The document discusses non-covalent interactions, which differ from covalent bonds by not involving electron sharing, and are crucial for the structure and function of biological systems. Non-covalent interactions are categorized into electrostatic, van der Waals forces, π-effects, and hydrophobic effects, each with unique characteristics and implications for molecular behavior. These interactions drive processes such as protein folding and molecular recognition, influencing the behavior of supermolecules in supermolecular chemistry.

1. NON COVALENT BONDS
VIPIN MOHAN
2011-09-112
College of Agriculture
Vellayani, TVm

2. COVALENT INTERACTIONS
• Covalent interactions (bonds) provide the glue that holds biopolymers together. Covalent
bond energies are on the order of 100 kcal/mole.

3. NON COVALENT INTERACTIONS
• A non-covalent interaction differs from a covalent bond in that it does not involve
the sharing of electrons, but rather involves more dispersed variations
of electromagnetic interactions between molecules or within a molecule.
• The energy released in the formation of non-covalent interactions is typically on the
order of 1-5 kcal/mol .
• Non-covalent interactions can be generally classified into 4
categories: electrostatic, π-effects, van der Waals forces, and hydrophobic effects.
• In fact, van der Waals forces are responsible for why geckos can walk up and down
walls!

4. • Non-covalent forces drive spontaneous folding of proteins and nucleic acids
and mediate recognition of complementary molecular surfaces.
• Noncovalent forces dictate conformation and interaction in biological
systems.
• Non-covalent interactions are the dominant type of interaction
between supermolecules in supermolecular chemistry.

5. • 1 Electrostatic Interactions
• Ionic
• H-bonding
• Halogen Bonding
• 2 Van der Waals Forces
• Dipole-Dipole
• Dipole-Induced Dipole
• London Dispersion Forces
• 3 π-effects
• π-π Interaction
• Cation-π & Anion-π
• Polar-π
• 4 Hydrophobic effect

6. IONIC
• It involve the attraction of ions or molecules with full permanent charges of
opposite signs.
• Ionic interactions occur between cations and anions.
• These bonds are non-directional, and strength depends on the distance of
separation (r) according to 1/r2. Strength also depends on the medium
(dielectric constant), and is less in polar than nonpolar solvents.

7. H-BONDING
• A hydrogen bond (H-bond), is a specific type of dipole-dipole interaction that
involves the interaction between a partially-positive hydrogen atom and a highly
electronegative atom .
• It is technically not a covalent bond, but instead electronegative, partially-negative
oxygen, nitrogen, sulfur, or fluorine is classified as a very strong dipole-dipole (non-
covalent) interaction.
• Most commonly, the strength of hydrogen bonds lies between 0 - 4 kcal/mol, but
can sometimes be as strong as 40 kcal/mol

9. HALOGEN BONDING
• Halogen bonding is a type of non-covalent interaction which does not
involve the formation nor breaking of actual bonds, but rather is similar to
the dipole-dipole interaction known as hydrogen bonding.
• In halogen bonding, a halogen atom acts as an electrophile, or electron-
seeking species, and forms a weak electrostatic interaction with
a nucleophile, or electron-rich species.
• The nucleophilic agent in these interactions tends to be
highlyelectronegative (such as oxygen, nitrogen, or sulfur), or may
be anionic, bearing a negative formal charge.
• As compared to hydrogen bonding, the halogen atom takes the place of the
partially-positively charged hydrogen as the electrophile.

10. VAN DER WAALS FORCES
• Van der Waals Forces are a subset of electrostatic interactions involving permanent
or induced dipoles (or multipoles). These include the following:
• permanent dipole-dipole interactions, alternatively called the Keesom force
• dipole-induced dipole interactions, or the Debye force
• induced dipole-induced dipole interactions, commonly referred to as London
dispersion forces
• Note: Although hydrogen bonding and halogen bonding are both forms of dipole-
dipole interactions, these are typically not classified as Van der Waals Forces by
convention.

11. Dipole-Dipole
Dipole-dipole interactions are electrostatic interactions between
permanent dipoles in molecules. These interactions tend to align
the molecules to increase attraction (reducing potential energy).
Normally, dipoles are associated with electronegative atoms,
including (but not limited to) oxygen, nitrogen, sulfur, and fluorine.

13. DIPOLE-INDUCED DIPOLE
• A dipole-induced dipole interaction (Debye force) is due to the approach of a
molecule with a permanent dipole to another non-polar molecule with no
permanent dipole.
• This approach causes the electrons of the non-polar molecule to
be polarized toward or away from the dipole (or "induce" a dipole) of the
approaching molecule.
•

14. LONDON DISPERSION FORCES
• London dispersion forces are the weakest type of non-covalent interaction.
• They are also known as "induced dipole-induced dipole interactions", and
form from molecules that inherently do not have permanent dipoles.
• They are caused by the temporary repulsion of electrons away from the
electrons of a neighboring molecule, leading to a partially-positive dipole on
one molecule and a partially-negative dipole on another molecule.
• Hexane is a good example of a molecule with no polarity or highly
electronegative atoms.

16. π-EFFECTS
• π-effects can be broken down into numerous categories, including ,
π-π interactions, cation-π & anion-π interactions, and polar-π interactions.
• In general, π-effects are associated with the interactions of molecules with the π-
systems of conjugated molecules such as benzene.

17. π-π INTERACTION
• π-π interactions are associated with the interaction between the π-orbitals of a
molecular system.
• For a simple example, a benzene ring, with its fully conjugated π cloud, will interact
in two major ways and one minor way’ with a neighboring benzene ring through a
π-π interaction.
• The two major ways that benzene stacks are edge-to-face, with an enthalpy of ~2
kcal/mol, and displaced (or slip stacked), with an enthalpy of ~2.3 kcal/mol.
Interestingly, the sandwich configuration is not nearly as stable of an interaction as
the previously two mentioned due to high electrostatic repulsion of the electrons in
the π orbitals.

18. CATION π&ANION π
• Cation-π interactions involve the positive charge of
a cation interacting with the electrons in a π-system of a
molecule.
• This interaction is surprisingly strong (as strong or stronger
than H-bonding in some contexts), and has many potential
applications in chemical sensors.
• For example, the sodium ion can easily sit atop the π cloud of a
benzene molecule, with C6 symmetry (for more on point groups
and molecular symmetry.

19. • Anion-π interactions are very similar to cation-π interactions,
but reversed.
• In this case, an anion sits atop an electron-poor π-system,
usually established by the placement of electron-withdrawing
substituents on the conjugated molecule.

20. POLAR-π
• Polar-π interactions involve molecules with permanent dipoles (such as water)
interacting with the quadrupole moment of a π-system (such as that in benzene .
• While not as strong as a cation-π interaction, these interactions can be quite
strong (~1-2 kcal/mol), and are commonly involved in protein folding and
crystallinity of solids containing both hydrogen bonding and π-systems.
• In fact, any molecule with a hydrogen bond donor (hydrogen bound to a highly
electronegative atom) will have favorable electrostatic interactions with the
electron-rich π-system of a conjugated molecule.

21. HYDROPHOBIC EFFECT
• The hydrophobic effect is the desire for non-polar molecules to
aggregate in aqueous solutions in order to separate from water.
• This phenomenon leads to minimum exposed surface area of
non-polar molecules to the polar water molecules (typically
spherical droplets), and is commonly used in biochemistry to
study protein folding and other various biological phenomenon.
• olive oil in water