegative atom for instance,
Hδ+ xδ here x is the electronegative atom. When this happens the hydrogen, due to its positive character, will attract another electronegative atom of the neighboring molecule forming a bond.
This bond is known as hydrogen bond. This is illustrated below
H δ+ X δ-----------H δ+ X δ------------H δ+ Xδ- ------------
v denotes covalent bond.
---------v denotes hydrogen bond.
In this chapter we will be studying the kinds of bonding, the consequence of hydrogen bonding on boiling or melting points and on water solubility of the compound containing hydrogen bonded elements.
Hydrogen bond can be termed as 'the attractive force that binds hydrogen atom of 1 molecule with electronegative atom of another molecule, usually of the same compound'. The hydrogen bond energy is only about 7 - 59 kg mol-1 compared to the usual covalent bond energy of 389 - 665 kg mol-1 for H-N, H-O and H-F bonds. This, hydrogen bond is much weaker than a covalent bond. Evidently, its length is as well much more than the covalent bond. There are 2 kinds of hydrogen bonding. These are as
Intermolecular hydrogen bonding:
In this case, 2 or more molecules of similar are involved in hydrogen bonding. Several common instances of intermolecular hydrogen bonding occurring between the molecules of similar compound are HF, H2O alcohols etc.
Instances of the intermolecular hydrogen bonding between 2 different kinds of molecules are as given.
Intramolecular hydrogen bonding:
Intramolecular hydrogen bond is formed between 2 atoms of similar molecule. As a effect of this, usually a 5-6 membered ring termed chelate ring is formed.
Intramolecular hydrogen bonding takes place in molecules like δ- Nitrophenol and Sallcylaldehyde.
Intramolecular hydrogen bonding does not take place in p- nitrophenol since of the large distance between the 2 groups (NO2 and QH) in p - nitrophenol
This kind of hydrogen bonding is not possible. It does though show the usual intermolecular hydrogen bonding.
It is important to note that the vast majority of intra molecular hydrogen bonding takes place wherever a 5-6 membered ring can be formed since of the stability associated with these rings.
Effects of Hydrogen Bonding:
Hydrogen bonding plays a very important role in determining the properties of compounds. In this division we shall discuss its consequence on the melting point, boiling point or solubility in water.
Boiling Point and Melting Point:
If we study the values for the melting or boiling points shown, we will see that the melting and boiling points of the hydrides of group 14 elements example CH4, SiH4, GeH4 and SnH4 show a common amplify with increase in molecular weight. Though, NH3 in Group 15, HF in Groupl7 and water contain abnormally high melting or boiling points as compared to other hydrides in their relevant groups in the periodic table. This anomaly is explained on the basis of hydrogen bond formation. In compounds anywhere the molecules are linked via hydrogen bonds, several extra energy is required to break the intermolecular hydrogen bond or this is responsible for their superior boiling and melting points Intra molecular hydrogen bond, though has the opposite consequences. For instance, in ortho nitrophenol the groups present in ortho position are involved in intramolecular hydrogen bonding thus preventing the intermolecular hydrogen bond formation, for example association of the molecule. Due to the intramolecular chelated structure, o-nitrophenol is seen to be volatile wherever as p-nitrophenol is not.
Solubility of a substance rises markedly whenever hydrogen bonding is possible between the solvent and the solute molecules. For instance, lower alcohols like, methanol, ethanol etc are highly miscible with water due to the hydrogen bonding with water molecules.
Polarising power of h+:
We familiar that the polarizing power of a cation, for example its ability to distort and polarize an anion is directly proportional to its positive charge or inversely proportional to its size. We can as well say that the polarizing power of a cation is proportional to the ratio of its charge to its size. This ratio is recognized as the ionic potential of the cation. As the hydrogen cation, for example proton is vanishgly small, it has a very high ionic potential and a huge polarizing power.
As consequence of this polarizing power for protons, H+ hardly exists freely. They are usually found related with other molecules. For example, with ammonia and water, these form species like NH4 +, H3O + H5O+, HgO4+etc. The aquated proton species are represented as H+. Enthalpy of formation of these aquated proton species is large (-1075kg mol-1). It is mainly since of this reason that many covalent hydrides (H-X) are acidic in aqueous solution, example they release H+ ions even though H-X bonds are often very strong in them.
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