CHEMICAL BONDING-COVALENT BONDING
2.1. Covalent bond
It is a chemical bond that involves the sharing of electrons between atoms. This is the strongest force. It is the bond which held the atoms together. A typical carbon-carbon (C-C) covalent bond has a bond length of 1.54 Å and bond energy of 85 kcal mol-1 (356 kJ mol-1).
Example of a covalent bond
A non-polar covalent bond is a chemical bond in which two atoms share a pair of an electron with each other.
the eg.-Non polar covalent bond is found in methane (CH4). The Lewis structure shows the electrons which are shared between C and H atoms.
Polar covalent bonding is a chemical bond where a pair of electron is equally shared between two atoms. It has significant ionic character. It means that the two shared electrons are closer to one of the atoms than the other, creating an imbalance of charge.
eg. Polar covalent bond is water (H2O)
2.2 Non-covalent bonding
Non-covalent bonds do not involve the sharing of electrons between atom like in covalent bond. It is a bond that is found between two macromolecules
It is a reversible interaction of biomolecules. Four types of noncovalent bonds are formed between macromolecules. The four fundamental noncovalent bonds are electrostatic interactions, hydrogen bonds, hydrophobic bond and Vander Walls interactions. They differ in geometry, strength, and specificity. These bonds are greatly affected in a different manner in the presence of water.
2.2.1. Electrostatic interactions
An electrostatic interaction depends on the electric charges on atoms. An ionic bond is a type of chemical bond formed through an electrostatic attraction between two oppositely charged ions. Ionic bond is formed due to the transfer of an electron from one donor atom to another acceptor atom like sodium atom to fluorine. In general, electrostatic interaction takes place between two stationary charges. The energy of an electrostatic interaction is given by Coulomb's law:
where E is the energy, q1 and q2 are the charges on the two atoms (in units of the electronic charge), r is the distance between the two atoms (in angstroms) and k is a constant.
2.2.2 Hydrogen bonds
Four conditions need to be fulfilled for hydrogen bonding
1. Hydrogen is a must for hydrogen bonding
2. The hydrogen must be sandwich between
two highly electronegative atom
3. The distance between the electronegative
atom and hydrogen must be 2.8 A°
4. The bond must be in a plane.
Hydrogen bonds are relatively weak interactions, which are crucial for biological macromolecules such as DNA and proteins. The solubility of any molecule in water is also dependent upon hydrogen bonding. Hydrogen bonds are fundamentally electrostatic interactions. Normally hydrogen bonding is formed by fluorine, oxygen and nitrogen. The electronegative atom (like F, O, N) to which the hydrogen atom is covalently bonded pulls shared electron density away from the hydrogen atom because of this the hydrogen develops a partial positive charge (+) and the electronegative atom develops partial negative charge (–). Thus, the covalently bonded hydrogen of one molecule can interact with another atom having a partial negative charge (–) through an electrostatic interaction.
The large difference in electronegativities between hydrogen and any of fluorine, nitrogen and oxygen, coupled with their lone pairs of electrons cause strong electrostatic forces between molecules. Hydrogen bonds are responsible for the high boiling points of water and ammonia with respect to their heavier analogues.
2.2.3. Van der Waals interactions.
The basis of Van der Waals interaction is that the distribution of electronic charge around an atom changes with time. The distribution changes with time are because of the movement of electrons in the orbit. The electron is not static, it revolves around the orbit.
Due to the movement of an electron at any instant, the charge distribution is not perfectly symmetric. This transient asymmetry in the electronic charge around an atom acts through electrostatic interactions to induce a complementary asymmetry in the electron distribution around its neighbouring atoms.
The resulting attraction between two atoms increases as they come closer to each other. The minimum distance between the molecule is known as Vander Walls contact distance. At a shorter distance, very strong repulsive forces become dominant because the outer electron clouds overlap to each other. Thus atoms repel to each other. That's the reason Vander Walls forces is the sum of attractive and repulsive forces.
Three types of Van der Walls forces
2. Dipole - induced Dipole
3. London - forces
London dispersion force :
A type of Van der Walls force is known as the London dispersion force. The London dispersion force arises due to instantaneous dipoles in neighbouring atoms. As the negative charge of the electron is not uniform around the whole atom, there is always a charge imbalance. This small charge will induce a corresponding dipole in a nearby molecule; causing an attraction between the two. The electron then moves to another part of the electron cloud and the attraction is broken. London dispersion force is the weakest intermolecular force. London dispersive force is a temporary attractive force that results when electrons in two atoms occupy positions that make the atoms form temporary dipoles. This force is also known as induced-dipole-induced dipole attraction.
2.2.4. Hydrophobic interaction
Hydrophobic force is not a force, is an interaction of non-polar molecules like lipids which come close to each other only in presence of water. The hydrophobic force is a tendency of water molecules to form a bigger cage around the lipid molecules. The hydrophobic force does not exist without water. It is a tendency of water molecule only. Hydrophobic is generally non-polar molecules and have a long chain of carbon molecule which does not interact with water molecules.
- Book COVER AND ABOUT US
- CHEMICAL BONDING
- AMINO ACIDS
- PROTEIN STRUCTURE
- RAMACHANDRAN PLOT
- PROTEIN STABILITY
- KINETIC ANALYSIS
- REGULATION OF GLYCOLYSIS
- TRICARBOXYLIC ACID CYCLE (TCA CYCLE)
- REGULATION OF THE CITRIC ACID CYCLE
- GLYOXYLATE CYCLE OR KREBS KORNBERG CYCLE
- ELECTRON-TRANSPORT CHAIN
- MECHANISMS OF OXIDATIVE PHOSPHORYLATION
- PENTOSE PHOSPHATE PATHWAY
- LIPID METABOLISM
- FATTY ACID OXIDATION
- DNA STRUCTURE
- NUCLEOTIDE BIOSYNTHESIS