Non covalent bonds

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.

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/r 2 . 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

8. 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 highly electronegative  (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.

9. 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.

10. 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 .

11. 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 .  

12. 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.

13. 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 .

14. 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.

15. 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  C 6  symmetry (for more on point groups and molecular symmetry.

16. 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 .

17. 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.

18. 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

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