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The Weak Noncovalent Interactions Have A Constitutive Role Biology Essay

Noncovalent interactions. The weak noncovalent interactions have a constitutive role in biological or biomimetic systems as well as in artificial supramolecular structures. Noncovalent or van der Waals interactions were first recognized by J. D. van der Waals in the nineteenth century . Their role in nature has been unravelled only during the past two decades.

In contrast to the covalent interactions that dominate in classical molecules, noncovalent interactions are weak interactions that bind together different kinds of building blocks into supramolecular entities . Covalent bonds are generally shorter than 2 Å, while noncovalent interactions function within range of several angstoms. The formation of a covalent bond require overlapping of partially occupied orbitals of interacting atoms, which share a pair of electrons. In noncovalent interactions, in turn, no overlapping is necessary because the attraction comes from the electrical properties of the building blocks.

The noncovalent interactions or van der Waals forces involved in supramolecular entities may be a combination of several interactions, e.g. ion-pairing, hydrogen bonding, cation−π , π −π interactions etc. . A wide range of attractive and repulsive forces is subsumed under term noncovalent. Noncovalent interactions comprise interactions between permanent multipoles, between a permanent multipole and an induced multipole, and between a time–variable multipole and an induced multipole. The stabilizing energy of noncovalent complexes is generally said to consist of the following energy contributions: electrostatic (or Coulombic), induction, charge transfer, and dispersion. These terms are basically attractive terms. The repulsive contribution, which is called exchange-repulsion, prevents the subsystems from drawing too close together. The term induction refers to general ability of charged molecules to polarize neigbouring species, and dispersion (London) interaction results from the interactions between fluctuating multipoles. In charge-transfer (CT) interactions the electron flow from the donor to the acceptor is indicated. The term van der Waals (vdW) forces is frequently used to describe dispersion and exchange-repulsion contributions, but sometimes also other long–range contributions are included in the definition. All of these interactions involve host and guest as well as their surroundings (e.g. solvation, crystal lattice and gas phase).

Noncovalent interactions are individually weak but collectively strong.

All three forms of noncovalent interactions are individually weak (on the order of 5 kcal/mole) as compared with a covalent bond (with its 90-100 kcal/mole of bond energy). And what strength these interactions do have requires that the interacting groups can approach each other closely (an angstrom or less). So we can conclude that all the examples given at the top of the page require:

a substantial number of noncovalent interactions working together to hold the structures together

a surface topography that enables substantial areas of two interacting surfaces to approach each other closely; that is, they must fit each other.

Examples are.-

The Interaction of Antibodies with Antigens

Antibodies are proteins

synthesized and secreted by B cells that

bind to antigens. Most antigens are macromolecules: proteins, polysaccharides, even DNA and RNA.

The interaction occurs:

by noncovalent forces (like that between enzymes and their substrate)

between

the antigen-combining site on the antibody and

a portion of the antigen called the antigenic determinant or epitope.

These photos show one type of interaction — precipitation — between antibodies and antigen. http://users.rcn.com/jkimball.ma.ultranet/BiologyPages/P/ppt_a-d.jpg

(a) The tube contains antibodies to the Type III pneumococcal polysaccharide isolated from the capsule surrounding the bacteria.

(b) A solution of the polysaccharide is added, and

(c) the formation of insoluble antigen-antibody complexes is revealed by the almost instantaneous appearance of turbidity.

(d) After an hour, the complexes settle out as a precipitate. If the proportion of antigen to antibody in the mixture is selected properly, the fluid above the precipitate will be devoid of both.

In the human body, this binding can literally be life-saving.

The capsule that surrounds pneumococci protects them from phagocytosis. (Pneumococci that fail to make a capsule — "R" forms — do not cause disease

If the appropriate antibodies are present in the body, they combine with the capsule. Coated with protein instead of polysaccharide, the pneumococci are now easy to ingest. http://users.rcn.com/jkimball.ma.ultranet/BiologyPages/P/pneumococci.jpg

These photomicrographs show phagocytosis of antibody-coated pneumococci.

Left: A neutrophil extends a pseudopod toward two pneumococci.

Center: these bacteria have been engulfed (arrows), and the neutrophil is beginning to engulf four more pneumococci at the upper right.

Right: Two pneumococci have escaped.

Other non-covalent interactions

There are several different types of interactions that are generally grouped under the term

non-covalent interactions. Typically pure electrostatic interactions (like Born or

Coulomb), while they are non-covalent themselves, are not counted in this category, but

they are instead referred to as electrostatic interactions. As it turns out the all the other

non-covalent interactions are actually electrostatic in nature as well and the little

contribution to these interactions that is not electrostatic turns out to be covalent in

nature. As you can see the whole nomenclature of these interactions is a terrible mess!!!

As you can tell from this little riff on the nomenclature, there is a good deal of confusion

about these so called “non-covalent” interactions. This confusion in name giving

represents a general lack of knowledge about how these interactions work in detail. So

instead of deriving these interactions from first principles, we will stick with more of a

qualitative description of these types of interactions.

Van der Waals interaction (dispersion energies)

Van der Waals interactions are probably the most basic type of interaction imaginable.

Any type of atom (as it turns out even macroscopic surfaces) experience Van der Waals

interactions. You get these interactions simply for being made of atoms.

The current understanding of Van der Waals interactions is a bit foggy and it is actually

very difficult to get very good tabulated numbers. Looking through the literature, one

gets the impression that if there is an interaction, but if all the other possible explanations

fail, then it must be Van der Waals interactions. In medical circles, I think this is referred

to as diagnosis by exclusion. As a matter of fact most of the literature on Van der Waals

interactions comes down to the non-ideal behavior of gasses that aught to be ideal (i.e.

noble gasses).

It turns out that if noble gasses are not extremely dilute, the ideal gas law does not really

work.

PV/T < R

The volume or the pressure is always a little lower or than you would expect from an

ideal gas. This indicates some type of process that holds the different atoms together. But

these gases certainly do not form covalent bonds with one another (full outer shell), they

also could not have electrostatic interactions, because they certainly are not ionic and

they do not have dipoles like water.

As it turns out though, the actual strength of the interactions and their distance

dependence is pretty accurately predicted by a theory of induced-induced dipole

interactions. As we will see even though VDW interactions are at their heart electrostatic,

their distance dependence is much steeper than that of Born or Coulomb effects.

What are induced dipoles?

Neutral atoms are made up of charged components

While noble gas atoms like most other atoms are not charged, they are of course made up

of positively and negatively charged parts: the nucleus and the electrons. The reason the

atom appears neutral to the outside, is because the electrostatic fields from the electrons

and the nucleus are co-centric.

External fields induce dipoles AND

attract them

If we now take such a neutral atom and we place it into an electro static field of a positive

charge then we will pull the electron to that charge and push the nucleus away from the

charge. The result is an induced dipole and this dipole, because the negatively charged

end of the dipole is closer to the external positive charge than the positive end of the

dipole, the induced dipole experiences a net attraction. Also notice that the interaction

energy depends on the distance between the centers of the positive and negative charges

of the dipole r.

van der Waals force

In physical chemistry, the van der Waals force (or van der Waals interaction), named after Dutch scientist Johannes Diderik van der Waals, is the sum of the attractive or repulsive forces between molecules (or between parts of the same molecule) other than those due to covalent bonds or to the electrostatic interaction of ions with one another or with neutral molecules. The term includes:

force between two permanent dipoles (Keesom force)

force between a permanent dipole and a corresponding induced dipole (Debye force)

force between two instantaneously induced dipoles (London dispersion force)

It is also sometimes used loosely as a synonym for the totality of intermolecular forces. Van der Waals forces are relatively weak compared to normal chemical bonds, but play a fundamental role in fields as diverse as supramolecular chemistry, structural biology, polymer science, nanotechnology, surface science, and condensed matter physics. Van der Waals forces define the chemical character of many organic compounds. They also define the solubility of organic substances in polar and non-polar media. In low molecular weight alcohols, the properties of the polar hydroxyl group dominate the weak intermolecular forces of van der Waals. In higher molecular weight alcohols, the properties of the nonpolar hydrocarbon chain(s) dominate and define the solubility. Van der Waals-London forces grow with the length of the nonpolar part of the substance.

Definition

Van der Waals forces include attractions between atoms, molecules, and surfaces. They differ from covalent and ionic bonding in that they are caused by correlations in the fluctuating polarizations of nearby particles (a consequence of quantum dynamics).

Intermolecular forces have four major contributions:

A repulsive component resulting from of the Pauli exclusion principle that prevents the collapse of molecules.

Attractive or repulsive electrostatic interactions between permanent charges (in the case of molecular ions), dipoles (in the case of molecules without inversion center), quadrupoles (all molecules with symmetry lower than cubic), and in general between permanent multipoles. The electrostatic interaction is sometimes called the Keesom interaction or Keesom force after Willem Hendrik Keesom.

Induction (also known as polarization), which is the attractive interaction between a permanent multipole on one molecule with an induced multipole on another. This interaction is sometimes called Debye force after Peter J.W. Debye.

Dispersion (usually named after Fritz London), which is the attractive interaction between any pair of molecules, including non-polar atoms, arising from the interactions of instantaneous multipoles.

Returning to nomenclature, different texts refer to different things using the term "van der Waals force". Some texts mean by the van der Waals force the totality of forces (including repulsion); others mean all the attractive forces (and then sometimes distinguish van der Waals-Keesom, van der Waals-Debye, and van der Waals-London); finally, some use the term "van der Waals force" solely as a synonym for the London/dispersion force. A common trend is that biochemistry and biology books, more frequently than chemistry books, use "van der Waals forces" as a synonym for London forces only.

All intermolecular/van der Waals forces are anisotropic (except those between two noble gas atoms), which means that they depend on the relative orientation of the molecules. The induction and dispersion interactions are always attractive, irrespective of orientation, but the electrostatic interaction changes sign upon rotation of the molecules. That is, the electrostatic force can be attractive or repulsive, depending on the mutual orientation of the molecules. When molecules are in thermal motion, as they are in the gas and liquid phase, the electrostatic force is averaged out to a large extent, because the molecules thermally rotate and thus probe both repulsive and attractive parts of the electrostatic force. Sometimes this effect is expressed by the statement that "random thermal motion around room temperature can usually overcome or disrupt them" (which refers to the electrostatic component of the van der Waals force). Clearly, the thermal averaging effect is much less pronounced for the attractive induction and dispersion forces.

The Lennard-Jones potential is often used as an approximate model for the isotropic part of a total (repulsion plus attraction) van der Waals force as a function of distance.

Van der Waals forces are responsible for certain cases of pressure broadening (van der Waals broadening) of spectral lines and the formation of van der Waals molecules. The London-van der Waals forces are related to the Casimir effect for dielectric media, the former being the microscopic description of the latter bulk property. The first detailed calculations of this were done in 1955 by E. M. Lifshitz.

Calculation

F(r)= \frac{\lambda}{r^s} - \frac{\mu}{r^t}

London dispersion force

Main article: London dispersion force

London dispersion forces, named after the German-American physicist Fritz London, are weak intermolecular forces that arise from the interactive forces between instantaneous multipoles in molecules without permanent multipole moments. London dispersion forces are also known as dispersion forces, London forces, or induced dipole–dipole forces. They increase with the molar mass, causing a higher boiling point especially for the halogen group.

Use by animals

http://upload.wikimedia.org/wikipedia/commons/thumb/1/16/Gecko_Leaftail_1.jpg/300px-Gecko_Leaftail_1.jpg

Gecko climbing glass using its natural setae

The ability of geckos - which can hang on a glass surface using only one toe - to climb on sheer surfaces has been attributed to van der Waals force, although a more recent study suggests that water molecules of roughly monolayer thickness (present on virtually all natural surfaces) also play a role. Efforts continue to create a dry glue that exploits this knowledge.

London dispersion force

http://upload.wikimedia.org/wikipedia/commons/thumb/0/04/Argon_dimer_potential.png/300px-Argon_dimer_potential.png

Interaction energy of argon dimer. The long-range part is due to London dispersion forces

London dispersion forces (LDF, also known as dispersion forces, London forces, induced dipole–induced dipole forces) is a type of force acting between atoms and molecules. They are part of the van der Waals forces. The LDF is named after the German-American physicist Fritz London.

The LDF is a weak intermolecular force arising from quantum induced instantaneous polarization multipoles in molecules. They can therefore act between molecules without permanent multipole moments.

London forces are exhibited by nonpolar molecules because of the correlated movements of the electrons in interacting molecules. Because the electrons from different molecules start "feeling" and avoiding each other, Electron density in a molecule becomes redistributed in proximity to another molecule, (see quantum mechanical theory of dispersion forces). This is frequently described as formation of "instantaneous dipoles" that attract each other. London forces are present between all chemical groups and usually represent main part of the total interaction force in condensed matter, even though they are generally weaker than ionic bonds and hydrogen bonds.

This is the only attractive intermolecular force present between neutral atoms (e.g., a noble gas). Without London forces, there would be no attractive force between noble gas atoms, and they wouldn't exist in liquid form.

London forces become stronger as the atom or molecule in question becomes larger. This is due to the increased polarizability of molecules with larger, more dispersed electron clouds. This trend is exemplified by the halogens (from smallest to largest: F2, Cl2, Br2, I2). Fluorine and chlorine are gases at room temperature, bromine is a liquid, and iodine is a solid. The London forces also become stronger with larger amounts of surface contact. Greater surface area means closer interaction between different molecules.

Quantum mechanical theory of dispersion forces

The first explanation of the attraction between noble gas atoms was given by Fritz London in 1930. He used a quantum mechanical theory based on second-order perturbation theory. The perturbation is the Coulomb interaction V between the electrons and nuclei of the two monomers (atoms or molecules) that constitute the dimer. The second-order perturbation expression of the interaction energy contains a sum over states. The states appearing in this sum are simple products of the excited electronic states of the monomers. Thus, no intermolecular antisymmetrization of the electronic states is included and the Pauli exclusion principle is only partially satisfied.

London developed the perturbation V in a Taylor series in \frac{1}{R}, where R is the distance between the nuclear centers of mass of the monomers.

This Taylor expansion is known as the multipole expansion of V because the terms in this series can be regarded as energies of two interacting multipoles, one on each monomer. Substitution of the multipole-expanded form of V into the second-order energy yields an expression that resembles somewhat an expression describing the interaction between instantaneous multipoles (see the qualitative description above). Additionally an approximation, named after Albrecht Unsöld, must be introduced in order to obtain a description of London dispersion in terms of dipole polarizabilities and ionization potentials.

In this manner the following approximation is obtained for the dispersion interaction E_{AB}^{\rm disp}between two atoms A and B. Here αA and αB are the dipole polarizabilities of the respective atoms. The quantities IA and IB are the first ionization potentials of the atoms and R is the intermolecular distance.

E_{AB}^{\rm disp} \approx -{3 \alpha^A \alpha^B I_A I_B\over 4(I_A + I_B)} R^{-6}

Note that this final London equation does not contain instantaneous dipoles (see molecular dipoles). The "explanation" of the dispersion force as the interaction between two such dipoles was invented after London gave the proper quantum mechanical theory. See the authoritative work for a criticism of the instantaneous dipole model and for a modern and thorough exposition of the theory of intermolecular forces.

The London theory has much similarity to the quantum mechanical theory of light dispersion, which is why London coined the phrase "dispersion effect".

Relative magnitude

Dispersion forces are usually dominant of the three van der Waals forces (orientation, induction, dispersion) between atoms and molecules, with the exception for molecules that are small and highly polar, like of water. The following contribution of the dispersion to the total intermolecular interaction energy has been given:

Contribution of the dispersion to the total intermolecular interaction energy

Molecule pair

Ne-Ne

CH4-CH4

HCl-HCl

HBr-HBr

96

HI-HI

CH3Cl-CH3Cl

NH3-NH3

H2O-H2O

HCl-HI

H2O-CH4

Mechanical bond

The mechanical bond is a type of chemical bond found in mechanically-interlocked molecular architectures such as catenanes and rotaxanes. Unlike classical molecular structures, interlocked molecules consist of two or more separate components which are not connected by chemical (i.e. covalent) bonds. These structures are true molecules and not a supramolecular species, as each component is intrinsically linked to the other – resulting in a mechanical bond which prevents dissociation without cleavage of one or more covalent bonds. “Mechanical bond” is a relatively new term and at this point has limited usage in chemical literature relative to more well established bonds, such as covalent, hydrogen, or ionic bonds.

Halogen bond

Halogen bonding (XB) is the non-covalent interaction that occurs between a halogen atom (Lewis acid) and a Lewis base. Although halogens are involved in other types of bonding (e.g. covalent), halogen bonding specifically refers to when the halogen acts as an electrophilic species.

Bonding

http://upload.wikimedia.org/wikipedia/commons/thumb/5/53/XBHBCompare.jpg/400px-XBHBCompare.jpg

http://bits.wikimedia.org/skins-1.5/common/images/magnify-clip.png

Figure 1: Comparison between hydrogen and halogen bonding. In both cases, D (donor) is the atom, group, or molecule that donates the electron poor species (H or X). H is the hydrogen atom involved in HB, and X is the halogen atom involved in XB. A (acceptor) is the electron rich species.

A parallel relationship can easily be drawn between halogen bonding and hydrogen bonding (HB). In both types of bonding, an electron donor/electron acceptor relationship exists. The difference between the two is what species can act as the electron donor/electron acceptor. In hydrogen bonding, a hydrogen atom acts as the electron acceptor and forms a non-covalent interaction by accepting electron density from an electron rich site (electron donor). In halogen bonding, a halogen atom is the electron acceptor. Electron density transfers results in a penetration of the van der Waals volumes.

http://upload.wikimedia.org/wikipedia/commons/thumb/c/cb/IodoChlorineAmine.jpg/80px-IodoChlorineAmine.jpg

http://bits.wikimedia.org/skins-1.5/common/images/magnify-clip.png

Figure 2: XB in complex between iodine monochloride and trimethylamine.

Halogens participating in halogen bonding include: iodine (I), bromine (Br), chlorine (Cl), and sometimes fluorine (F). All four halogens are capable of acting as XB donors (as proven through theoretical and experimental data) and follow the general trend: F < Cl < Br < I, with iodine normally forming the strongest interactions.

Dihalogens (I2, Br2, etc.) tend to form strong halogen bonds. The strength and effectiveness of chlorine and fluorine in XB formation depend on the nature of the XB donor. If the halogen is bonded to an electronegative (electron withdrawing) moiety, it is more likely to form stronger halogen bonds.

For example, iodoperfluoroalkanes are well-designed for XB crystal engineering. In addition, this is also why F2 can act as a strong XB donor, but fluorocarbons are weak XB donors because the alkyl group connected to the fluorine is not electronegative. In addition, the Lewis base (XB acceptor) tends to be electronegative as well and anions are better XB acceptors than neutral molecules.

Halogen bonds are strong, specific, and directional interactions that give rise to well-defined structures. Halogen bond strengths range from 5-180 kJ/mol. The strength of XB allows it to compete with HB, which are a little bit weaker in strength. Halogen bonds tend to form at 180° angles, which was shown in Odd Hassel’s studies with bromine and 1,4-dioxane in 1954. Another contributing factor to halogen bond strength comes from the short distance between the halogen (Lewis acid, XB donor) and Lewis base (XB acceptor). The attractive nature of halogen bonds result in the distance between the donor and acceptor to be shorter than the sum of van der Waals radii. The XB interaction becomes stronger as the distance decreases between the halogen and Lewis base.

Entropic force

A standard example of an entropic force is the elasticity of a freely-jointed polymer molecule: If the molecule is pulled into an extended configuration, the fact that more contracted, randomly coiled configurations are overwhelmingly more probable (i.e. have greater entropy) results in the chain eventually returning (through diffusion) to such a configuration. To the macroscopic observer, the precise origin of the microscopic forces that drive the motion is irrelevant: The observer simply sees the polymer contract into a state of higher entropy, as if driven by an elastic force.

Entropic forces also occur in the physics of gases and solutions, where they generate the pressure of an ideal gas (the energy of which depends only on its temperature, not its volume), the osmotic pressure of a dilute solution, and in colloidal suspensions, where they are responsible for the crystallization of hard spheres.

Hydrophobic force

A very frequently citedexample of an entropic force is the hydrophobic force. It originates from the entropy of the hydrogen bonded three-dimensional network of water molecules at room temperature. Since each water molecule is capable of donating two hydrogen bonds through the two protons and accepting two more hydrogen bonds through the two sp3 hybridized lone pairs, water molecules can form an extended three-dimensional network, unlike the case of hydrogen fluoride (which can accept 3 but donate only 1) or ammonia (which can donate 3 but accept only 1), which mainly form linear chains. Introduction of a non-hydrogen-bonding surface disrupts this network and the water molecules rearrange themselves around the surface so as to minimize the number of disrupted hydrogen bonds.

If the introduced surface had an ionic or polar nature, there would be water molecules standing more or less normal to the surface. But a non-hydrogen-bonding surface forces the surrounding hydrogen bonds to be tangential and they are locked in a clathrate-like basket shape. Water molecules involved in this clathrate-like basket around the non-hydrogen-bonding surface are constrained in their orientation. Thus, any event that would minimize such a surface is entropically favored. For example, when two such hydrophobic particles come very close, the clathrate-like baskets surrounding them merge, releasing some of the water molecules into the bulk of the water, leading to an increase in entropy. This is the basis of the so-called "attraction" between hydrophobic objects in water.

Hydrogen bond

http://upload.wikimedia.org/wikipedia/commons/thumb/2/2d/Hydrogen_Bond_Quadruple_AngewChemIntEd_1998_v37_p75.jpg/300px-Hydrogen_Bond_Quadruple_AngewChemIntEd_1998_v37_p75.jpg

An example of intermolecular hydrogen bonding in a self-assembled dimer complex reported by Meijer and coworkers.

http://upload.wikimedia.org/wikipedia/commons/thumb/9/9e/Acac.png/300px-Acac.png

Intramolecular hydrogen bonding in acetylacetone helps stabilize the enol tautomer

A hydrogen bond is the attractive interaction of a hydrogen atom with an electronegative atom, such as nitrogen, oxygen or fluorine, that comes from another molecule or chemical group. The hydrogen must be covalently bonded to another electronegative atom to create the bond. These bonds can occur between molecules (intermolecularly), or within different parts of a single molecule (intramolecularly). The hydrogen bond (5 to 30 kJ/mole) is stronger than a van der Waals interaction, but weaker than covalent or ionic bonds. This type of bond occurs in both inorganic molecules such as water and organic molecules such as DNA.

Intermolecular hydrogen bonding is responsible for the high boiling point of water (100 °C) compared to the other group 16 hydrides that have no hydrogen bonds. Intramolecular hydrogen bonding is partly responsible for the secondary, tertiary, and quaternary structures of proteins and nucleic acids. It also plays an important role in the structure of polymers, both synthetic and natural.

Refrences

Kimball's biology (textbook), 1995 ed.

Molecular Cell Biology (textbook), Lodish, Berk, Zipursky, Matsudaira, Baltimore, Darnell.


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