Types of Chemical Bonding

Several atoms are assembled and held together to form thousands of molecules which participate in the building and function of physical and biological systems. A bond is any force which holds two atoms together. The formation of bond between two atoms is due to some redistribution or regrouping of electrons to form a more stable configuration. The regrouping of electrons in the combining atoms may take place in either of the 3 ways: 

1. by a transfer of one or more electrons from one atom to another — electrovalent bonding 
2. by a sharing of one or more pairs of electrons between the combining atoms — covalent bonding
3. by a combination of the two processes of transfer and sharing — coordinate bonding.

Strength of bonds common in biomolecules

Type of Bonds

Covalent Bond

A covalent bond is formed when two atoms share one or more pair of electrons. Covalent bonds are strong bonds and very stable in nature. A covalent bond may also be termed a molecular bond. Covalent bonds form between two nonmetal atoms with identical or relatively close electronegativity values. This type of bond may also be found in other chemical species, such as radicals and macromolecules. 

Two important types of covalent bonds are-

1. Nonpolar or pure covalent bonds
2. Polar covalent bonds.

1. Nonpolar covalent bonds

Nonpolar bonds occur when atoms equally share electron pairs. Since only identical atoms (having the same electronegativity) truly engage in equal sharing, the definition is expanded to include covalent bonding between any atoms with an electronegativity difference less than 0.4. Examples of molecules with nonpolar bonds are H₂, N₂, and CH₄.

As the electronegativity difference increases, the electron pair in a bond is more closely associated with one nucleus than the other. If the electronegativity difference is between 0.4 and 1.7, the bond is polar. If the electronegativity difference is greater than 1.7, the bond is ionic.

A hydrogen molecule, H₂, consists of two hydrogen atoms joined by a covalent bond. Each hydrogen atom needs two electrons to achieve a stable outer electron shell. The pair of electrons is attracted to the positive charge of both atomic nuclei, holding the molecule together.

2. Polar covalent bonds

Polar covalent bonds may be characterized as transition state between ionic and covalent bonds. In this case, there is neither the complete transfer of electrons from one atom to the other nor equal sharing. When a covalent bond is formed between two atoms with different electronegativities, the electrons involved in the bond are not shared equally. The atom with higher electronegativity pulls the bonding electrons closer to it. In other words, the electron density in the molecular orbital would be greater around the atom with higher electronegativity. The result of this displacement of the molecular orbital toward the more electronegative atom will acquire a small negative charge and the less electronegative atom will acquire a small positive charge. For example, in the case of C-Cl bond, chlorine is more electronegative than carbon. As such, the electron density in the molecular orbital would be higher around chlorine atom than around the carbon atom. Thus, the chlorine atom acquires a small negative charge and the carbon atom acquires a small positive charge.

Non-covalent bonds

It includes ionic bonds, hydrogen bonds, van der Walls forces and hydrophobic interactions. There are weak interactions. Energy required to break non-covalent bonds is only 1-5 kcal/mol.

Ionic Bond

Ionic bonding is the complete transfer of valence electron(s) between atoms. It is a type of chemical bond that generates two oppositely charged ions. In ionic bonds, the metal loses electrons to become a positively charged cation, whereas the nonmetal accepts those electrons to become a negatively charged anion. Ionic bonds require an electron donor, often a metal, and an electron acceptor, a nonmetal. Ionic bond formation takes place between atoms of strongly electropositive and strongly electronegative elements. An element preceding an inert gas in the periodic table is strongly electronegative and the element immediately following the inert gas is strongly electropositive.

The predicted overall energy of the ionic bonding process, which includes the ionization energy of the metal and electron affinity of the nonmetal, is usually positive, indicating that the reaction is endothermic and unfavorable. However, this reaction is highly favorable because of the electrostatic attraction between the particles. At the ideal inter-atomic distance, attraction between these particles releases enough energy to facilitate the reaction. Most ionic compounds tend to dissociate in polar solvents because they are often polar. This phenomenon is due to the opposite charges on each ion.

Hydrogen Bonds

Hydrogen bonds can be formed between uncharged molecules as well as charged ones. In a hydrogen bond, a hydrogen atom is shared by two other atoms. The atom to which the hydrogen is more tightly linked is called the hydrogen donor, whereas the other atom is called as the hydrogen acceptor. The acceptor has a partial negative charge that attracts the hydrogen atom. In fact, a hydrogen bond can be considered as an intermediate in the transfer of a proton from an acid to a base. 

The donor in a hydrogen bond in biological systems is an oxygen or nitrogen atom that has a covalently-attached hydrogen atom. The acceptor is either oxygen or nitrogen. The bond energies of hydrogen bonds range between 2 and 7 kcal/ mol. Hydrogen bonds are stronger than the van der Waals but much weaker than covalent bonds. The strongest hydrogen bonds are those in which the donor, hydrogen and acceptor atoms are colinear. 

An important feature of hydrogen bonds is that they are highly directional. Hydrogen bonding is of 2 types: intramolecular (within a molecule) and intermolecular (between two molecules). Both types are common to many macromolecules such as proteins and DNA molecule. Intramolecular bonding gives rise to chelation, i.e., ring formation and this normally occurs only with the formation of 5-, 6- or 7- membered rings. Intermolecular bonding, however, gives rise to association, thereby raising the boiling point; it also raises the surface tension and the viscosity, but lowers the dielectric constant. It may exist in compounds in the liquid or solid state and its formation is very much affected by the shape of the molecular, i.e., by the steric factor.

Hydrophobic Interactions

The tendency of nonpolar molecules in a polar solvent (usually water) to interact with one another is called the hydrophobic effect. The interactions between the nonpolar molecules are called hydrophobic interactions. The relative hydrophobicity of amino acid residues is defined by a system known as hydrophobicity scales. The interactions between nonpolar molecules and water molecules are not as favorable as interactions amongst just the water molecules, due to the inability of nonpolar molecules to form hydrogen bonding or electrostatic interactions. When nonpolar molecules are introduced to the water molecules, the water molecules will initially surround the nonpolar molecules, forming a "cages" around the molecules. However, the tendency of nonpolar molecules to associate with one another will draw the nonpolar molecules together, forming a nonpolar aggregate.

Hydrophobic interactions are a major driving force in the folding of macromolecules, the binding of substrates to enzymes and the formation of membranes that define the boundaries of cells and their internal compartments. In macromolecules such as proteins, the acceptance or rejection by the aqueous environment of the hydrophilic and hydrophobic moieties respectively exerts a dominant influence on their final conformation. Here the nonpolar side chains of neutral amino acids tend to be closely associated with one another. The relationship is nonstoichiometric; hence no true bond may be said to exist. This clustering together of nonpolar molecules or groups in water is called hydrophobic interaction. The familiar sight of dispersed oil droplets coming together in water to form a single large oil drop is an analogous process.

Van der Waals Interactions

Van der Waals forces are driven by induced electrical interactions between two or more atoms or molecules that are very close to each other. Van der Waals interaction is the weakest of all intermolecular attractions between molecules. However, with a lot of Van der Waals forces interacting between two objects, the interaction can be very strong. van der Waals interactions (named after J. D. van der Waals) are weak, nonspecific, interatomic attractions and come into play when any two uncharged atoms are 3 to 4 Å apart. Though weaker and less specific than electrostatic and hydrogen bonds, van der Waals interactions are no less important in biological systems. The basis of a van der Waals bond is that the distribution of electronic charge around an atom changes with time.

All types of molecules exhibit van der Waals forces which arise from the attraction of the bound electrons of one atom for the nucleus of another. When two atoms are far apart, there is a very weak attraction which becomes stronger as the atoms move closer together. However, if the atoms move close enough for their outer electron shells to overlap, then a force of repulsion occurs. At a certain distance, defined as the van der Waals contact radius, there is a balance between the forces of attraction and those of repulsion. Each type of atom has a specific van der Waals contact radius. At this point of balance, the two atoms are separated by the van der Waals contact distance. The contact distance between an oxygen and carbon atom, for example, is 3.4 Å which is obtained by adding 1.4 and 2.0 Å, the contact radii of the O and C atoms, respectively. 

Dipole-Dipole Interaction

Dipole-Dipole interactions occur between molecules that have permanent dipoles; these molecules are also referred to as polar molecules. The figure below shows the electrostatic interaction between two dipoles.

Induced Dipoles

An induced dipole moment is a temporary condition during which a neutral nonpolar atom (i.e. Helium) undergo a separation of charges due to the environment. When an instantaneous dipole atom approaches a neighboring atom, it can cause that atom to also produce dipoles. The neighboring atom is then considered to have an induced dipole moment.
 
Even though these two atoms are interacting with each other, their dipoles may still fluctuate. However, they must fluctuate in synchrony in order to maintain their dipoles and stay interacted with each other. Result of synchronizing fluctuation of dipoles:

Spontaneous Dipole-Induced Dipole Interaction

Spontaneous dipole-induced dipole interactions are also known as dispersion or London forces (name after the German physicist Fritz London). They are large networks of intermolecular forces between nonpolar and non-charged molecules and atoms (i.e. alkanes, noble gases, and halogens). Molecules that have induced dipoles may also induce neighboring molecules to have dipole moments, so a large network of induced dipole-induced dipole interactions may exist. The image below illustrates a network of induced dipole-induced dipole interactions.
 

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