Organic Chemistry

Module 1. Hydrogen bonding and hydrophobic interactions

Lesson 2
INTER AND INTRA MOLECULAR HYDROGEN BONDING IN ALCOHOLS, CARBOXYLIC ACIDS AND OTHER MOLECULES AND THEIR SIGNIFICANCE

2.1 Introduction

• When hydrogen bonding is formed between atoms of different molecules, it is called intermolecular hydrogen bonding.
• All the examples given so far in Lesson 1 are of intermolecular hydrogen bonding (HF , H₂O etc).
• Intermolecular- between different molecules.
• When hydrogen bonding is formed between atoms of same molecule, it is called intramolecular hydrogen bonding.
• Intramolecular- within the molecule
• Examples are given below
• Intramolecular hydrogen bonding forms when –O-H group and electronegative element atom in the same molecule are present close together in such a position that a ring can form without disturbing the normal bond angles.
• The new five- or six- membered rings formed above as a result of intramolecular hydrogen bonding are known as chelate rings and such compounds are often referred to as chelate compounds (phenomenon as chelation).

2.2 Conditions For Hydrogen Bonding

• An effective hydrogen bond will form only when hydrogen atom is covalently attached to an atom which is

• Strongly electronegative and
• Small in size

• When the electronegativity is not high or if that atom has a large atomic radius

• The electrostatic forces of attraction will be weak and
• As a result hydrogen bond will not be very effective

• Fluorine, oxygen and nitrogen are the only three elements

• Which have sufficiently high electronegativity and
• Are small enough to form effective hydrogen bonds

• Chlorine has electronegativity comparable to that of nitrogen, still it does not form effective hydrogen bonding due to its relatively larger size
• Bromine and iodine are not as highly electronegative as required for hydrogen bonding.

2.3 Strength of Hydrogen Bonding

• Being electrostatic in nature, they are much weaker than the covalent bonds
• The strength of a hydrogen bonding is of the order of

• 2 to 10 kcal/mole or
• 10 to 40 kJ/mole (i.e. per 6.022 x 10²³ bonds)

• The strength of a normal covalent bond is of the order of

• 50-100 kcal/mole

• Greater the electronegativity and smaller the size of the electronegative atom, stronger is the hydrogen bond
• Therefore, hydrogen bonding involving F, O or N atoms have strengths of 10, 7 and 2 kcal/mole respectively.

2.4 Effects of Hydrogen Bonding on Physical Properties

• The existence of hydrogen bonding in molecules has a marked effect on their physical properties such as melting and boiling points, solubility, spectral characteristics, density etc.

2.4.1 Effect on melting and boiling points

• Ordinarily, compounds with similar molecular weights have similar melting/boiling points and there is regular increase in these physical constants with rising molecular weights.
• Compounds containing intermolecular hydrogen bonding show unusually high melting/boiling points.
• This is evident from boiling points of hydrogen compounds with elements of group V, VI, VII- shown in the Table below

Table 2.1 Boiling point of hydrogen bonded compounds

• In the series of hydrogen compounds given below, all the compounds have almost similar mol. wt., but there is a wide difference in their boiling points.

Table 2.2

• n-Butane contains only hydrogen and carbon lacks protonic hydrogen therefore non-polar covalent bond- no hydrogen bond formation- no molecular association- lower boiling point.
• Other compounds contain strongly polar covalent bond (N-H or O-H).
• Leads formation of intermolecular hydrogen bonding causes large molecular association - therefore much more energy would be needed to break them apart - therefore higher boiling points.
• The effect of intramolecular hydrogen bonding on melting/boiling points is opposite to that observed in the case of intermolecular hydrogen bonding
• Example = o-Nitrophenol (m.p. 44 °C) (steam volatile) and p-Nitrophenol (m.p. 114 °C) (non-volatile in steam) (Fig. 2.2)
• In ortho derivatives presence of intramolecular hydrogen bonding prevents formation of intermolecular hydrogen bonding- therefore intermolecular association does not takes place- therefore they have lower melting point and higher volatility.
• In para derivatives intermolecular hydrogen bonding formation takes places- leads to intermolecular association resulting into an increase in melting point and lower volatility.
• Another example : cis- and trans isomers of HOOC-CH=CH-COOH (Fig. 2.3 Melting points of Maleic acid and Fumaric acid on the basis of hydrogen bonding)
  

2.4.2 Effect on solubility

• Compounds which can form intermolecular hydrogen bonds with a solvent (water) would be generally soluble in that solvent
• Compounds in which hydrogen bonding with solvent (water) molecules is restricted (prevented) due to intramolecular hydrogen bonding would be less soluble or insoluble in that solvent
• Examples

• Hydroxy compounds like methyl alcohol, ethyl alcohol, sugars, etc are highly soluble in water due to formation of hydrogen bonds between molecules of these compounds and molecules of water
• In o-nitrophenol, - OH group is not available for hydrogen bonding with water- therefore it is less soluble in water. Whereas, in p-nitrophenol – OH group is available for hydrogen bonding with water - therefore it is more soluble in water

2.4.3 Stability of crystal structure

• Intramolecular hydrogen bonding in maleic acid halves its ability to from intermolecular bonds
• In fumaric acid all hydrogen bonds are intermolecular = therefore it gives stronger, interlinked crystal structure

2.4.4 Effect on spectral characteristics

• Spectral characteristic of compounds are significantly affected if these compounds contain hydrogen bonds- examples.
• An Infra-red (IR) spectral study of free –OH group at 3500 cm^-1 is shifted to lower frequencies if the –OH group involves hydrogen bonding -around 3200 cm^-1 in alcohol and around 3000 cm^-1 in carboxylic acids
• Much higher shift in the acids compared to alcohols- suggests that H bonds in the acids are stronger than that in alcohols.

2.4.5 Unusual behaviors of water

• Two unusual behaviors of water

1. Lower density in solid state (ice) than that in liquid state - therefore ice floats over water – unusual – because density of solid is lower
2. Water contract when heated between 0°C and 4°C - volume is minimum at 4°C - therefore its density is maximum at 4°C - unusual= substances expand when heated in all temperature ranges

• Both these peculiar features of water are attributed to formation of hydrogen bonding by water molecules (Fig. 2.4 The tetrahedral hydrogen bonded cage-like structure of ice)
  

• Hydrogen atoms
• Central oxygen atom
• Four oxygen atoms surrounding the central hydrogen atom
• Hydrogen bonding between water molecules is more extensive in ice than in liquid water.
• Because a substance in solid state has a definite structure and the molecules are more rigidly fixed relative to one another than in liquid state.
• In ice, the water molecules are tetrahedrally oriented with respect to one another
• Each oxygen atom is surrounded tetrahedrally by four hydrogen atoms, two of these are bonded covalently and the other two by hydrogen bonds
• The hydrogen bonds are weaker- therefore longer than the covalent bonds- this arrangement gives rise to an open cage-like structure – as indicated in the figure given above.
• There are number of holes or open spaces – because the hydrogen bonds holding the water molecules in ice are directed in definite angles
• In liquid water such hydrogen bonds are fewer in number
• As a result, when ice melts a large number of hydrogen bonds are broken and molecules move into the ‘holes’ or ‘open space’ - therefore come closer to one another than they were in solid state.
• As liquid water is heated from 0 to 4°C, hydrogen bonds continue to be broken and the molecules come closer and closer together
• This leads to contraction – decrease in volume- increase in density
• On rise of temperature from 0 to 4°C some expansion takes place, but contraction effect predominates - therefore there is net contraction.
• Above 4°C - expansion effect predominates – rise in volume (d=m/v). So water has maximum density at 4°C since v (volume) is minimum at 4°C.

2.5 Importance of Hydrogen Bonding

• Phenomenon of hydrogen bonding formation bears a great significance in various aspects

• Retention of water on the earth in large amount

• Without hydrogen bonding, water would have existed as a gas like hydrogen sulphide
• In that case, no life would have been possible on this globe

• Determines and maintains structure of various proteins which are so essential for life
• Makes wood fibers more rigid and thus makes it an article of great utility to meet requirements of housing, furniture etc
• Provides rigidity and tensile strength to cotton and silk fibers which are of vital importance for our clothing
• Forms linkage with water – cotton cutting takes more time for drying
• Affects properties of food constituent – changes viscosity, solubility and stability- desirable or undesirable
• Involves in –chemotherapeutic action of drugs, binding of dyes to textiles, adhesive action of paints, lacquers etc
• Collects water in animal and vegetable cells – required for various activities of the cells.
• Exists a part in biomolecules of living cells – DNA, RNA etc
• Formation retains moisture in each crust – i.e. with clay
• Survival of marine life in aquatic regions.

• Water can form hydrogen bond with very large number of compound – because it can form hydrogen bond by

1. Providing proton

e.g. its hydrogen bond formation with oxygen atom of carbonyl group (Fig. 2.5)

2. Accepting proton

e.g. its hydrogen bond formation with hydrogen atom of amine groups (Fig. 2.6)