Chemical bonding

1. Chemical Bonding
Course: B.Sc.
Microbiology/Biotechnology/Biochemistry
Sem II
Sub: Inorganic Chemistry
Unit 3.1

2. HYDROGEN BONDING
• A hydrogen atom : one electron - covalently bonded to only one
atom.
• hydrogen atom can involve itself in an additional electrostatic
bond with a second atom of highly electronegative character
such as fluorine or oxygen. This second bond permits a
hydrogen bond between two atoms or strucures.
• The strength of hydrogen bonding varies.
 Hydrogen bonds connect water
molecules in ordinary ice.
Hydrogen bonding is also very
important in proteins and nucleic
acids and therefore in life
processes.

3. Types of Hydrogen bond
• occur within one single molecule, between two like molecules, or between two
unlike molecules.
• 1. Intramolecular hydrogen bonds:
- occur within one single molecule.
- occurs when two functional groups of a molecule can form hydrogen bonds with
each other.
- both a hydrogen donor and acceptor must be present within one molecule
For example : ethylene glycol (C2H4(OH)2) between its two hydroxyl groups
• 2. Intermolecular hydrogen bonds
• occur between separate molecules in a substance.
• occur between any number of like or unlike molecules as long as hydrogen donors
and acceptors are present an in positions in which they can interact.
• For example, between NH3 molecules alone, between H2O molecules alone, or
between NH3 and H2O molecules.

4. Properties & effects of Hydrogen bond•
1. Boiling points of molecules : molecules with larger molar masses to have
higher normal boiling points than molecules with smaller molar masses.
This, without taking hydrogen bonds into account, is due to greater
dispersion forces. Larger molecules have more space for electron
distribution and thus more possibilities for an instantanous dipole
moment.
H2O, HF, and NH3 : each have higher boiling points than the same
compound formed between hydrogen.
• This is because H2O, HF, and NH3 all exhibit hydrogen bonding, whereas
the others do not.
•
• Furthermore, H2O has a smaller molar mass than HF but partakes in more
hydrogen bonds per molecule, so its boiling point is consequently higher.

5. 2. Viscosity
Those subtances which are capable of
forming hydrogen bonds tend to have a
higher viscosity than those that do not.
Substances which have the possibility for
multiple hydrogen bonds exhibit even higher
viscosities.

6. Factors preventing Hydrogen bonding
• 1. Electronegativity
• cannot occur without significant electronegativity differences between hydrogen and
the atom it is bonded to.
• Thus, we see molecules such as PH3, which no not partake in hydrogen bonding. PH3
exhibits a trigonal pyramidal molecular geometry like that of ammmonia, but unlike
NH3 it cannot hydrogen bond. This is due to the similarity in the electronegativities of
phosphorous and hydrogen. Both atoms have an electronegativity of 2.1, and thus, no
dipole moment occurs. This prevents the hydrogen bonding from acquiring the partial
positive charge needed to hydrogen bond with the lone electron pair in another
molecule.
• 2. Atom Size
• The size of donors and acceptors can also effect the ability to hydrogen bond.
• low ability of Cl to form hydrogen bonds.
• When the radii of two atoms differ greatly or are large, their nuclei cannot achieve close
proximity when they interact, resulting in a weak interaction.

7. 4 - VAN DER WAALS
BONDING
• weak bond.
• The intermolecular attractive forces operative between all
molecules, when they are close to one another –van dar waals
forces.
• Natural fluctuation in the electron density of all molecules and
these cause small temporary dipoles within the molecules.
• It is these temporary dipoles that attract one molecule to
another. They are called van der Waals' forces.
• The bigger a molecule is, the easier it is to polarise (to form a
dipole), and so the van der Waal's forces get stronger, so bigger
molecules exist as liquids or solids rather than gases.

9. • Water molecules- attracted to each other by
electrostatic forces.
• Even though the water molecule as a whole is
electrically neutral, the distribution of charge
in the molecule is not symmetrical and leads
to a dipole moment - a microscopic separation
of the positive and negative charge centers.
This leads to a net attraction between such
polar molecules.

10. • shape : influences its ability to form temporary dipoles.
• Long thin molecules can pack closer to each other than
molecules that are more spherical.
• bigger the 'surface area' of a molecule : the greater the van
der Waal's forces will be and the higher the melting and
boiling points of the compound will be.
Homonuclear molecules,
such as iodine, develop
temporary dipoles due to
natural fluctuations of electron
density within the molecule
Heteronuclear molecules,
such as H-Cl have permanent
dipoles that attract the opposite
pole in other molecules.

11. The dipoles can be formed as a result of unbalanced distribution
of electrons in asymettrical molecules. This is caused by the
instantaneous location of a few more electrons on one side of the
nucleus than on the other.
symmetric asymmetric
Therefore atoms or molecules containing dipoles are attracted to
each other by electrostatic forces.

12. Ionic solids
• composed of oppositely charged ions.
• consist of positively charged cations and negatively
charged anions .
• dissolved in water : seperate,
• free to move about in the water allowing the
solution to conduct electrical current.

13. Properties of Ionic Compounds
– Have high melting and boiling temperatures.
– Are hard but brittle
• They also:
– Do NOT conduct electricity in the solid state
– They will only conduct electricity if they are melted or
dissolved in water
• The physical properties of ionic compounds are very
different from metals.

14. Lattice energy
• type of potential energy.
• Lattice energy is the energy required to break apart an ionic
solid and convert its component atoms into gaseous ions.
• Value for the lattice energy – positive : an endothermic
reaction.’

15. • “lattice energy is the reverse process, meaning it is the energy
released when gaseous ions bind to form an ionic solid.
• this process will always be exothermic, and thus the value for
lattice energy will be negative.
• units - kJ/mol.”

16. Born-Haber Cycle
• A series of hypothetical steps and their enthalpy changes
needed to convert elements to an ionic compound and
devised to calculate the lattice energy.
• Using Hess’s law.

18. Born-Haber Cycle Steps
1. Elements (standard state) - gaseous atoms
2. form cations and anions
3. Combining gaseous anions and cations to form a solid ionic
compound

19. Step 1: Atomisation
• The standard enthalpy change of
atomisation is the ΔH required to produce
one mole of gaseous atoms.
• Na(s)  Na(g) ΔHo
at = +109 kJmol-1

20. • NOTE: for diatomic gaseous elements, Cl2, ΔHo
at is
equal to half the bond energy (enthalpy).
• Cl2(g)  Cl(g) ΔHo
at = ½ E (Cl-Cl)
ΔHo
at = ½ (+242 )
ΔHo
at = +121 kJmol-1

21. Step 2: Formation of gaseous ions
• Electron Affinity
– Enthalpy change when one mole of gaseous atoms or
anions gains electrons to form a mole of negatively
charged gaseous ions.
• Cl(g) + e-  Cl-(g) ΔHo = -364 kJmol-1
– For most atoms = exothermic, but gaining a 2nd electron is
endothermic due to the repulsion between the anion and
the electron.

22. Becoming cations
• Ionisation energy
– Enthalpy change for one mole of a gaseous
element or cation to lose electrons to form a mole
of positively charged gaseous ions
• Na(g)  Na+(g) + e- IE1= +494 kJmol-1

23. Lattice Enthalpy
• Energy required to convert one mole of the solid
compound into gaseous ions.
• NaCl (s)  Na+(g) + Cl-(g)
• ΔHo
lat = +771kJmol-1
• highly endothermic
cannot directly calculate ΔHo
lat , but values are
obtained indirectly through Hess’s law for the
formation of the ionic compound

24. References
• Essentials of Physical chemistry by Bahl Arun, S Chand, 2012
• General chemistry by Ebbing Darrell D, 5th, A I T B S Publishers,
2002