Non-covalent interaction
In
Non-covalent interactions
The non-covalent interactions may occur between different parts of the same molecule (e.g. during protein folding) or between different molecules and therefore are discussed also as intermolecular forces.
Electrostatic interactions
Ionic
These interactions can also be seen in molecules with a localized charge on a particular
Hydrogen bonding
A hydrogen bond (H-bond), is a specific type of interaction that involves dipole–dipole attraction between a partially positive hydrogen atom and a highly electronegative, partially negative oxygen, nitrogen, sulfur, or fluorine atom (not covalently bound to said hydrogen atom). It is not a covalent bond, but instead is classified as a strong non-covalent interaction. It is responsible for why water is a liquid at room temperature and not a gas (given water's low molecular weight). Most commonly, the strength of hydrogen bonds lies between 0–4 kcal/mol, but can sometimes be as strong as 40 kcal/mol[3] In solvents such as chloroform or carbon tetrachloride one observes e.g. for the interaction between amides additive values of about 5 kJ/mol. According to Linus Pauling the strength of a hydrogen bond is essentially determined by the electrostatic charges. Measurements of thousands of complexes in chloroform or carbon tetrachloride have led to additive free energy increments for all kind of donor-acceptor combinations.[11][12]
Halogen bonding
Halogen bonding should not be confused with halogen–aromatic interactions, as the two are related but differ by definition. Halogen–aromatic interactions involve an electron-rich
Van der Waals forces
- permanent Keesom force
- dipole-induced dipole interactions, or the Debye force
- induced dipole-induced dipole interactions, commonly referred to as London dispersion forces
Hydrogen bonding and halogen bonding are typically not classified as Van der Waals forces.
Dipole–dipole
For example,
Often molecules contain dipolar groups, but have no overall
Dipole-induced dipole
A dipole-induced dipole interaction (
London dispersion forces
π-effects
π–π interaction
π–π interactions are associated with the interaction between the π-orbitals of a molecular system.[3] The high polarizability of aromatic rings lead to dispersive interactions as major contribution to so-called stacking effects. These play a major role for interactions of nucleobases e.g. in DNA.[19] 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 (see figure 3). 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.[3] 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.[3]
Cation–π and anion–π interaction
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[21]
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 (see figure 5). 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.[3] 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.[citation needed]
Hydrophobic effect
The hydrophobic effect is the desire for non-polar molecules to aggregate in aqueous solutions in order to separate from water.[22] 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.[22] The effect is also commonly seen when mixing various oils (including cooking oil) and water. Over time, oil sitting on top of water will begin to aggregate into large flattened spheres from smaller droplets, eventually leading to a film of all oil sitting atop a pool of water. However the hydrophobic effect is not considered a non-covalent interaction as it is a function of entropy and not a specific interaction between two molecules, usually characterized by entropy.enthalpy compensation.[23][24][25] An essentially enthalpic hydrophobic effect materializes if a limited number of water molecules are restricted within a cavity; displacement of such water molecules by a ligand frees the water molecules which then in the bulk water enjoy a maximum of hydrogen bonds close to four.[26][27]
Examples
Drug design
Most pharmaceutical drugs are small molecules which elicit a physiological response by "binding" to
Non-covalent metallo drugs have been developed. For example, dinuclear triple-helical compounds in which three ligand strands wrap around two metals, resulting in a roughly cylindrical tetracation have been prepared. These compounds bind to the less-common nucleic acid structures, such as duplex DNA, Y-shaped fork structures and 4-way junctions.[29]
Protein folding and structure
The folding of
Single tertiary protein structures can also assemble to form protein complexes composed of multiple independently folded subunits. As a whole, this is called a protein's
Boiling points
Non-covalent interactions have a significant effect on the
The predominant non-covalent interactions associated with each species in solution are listed in the above figure. As previously discussed,
References
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