Title: Hydrogen Bonding
Theory:
When hydrogen is bonded to a highly electronegative atom such as fluorine, oxygen and
nitrogen, the bonding electron pair is drawn towards the electronegative atom. As hydrogen
has no inner shell electron and is very small in size, the positive charge density developed is
high. The attraction between the lone pair electrons of the electronegative atom and slightly
positive charged hydrogen atom of another molecule is called a hydrogen bond. It is much
weaker than the covalent bond but much stronger than van der Waals’ forces. Its strength is
about 1/10 to 1/20 that of a typical covalent bond.
For Example, the intermolecular force in hydrogen fluoride is provided by hydrogen bonding.
As fluorine is the most electronegative element, the dipole moment of HF molecule is very
large. Actually, the particularly strong dipole-dipole interaction between HF molecules is
called hydrogen bonding.
The essential requirements for the formation of a hydrogen bond are:
1. A hydrogen atom must be directly bonded to a highly electronegative atom. Only the
most electronegative atoms — F, O and N can bring about strong interactions called
hydrogen bonding. In molecules such as HCl, the interaction is said to be of the van der
Waals’ type but not hydrogen bonding.
2. An unshared pair of electrons (lone pair electrons) on the electronegative atom.
Importance of Hydrogen Bonding in Physical Phenomena
Hydrogen bonding plays an important role in determining the structure and physical
properties of many compounds. There are two types of hydrogen bonding. One is the
intermolecular hydrogen bond formed between two different molecules. The other is the
intramolecular hydrogen bond formed between two different atoms in the same molecule.
The consequence of forming intramolecular hydrogen bonds is that the number of
intermolecular hydrogen bonds formed will be reduced.
About the experiment
Breaking or formation of intermolecular hydrogen bonds between molecules in liquids would
cause an enthalpy change when the liquids are mixed. This experiment is to investigate such
enthalpy changes and to measure approximate strengths of hydrogen bonds form between
molecules of ethanol and those between molecules of trichloromethane and ethyl ethanoate
using simple calorimetric method.
Experiments A and B show the dissociation of hydrogen bond of ethanol.
Experiment C and D show the formation of hydrogen bond between ethyl ethanoate and
trichloromethane.
The strength of the hydrogen bond formed between trichloromethane and ethyl ethanoate can
be determined by measuring the enthalpy change of mixing the two organic liquids.
Both trichloromethane and ethyl ethanoate are polar molecules. In their pure forms, there is
only dipole-dipole attraction between the molecules. However, after mixing them together, the
hydrogen atom of trichloromethane molecule and the oxygen atom of ethyl ethanoate
molecule become linked together by hydrogen bonding. The increase in strength of the
intermolecular force after mixing is accompanied by the release of heat (i.e. a rise in
temperature).
The hydrogen atom in trichloromethane is directly bonded to a carbon atom which has an
electronegativity value similar to that of hydrogen. But there are three C - Cl bonds in the
molecule.
The electronegative chlorine atoms tend to draw the pair of
bonding electrons away from the carbon atom. As a result, the
carbon atom attains positive charge character and the degree
of polarization of the C - H bond increases. The positive
charge density of the hydrogen atom becomes so large that
formation of the hydrogen bond is feasible.
An approximate value for the enthalpy change of formation for a hydrogen bond can be found
by a simple calorimetry experiment. An excess of one reagent (say, ethyl ethanoate) is mixed
with a fixed amount of trichloromethane and the temperature rise is recorded. It can be
assumed that all the trichloromethane molecules form hydrogen bonds. From the quantities of
reagents used, their specific heat capacities and temperature change on mixing, the enthalpy
change for one mole of trichloromethane that forms hydrogen bonds can be calculated. The
value found experimentally is about -2 to -4kJ/mol.
Chemicals: Ethanol, cyclohexane, ethyl ethanoate, trichloromethane
Apparatus: 10 cm3 and 25 cm3 cube measuring cylinders, 50 cm3 beaker, -10℃ to 110 ℃
thermometer.
Warnings!
Ethanol, cyclohexane, and ethyl ethanoate are flammable, trichloromethane is harmful. All
organic solvent should avoid inhaling vapour and skin contact.
Procedure:
This experiment is divided into 4 separate parts (A-D).
A. To discover the existence of hydrogen bonds between ethanol molecules
1.
2.
By the use of a measuring cylinder, 10cm3 of ethanol was added into an insulated 50cm3
breaker (insulated with paper towel). The temperature of the liquid was recorded.
Then 10cm3 of cyclohexane was added to the ethanol in the beaker, it was mixed well
and the lowest temperature attained was recorded.
B. To measure the strength of hydrogen bond formed between ethanol molecules
Repeat steps (1) and (2) in experiment A above using the same volume of ethanol but
20cm3 of cyclohexane. From the temperature drop estimate the hydrogen bond strength
(in kJ/mol) in ethanol.
C. To investigate the formation of hydrogen bonds between molecules of ethyl
ethanoate and trichloromethane
1. 10cm3 of ethyl ethanoate was measured into an insulated beaker. The temperature was
recorded.
2. 10cm3 of trichloromethane was added to this and mixed well. The highest temperature
attained was recorded.
D. To measure the strength of hydrogen bonds formed between molecules of ethyl
ethanoate and trichloromethane
Repeat steps (1) and (2) in experiment C above using either one liquid in excess. From
the temperature change estimate the hydrogen bond strength formed between molecules
of ethyl ethanoate and trichloromethane.
Results:
Given:
Liquid
Formula
Ethanol
CH3CH2OH
Cyclohexane
C6H12
Trichloromethane
CHCl3
Ethyl ethanoate CH3CO2CH2CH3
Relative
atomic
mass
Density
(kg/dm3 )
Specific heat
capacity
(kJ/kg K)
46
84
119.5
88
0.81
0.78
1.48
0.90
2.44
1.83
0.98
1.92
Experiment A:
(Ⅰ)
Temperature before the reaction
K
(Ⅱ)
Lowest temperature after the reaction
K
(Ⅲ)
Temperature Change
(Ⅱ)－(Ⅰ)
K
Therefore it is an (exothermic/endothermic) reaction.
Experiment B:
(Ⅰ)
Temperature before the reaction
K
(Ⅱ)
Lowest temperature after the reaction
K
(Ⅲ)
Temperature Change
(Ⅱ)－(Ⅰ)
K
Number of moles of ethanol
= mass/molar mass
=10g/46.00g
=
Energy absorbed by ethanol
mol
= mC
= 0.01×2.44×
=
kJ
Energy absorbed by cyclohexane = mC
= 0.02×1.83×
=
kJ
Total energy absorbed = total energy released from bond forming
+
=
kJ
Hydrogen bond strength
= total energy released from bond forming ÷ number of moles
=
÷
=
kJ/mol
Experiment C:
(Ⅰ)
Temperature before the reaction
K
(Ⅱ)
Highest temperature after the reaction
K
(Ⅲ)
Temperature Change
(Ⅱ)－(Ⅰ)
Therefore it is an (exothermic/endothermic) reaction.
K
Experiment D:
dm3
is in excess, its volume is
(Ⅰ)
Temperature before the reaction
K
(Ⅱ)
Highest temperature after the reaction
K
(Ⅲ)
Temperature Change
Number of moles of
(the one not in excess)
Energy released by
(the one in excess)
Energy released by
(Ⅱ)－(Ⅰ)
K
= mass/molar mass
=10g /
g
=
mol
= mC
=
=
kJ
=mC
=
=
kJ
Total energy released = total energy absorbed from bond forming
+
=
kJ
Hydrogen bond strength
= total energy absorbed from bond forming ÷ number of moles
=
÷
=
kJ/mol
Discussion:
1. If we use tetrachloromethane replacing trichloromethane in experiment C, explain the
different.
2. When we select an excess solvent to investigate the effect on H-bonding, what important
consideration should be applied?