Intermolecular forces- Organic Chemistry

1. Intra- & Intermolecular
forces
Dr. Rajasekhar Reddy A
K L COLLEGE OF PHARMACY,
KLEF DEEMED TO BE UNIVERSITY

2. Intermolecular Forces
• Intermolecular forces are the interactions that exist between
molecules.
• A functional group determines the type and strength of
these interactions.
• Intermolecular forces are also referred to as noncovalent
interactions or nonbonded interactions.

3. Intermolecular Forces in Ionic
Compounds
• Ionic compounds contain oppositely charged particles held together
by extremely strong electrostatic interactions.
• These ionic interactions are much stronger than the intermolecular
forces present between covalent molecules, so it takes a great deal
of energy to separate oppositely charged ions from each other

4. Intermolecular Forces in Covalent
Compounds
• Covalent compounds are composed of discrete molecules.
• The nature of the forces between the molecules depends on
the functional group present.
• There are three different types of interactions, presented here
in order of increasing strength:
• van der Waals forces
• dipole–dipole interactions
• hydrogen bonding

5. van der Waals Forces
• van der Waals forces, also called London forces, are very weak
interactions caused by the momentary changes in electron density in a
molecule.
• van der Waals forces are the only attractive forces present in nonpolar
compounds.
• For example, although a nonpolar CH4 molecule has no net dipole, at
any one instant its electron density may not be completely symmetrical,
creating a temporary dipole.
• This can induce a temporary dipole in another CH4 molecule, with the
partial positive and negative charges arranged close to each other.
• The weak interaction of these temporary dipoles constitutes van der
Waals forces. All compounds exhibit van der Waals forces.

6. • The surface area of a molecule determines the strength of the van der
Waals interactions.
• The larger the surface area, the larger the attractive force between two
molecules, and the stronger the intermolecular forces.
• Long, sausage-shaped molecules such as CH3CH2CH2CH2CH3 (pentane)
have stronger van der Waals interactions than compact spherical ones
like C(CH3)4 (neopentane).
• Another factor affecting the strength of van der Waals forces is
polarizability.
van der Waals Forces

7. van der Waals Forces

8. • Larger atoms like iodine, which have more loosely held valence
electrons, are more polarizable than smaller atoms like fluorine, which
have more tightly held electrons.
• Because larger atoms have more easily induced dipoles, compounds
containing them possess stronger intermolecular interactions.
• Thus, two F2 molecules have little force of attraction between them,
because the electrons are held very tightly and temporary dipoles are
difficult to induce.
• On the other hand, two I2 molecules exhibit a much stronger force of
attraction because the electrons are held much more loosely and
temporary dipoles are easily induced
van der Waals Forces

9. Dipole–Dipole Interactions
• Dipole–dipole interactions are the attractive forces between the
permanent dipoles of two polar molecules.
• In acetone, (CH3)2C=O, for example, the dipoles in adjacent molecules
align so that the partial positive and partial negative charges are in close
proximity.
• These attractive forces caused by permanent dipoles are much stronger
than weak van der Waals forces.

10. Hydrogen Bonding
• Hydrogen bonding typically occurs when a hydrogen atom bonded
to O, N, or F, is electrostatically attracted to a lone pair of electrons
on an O, N, or F atom in another molecule.
Thus, H2O molecules can hydrogen bond to each other. When they do, an
H atom covalently bonded to O in one water molecule is attracted to a lone
pair of electrons on the O in another water molecule.
• Hydrogen bonds are the strongest of the three types of intermolecular
forces, though they are still much weaker than any covalent bond.

11. Summary of Types of Intermolecular
Forces

14. Intramolecular forces – Ionic Bond
• Atoms combine with one another to give compounds having properties
different from the atoms they contain.
• The attractive force between atoms in a compound is a chemical bond.
• One type of chemical bond, called an ionic bond, is the force of
attraction between oppositely charged species (ions).
• Ions that are positively charged are referred to as cations; those that
are negatively charged are anions.
In forming ionic compounds, elements at the left of the periodic table typically lose
electrons, forming a cation that has the same electron configuration as the nearest
noble gas.

15. • A large amount of energy, called the ionization energy, must be added to any
atom in order to dislodge one of its electrons.
• The ionization energy of sodium, for example, is 496 kJ/mol (119 kcal/mol).
• Processes that absorb energy are said to be endothermic.
• In general, ionization energy increases across a row in the periodic table.
• Elements at the right of the periodic table tend to gain electrons to reach the
electron configuration of the next higher noble gas.
• Energy-releasing reactions are described as exothermic, and the energy
change for an exothermic process has a negative sign.
• The energy change for addition of an electron to an atom is referred to as its
electron affinity and is -349 kJ/mol (-83.4 kcal/mol) for chlorine.
Intramolecular forces – Ionic Bond

16. • Transfer of an electron from a sodium atom to a chlorine atom
yields a sodium cation and a chloride anion, both of which have a
noble gas electron configuration:
• The ionization energy of sodium (496 kJ/mol) and the electron
affinity of chlorine (-349 kJ/mol), the overall process is
endothermic with DH°= +147 kJ/mol.
• The energy liberated by adding an electron to chlorine is
insufficient to override the energy required to remove an electron
from sodium.
• This analysis, however, fails to consider the force of attraction
between the oppositely charged ions Na+ and Cl–, which exceeds
500 kJ/mol and is more than sufficient to make the overall process
exothermic.
• Attractive forces between oppositely charged particles are termed
Intramolecular forces – Ionic Bond

17. COVALENT BONDS
• The covalent, or shared electron pair, model of chemical bonding was
first suggested by G. N. Lewis of the University of California in 1916.
• Lewis proposed that a sharing of two electrons by two hydrogen
atoms permits each one to have a stable closed-shell electron
configuration analogous to helium.
• The amount of energy required to dissociate a hydrogen molecule H2
to two separate hydrogen atoms is called its bond dissociation energy
(or bond energy).
• For H2 it is quite large, being equal to 435 kJ/mol (104 kcal/mol).

18. DOUBLE BONDS AND TRIPLE BONDS
• Lewis’s concept of shared electron pair bonds allows for 4-electron
double bonds and 6-electron triple bonds.
• Carbon dioxide (CO2) has two carbon–oxygen double bonds, and the
octet rule is satisfied for both carbon and oxygen.
• Similarly, the most stable Lewis structure for hydrogen cyanide (HCN)
has a carbon–nitrogen triple bond.

19. Hydrogen Bonding - Intramolecular
• A hydrogen bond is a special type of dipole-dipole attraction which
occurs when a hydrogen atom bonded to a strongly electronegative
atom exists in the vicinity of another electronegative atom with a lone
pair of electrons.
• These bonds are generally stronger than ordinary dipole-dipole and
dispersion forces, but weaker than true covalent and ionic bonds.

20. The cohesion-adhesion theory of transport in vascular plants uses
hydrogen bonding to explain many key components of water movement
through the plant's xylem and other vessels.

21. • Hydrogen bonding is present abundantly in the
secondary structure of proteins, and also sparingly
in tertiary conformation.
• The secondary structure of a protein involves
interactions (mainly hydrogen bonds) between
neighboring polypeptide backbones which contain
N-H bonded pairs and oxygen atoms.

22. Inductive effects, Resonance, and Hyper
conjugation