Chapter 3: – Chemical Bonding 1st year Book
(Page 10)

Hydrogen Bonding

Latimer and Radebush 1920 introduced the concept of hydrogen bonding. According to them, it is formed when a slightly acidic hydrogen atom already bonded to a strongly electronegative atom such as F, N, or O is linked with weak electrostatic force by the n bonded pair of electrons of another atom.

Cause of hydrogen bond:

When hydrogen is covalently linked with most electronegative atoms like N, O, or F times Cl also) then the shared pair of electrons is attracted by the most electronic: atom due to which a partial negative charge is developed on it and a partial positive charge on the H atom as shown in HF, NH3, H2O, etc.
hydrogen bonding
Such type of development of opposite charges form the molecules polar and thus the molecules behave as dipoles. When a large number of such polar molecules break; together, the positive pole of one molecule attracts the negative pole of other molecules.

H2O molecules hydrogen bonding

Hence, the hydrogen bond is defined as the force of attraction between the positive pole of the H-atom of one molecule and the negative pole of the electronegative atom of the other molecule generally of the same substance is called a “Hydrogen bond”. The hydrogen bond is represented as dotted lines (……) and it forms a bridge between two highly electronegative atoms.
+δ−xδ+δ+δH Bridged H-atom 
In the case of Cl, Br, and I which are less electronegative as compared to F, the molecules HCl, HBr, and HI, the shared electron pair do not lie so far away from H-atom hence they do not behave as a dipole like H−F. Thus, they do not form hydrogen bonding and therefore are more volatile than HF.

Characteristic properties Hydrogen bonds:

Hydrogen bonds are much weaker than a normal covalent bond but stronger than Van der Waal’s forces. It is evident from their bond energies as shown below:
Forces : Van der Waals force hydrogen bond Covalent bond Ionic bond
Bond energy (kcal mole ) 1.00 3 to 10 50 to 100 100


Hydrogen bond increases the boiling points of liquids due to which H2O is a liquid and H2S, H2Se, and H2Te all are gases at ordinary temperature. The reason for this is that in the case of water, hydrogen bond causes the association of the H2O molecules, and some extra energy is needed to break their bonds.
The strength of the hydrogen bond formed depends on the electronegativity difference between hydrogen and electronegative atom. The higher the electronegativity of the electronegative atom greater will be its polarity with a hydrogen atom and the more easily this polar molecule forms a hydrogen bond. Thus the order of electronegativity is F>O>N>Cl.
If the hydrogen is substituted by the polar group then due to an increase in the share of a negative charge, carbon can also take part in hydrogen bonding e.g. in carboxylic acids, amino acids, etc.

Types of hydrogen bonding:

There are two types of hydrogen bonding.

(a) Intermolecular hydrogen bonding :

Such a type of hydrogen bonding takes place between two or more similar or dissimilar molecules. The following example illustrates this type of hydrogen bonding.

(1) In H2O molecules :

The oxygen end carries a partial negative charge (−δ) and the hydrogen end carries the part positive (+δ) charge. therefore H2Omolecules form hydrogen bonds as shown below:
Hydrogen bonding in H2O
Hydrogen bonding provides some specific characteristics and properties of the H2O molecule. e.g. When water is cooled, its thermal mobility is reduced. which results in the molecules coming closer and the density increasing. Its density increases up to 4∘C but on further cooling, the density decreases. Because H2O molecules align themselves in such a way that the oxygen atoms are tetrahedrally surrounded by lour hydrogen atoms leaving some empty space as shown in figure 3.13. Consequently, the density begins to decrease as the water freezes and becomes minimum for ice.
Intermolecular hydrogen bonding in ice.

(2) In lower carboxylic acids :

When two formic acid molecules come closer to each other they form intermolecular hydrogen bonding and dimerize. This results, it has a unique property from the other monocarboxylic acids. For example, the High boiling point of this acid as compared to other organic acids is due to intermolecular hydrogen bonding.
(日) Dimeric structure of formic acid, HCOOH (b) Dimeric structure of acetic acid. CH3COOH Figure 3.14: Intermolecular hydrogen bonding in lower carboxylic acids
Like formic acid hydrogen bonding between two molecules is apparent in acetic acid also. As a result of hydrogen bonding, this carboxylic acid also exists as a dimer and its molecular weight is found to be double than calculated from its simple formula. e.g. The molecular weight, from its formula is calculated to be 60 but in fact, determined from its vapor density(V.D.) is 120. Other molecules in which such type of hydrogen bonding is formed are para-nitrophenol simple alcohols, alcohol-water, urea-water, para chlorophenol, sugars with water, NH3, etc.

Cheitical nonding (V.D.) is 120. Other molecules in which such type of hydrogen bonding formed are para-nitrophenol simple alcohols, alcohol-water, urea-water, para chlorophenol, sugars with water, etc. Hydrogen bonding in ammoria 2 , and other alkyl group, Hydrogen bonding in alcohols 3, Hydrogen bonding in Urea - waler systom Hydrogen bonding in tomaldehyde - water system Hydrogen bonding in alcohot - water system B. Life witl exist due to hydrogen bonding. Most of the biomolecules are held together by weak hydrogen bonding which breaked and furthor linked as when required. The mędicines work's on the same principle:

  • A life well exists due to hydrogen bonding. Most of the biomolecules are held together by weak hydrogen bonding which breaks and is further linked when required. The mędicines work’s on the same principle:

(b) Intramolecular hydrogen bonding :

When a hydrogen bond is formed within a single molecule between two different functional groups is called intramolecular hydrogen bonding or ‘internal hydrogen bonding: This leads to the formation of a 5 or 6-membered ring structure. Thus, this hydrogen bonding is also known as chelation. For example ortho-nitrophenol, ortho-chloro-phenol, salicylic acid, salicylaldehyde, etc. form chelates due to intramolecular hydrogen bonding as shown below:
Intramolecular hydrogen bonding
The bolling point of O-nitrophenol is 214∘C whereas its para and meta isomers have the boiling points of 270∘C and 290∘C respectively. The reason for the low boiling point of ortho-nitro-phenol is due to the intramolecular hydrogen bond whereas the other two isomers show intermolecular hydrogen bonds and therefore possess higher boiling points.
The evidence for hydrogen bonding is determined on the basis of spectroscopic studies like X-ray, neutron diffraction, electron diffraction, and infra-red studies. For example, the dimeric structure of formic acid is confirmed by electron diffraction studies.

Properties of hydrogen bond :

(i) Hydrogen bond length is longer than a normal covalent bond.
(ii) Bond energy of a hydrogen bond is in the range of 3−10 kcal mole while that of a normal covalent bond is in the range of 5−100kcal/mole.
(iii) The formation of a hydrogen bond does not involve any sharing of electron pairs:
(iv) They are formed in highly polar molecules like H2O, NH3, HF, etc.

Effect of Hydrogen Bonding(H.B) on physical properties :

1- Unique behavior of H2O:

As already discussed above in the crystal structure of ice that each oxygen atom is linked with four H-atoms; two of these H-atoms are covalently bonded with an oxygen atom and the remaining two with hydrogen bonds. Now in the structure of Ice, each water molecule is associated with four other water molecules in the tetrahedral arrangement by hydrogen bonding. Thus, during the formation of such a structure, there is a large empty space between the four water molecules as depicted in figure 3.13. As soon as the ice melts these hydrogen bonds are broken and the space enclosed between four H2O molecules disappears due to thermal expansion. This is why the density of ice is Jess than that of liquid water and ice floats on liquid water. It is also observed that the density of water increases from 0 to 4∘. It is because as the temperature rises from 0∘C melting and thermal occurs which brings water molecules to close together. Consequently, the density increases. But above 4∘C the density of liquid further decreases with an increase in kinetic energy of H2O molecules.

2. Boiling points of hydrides :

It was observed that hydrides of most electronegative elements of VA, VIA, and VIIA like

N, O, and F show many high values of boiling points. The reason is an association of large Tanber of molecules due to intermolecular hydrogen bonding. As we move downward in any particular of these groups, there is a sharp decrease in boiling point but later on, the increasing trend is observed as shown in table 3.3.
Note The highest value of boiling points of hydrides of most electronegative elements N. O and F is due to the hydrogen bonding.

boiling point of hydried VA, VIA, and VIIA
3. The viscosity of liquids :

In the case of molecules having a greater tendency to form intermolecular hydrogen bonding, there is a decrease in the tendency of liquid to flow smoothly. This is due to the fact that a large number of liquid molecules associate together to form a cluster, This cluster formation hinders their smooth flow and increases the viscosity. Let us consider an example of glycerol and alcohol like propanol, both having three C-atoms in the molecule. Glycerol has three OH groups whereas propanol has only one OH group. Hence, the association of glycerol s higher than propanol showing a very high value of viscosity as compared to propanol. Glycerol has a viscosity of 104 millipores whereas propanol has a value of viscosity below 10 millipores.

4- Formation of HF2− ion :

HF molecule combines with F−ion and forms HF2− ion,
This reaction is possible due to the high electronegativity of fluorine and its small size. On the other hand, other halides like Cl, Br, 1 – ions do not form such types of ions. The reason is the lack of hydrogen bond between these ions and their acids.

5- Determination of molecular crystal structure :

Due to the hydrogen bonding, many molecular crystals have their linear chain structure like HCN, zig-zag chain structure of HFCH3OH formic acids, ethers, sheet structure of oxalic acid, and tetrahedral structure of ice. All these structures are possible because of hydrogen bonding.

6- Heat of Vapourization :

The heat required to convert a liquid into its vapor is known as the heat of vaporization. If a compound has hydrogen bonding then it will require more energy to rupture this bonding and the heat of vaporization will be high.

7 – Comparison between o, m– and para isomers of aromatic compounds :

O-nitro phenol forms intramolecular hydrogen bonding whereas meta and para isomers show intermolecular hydrogen bonding that is polymerization. Due to this reason, the m.p.’s of o-nitrophenol is low whereas the other two have higher m.p.’s. There is also a difference in their solubility and volatility as shown in table 3.4
Table 3.4: Melting point and other physical properties of nitrophenols.
Nitro phenol M.P. Solubility in H2O Volatility
o 45∘ less soluble highly volatile
m 96∘ more soluble less volatile
p 114∘ highly soluble non-volatile


8- Solubility in water:

When the number of hydroxyl groups in an organic compound is more than the hydrocarbon groups, the substance will be soluble in water to a greater extent. For example, sugar and glycerol are highly soluble, and substances with very high molecular weight which contain (−CH2−CHOH−) n linkage are also sparingly soluble. Detergents that find extensive use in our daily life owe their cleansing action due to hydrogen bonding.
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