Non-covalent interaction

Source: Wikipedia, the free encyclopedia.
(Redirected from
Non-covalent
)

In

van der Waals forces, and hydrophobic effects.[3][2]

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

Scheme 1. Process of NaF formation -- example of an electrostatic interaction

Ionic interactions involve the attraction of ions or molecules with full permanent charges of opposite signs. For example, sodium fluoride involves the attraction of the positive charge on sodium (Na+) with the negative charge on fluoride (F).[9] However, this particular interaction is easily broken upon addition to water, or other highly polar solvents. In water ion pairing is mostly entropy driven; a single salt bridge usually amounts to an attraction value of about ΔG =5 kJ/mol at an intermediate ion strength I, at I close to zero the value increases to about 8 kJ/mol. The ΔG values are usually additive and largely independent of the nature of the participating ions, except for transition metal ions etc.[10]

These interactions can also be seen in molecules with a localized charge on a particular

cation
(Na+).

Hydrogen bonding

Hydrogen-bonding-in-water

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

Figure 1. Anionic Lewis base forming a halogen bond with electron-withdrawn bromine (Lewis acid)

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.[citation needed
]

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

aromatic π-cloud as a nucleophile; halogen bonding is restricted to monatomic nucleophiles.[5]

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
    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

Figure 2. Dipole–dipole interactions between two acetone molecules, with the partially negative oxygen atom interacting with the partially positive carbon atom in the carbonyl.

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 oxygen, nitrogen, sulfur, and fluorine
.

For example,

carbonyl
(see figure 2). Since oxygen is more electronegative than the carbon that is covalently bonded to it, the electrons associated with that bond will be closer to the oxygen than the carbon, creating a partial negative charge (δ) on the oxygen, and a partial positive charge (δ+) on the carbon. They are not full charges because the electrons are still shared through a covalent bond between the oxygen and carbon. If the electrons were no longer being shared, then the oxygen-carbon bond would be an electrostatic interaction.

Often molecules contain dipolar groups, but have no overall

tetrachloromethane. Note that the dipole-dipole interaction between two individual atoms is usually zero, since atoms rarely carry a permanent dipole. See atomic dipoles
.

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.[13] Specifically, the dipole can cause electrostatic attraction or repulsion of the electrons from the non-polar molecule, depending on orientation of the incoming dipole.[13] Atoms with larger atomic radii are considered more "polarizable" and therefore experience greater attractions as a result of the Debye force.[citation needed
]

London dispersion forces

sublimation
heat of crystals is a measure of the dispersive interaction. While these interactions are short-lived and very weak, they can be responsible for why certain non-polar molecules are liquids at room temperature.

π-effects

cation-π and 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.[3]

π–π interaction

Figure 3. Various ways that benzene can interact intermolecularly. Note, however, that the sandwich configuration is not a favorable interaction compared to displaced or edge-to-face

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

Figure 4

cation interacting with the electrons in a π-system of a molecule.[3] This interaction is surprisingly strong (as strong or stronger than H-bonding in some contexts),[3] and has many potential applications in chemical sensors.[20] For example, the sodium ion can easily sit atop the π cloud of a benzene molecule, with C6 symmetry
(See figure 4).

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]

Figure 5.

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

dipole–dipole interactions
.

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

angle strain
also play major roles in the folding of a protein from its primary sequence to its tertiary structure.

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

electrostatic interactions
.

Boiling points

Non-covalent interactions have a significant effect on the

n-butanol
(C4H9OH).

Figure 8. Boiling points of 4-carbon compounds

The predominant non-covalent interactions associated with each species in solution are listed in the above figure. As previously discussed,

intermolecular forces
each molecule experiences in its liquid state.

References