Unit 1: Biomolecules and Bioenergetics

Unit-1: Biomolecule and Bioenergenetics (Part-1) Chemical Bonds Each living and non-living object is made up of molecules. Further, each molecule is made up of atoms, the atoms being jointed by chemical bonds. These bonds are produced due to the redistribution and regrouping of electrons to form more stable configurations. Types of Chemical Bonds There are two types of chemical bonds such as: Primary bond Ionic or Electrovalent bond Covalent bonds a. Polar bond b. Polar covalent bond c. Coordinate covalent bond Secondary Bond Hydrogen bond Hydrophobic bond Disulfide bond Primary Bonds Primary bonds bind one atom to the other forming small molecules. They are mainly ionic & covalent bonds. The differential bonds are as follows: Ionic bond: The bond between two oppositely charged ions is known as an ionic bond. Generally, ionic bonds are present in inorganic salts which have definite crystalline structures. The formation of ionic bonds is based on Kossel's theory. The ionic bonds are produced due to the force of attraction between two oppositely charged ions. The strength of the bond is approximately proportional to the product of the charges on the ions. If transfer of more than one electron occurs, ionic bonds can be formed. Covalent Bonds: A covalent bond, also called a molecular bond, is a chemical bond that involves the sharing of electron pairs between atoms. These electron pairs are known as shared pairs or bonding pairs, and the stable balance of attractive and repulsive forces between atoms when they share electrons is known as a covalent bond. Covalent bonds are present in the case of non-metals having very high ionisation potentials. A covalent bond is formed by the formation of electrostatic attraction between the two atoms. A covalent bond may be formed by electron pairs distributed over two or more atoms; in this case, it is a shared electron pair. Covalent bonds are stable bonds which hold the adjacent atoms together. Example: In the case of a hydrogen molecule (H₂), two atoms of hydrogen are combined together by sharing their single electron with each other. ​ Generally, the molecules having covalent bonds are gases or liquids under normal conditions. However, a few molecules with high molecular weight having covalent bonds may be solids (for example, paraffin wax). Examples of covalent molecules: Polar Covalent Bonds: A polar covalent bond is a type of chemical bond where a pair of electrons is unequally shared between two atoms. In a polar covalent bond, the electrons are not equally shared because one atom spends more time with the electrons than the other atom. In a polar covalent bond, one atom has a stronger pull than the other atom and attracts electrons. Polar covalent bonds: When electrons spend more time with one atom, it causes that atom to carry a partial negative charge. The atom that does not spend as much time with the electrons carries a partial positive charge. Polar compounds exhibit properties intermediate between ionic & covalent compounds. Polar compounds have dipole attraction between the elements for which they have high melting & boiling points. Example: Bromine chloride: Cl atom has a higher attraction force for electrons than Br, for which the electrons are more dense toward Cl end than Br end. (Base molecule) Water (H₂O) molecule is also a polar compound having a polar covalent bond between 1 oxygen and 2 hydrogen atoms as shown below: (H₂O molecule) 4. Coordinate covalent bonds: A coordinate covalent bond (also known as a dative bond or a coordinate bond) is a kind of 2-center, 2-electron covalent bond in which the two electrons derive from the same atom. The bonding of a metal ion to ligands involves this kind of interaction. This type of interaction is central to Lewis theory. The coordinate covalent bonds are normally indicated by an arrow (→) with the head pointing towards the electron acceptor. Such types of bonds are found in complex compounds. Example: a) Ammonium boron trifluoride complex: Secondary bond: Secondary bonds are formed by primary bonds. However, in biological smaller molecules are formed by primary bonds and present mainly secondary systems. Some larger molecules and present mainly tertiary and quaternary foldings, for example, proteins. 1) Hydrogen bond: Hydrogen bonds are formed between the hydrogen atom of one molecule and oxygen or nitrogen or fluorine atom of another molecule. Such a kind of bonding is known as a hydrogen bond. Here a partial positive charge is developed in the hydrogen atom as oxygen or nitrogen or fluorine atoms have stronger affinity for the electrons of the hydrogen, for which these atoms (O/N/F) develop a partial negative charge (δ-). Hydrogen bonds are formed in substances which may be solid, liquid, or gas. A very good example is the presence of hydrogen bonds in water present in liquid or solid state. In water, a hydrogen bond is produced between the hydrogen atom of one water molecule and the oxygen atom of another water molecule. 2) Hydrophobic bonds: Hydrophobic bonds are generated between hydrophobic compounds. The force of attraction for such bonds is equal to that of van der Waals forces. 3. Disulphide bonds: A disulphide refers to a functional group with the structure R-S-S-R'. The linkage (S-S) is called an SS bond or sometimes a disulphide bridge and usually derived by the coupling of two thiol groups. Example: The three-dimensional structure of protein i.e., tertiary and quaternary structure, are stabilized at some sites by disulphide bridges (or bonds (S-S)). Human insulin is a dimer protein having 2 chains; A & B joined together by disulphide bonds. Significance of chemical bonds: Chemical bonds are responsible for holding atoms together in a molecule or compound. The definite crystalline structure of inorganic salts is due to ionic bonds. The very high ionization potential of non-metals is due to covalent bonds. Under normal conditions, macromolecules and gases or liquids states are soluble due to covalent bonds. The melting & boiling points of the molecule are high or low depending upon the covalent bond. The polarity or non-polarity of the molecule is due to polar covalent & non-polar covalent bonds. In the biological system, some molecules like proteins have secondary, tertiary, and quaternary folding which are stabilized by secondary bonds like hydrogen, hydrophobic & disulphide bonds. The double helix of DNA & secondary structure of RNA are stabilized by hydrogen bonds. The heat of fusion, heat of vaporization, melting & boiling point, viscosity, and surface tension of water are high due to the presence of hydrogen bonds between water molecules. In proteins, cation-anion balance is stabilized by the bond of disulphide bonds with other disulphide bridges. Structure of Water A water molecule has two hydrogen atoms and one oxygen atom being joined by covalent bonds having bond energy value of 110 Kcal. The hydrogen atom is positively charged (δ+) while the oxygen atom is negatively charged (δ-). Thus, water molecule exhibits electrical polarity and it behaves as an electrical dipole and hence can interact with other dipolar molecules. Due to this both negative & positive charge, water molecule is electrically neutral. The distance of a covalent bond is 0.99Å. The bond angle between the covalent bond is 105° rather than the tetrahedral angle of 109° 28'. The shape of the water molecule (caused by angle) is bent or V-shaped (non-linear). The number of electrons from atoms, which may be shared during a covalent bond formation & a covalency of that atom, e.g., the covalency of oxygen is 2 and that of hydrogen is 1. Molecules possessing covalent bonds are generally found in gases and liquids under normal conditions. Covalent bonds are comparatively weaker than ionic bonds. Properties of Water Melting point: The melting point of water is 0°C, which is higher than most compounds and common gases like ammonia, hydrogen sulphide, etc. At melting point, water is converted from solid state (ice) to liquid state (water) and vice versa. The high melting point/freezing point of water is due to strong intermolecular forces acting between multiple water molecules. Boiling point: The boiling point of water is 100°C, which is higher in comparison with most other substances. At boiling point, water is converted into gaseous state from liquid state and vice versa. Specific heat capacity: The specific heat capacity of water is defined as the number of calories required to raise the temperature of one (1) gram of water from 15°C to 16°C. Sp. H. C. Sp. H. C. for water, which is 1, is the highest among all other compounds (except liquid ammonia). Heat of fusion: The heat of fusion may be defined as the number of calories (cal) or energy that should be given out by one gram of a substance to convert it from liquid state to solid state or vice versa. In other words, the number of calories required to convert one gram of a substance from solid state to liquid state at its melting point. The heat of fusion of water is 80 cal/g. Heat of vaporization: The latent heat of vaporization is defined as the number of calories per gram required to overcome the attractive forces between adjacent molecules in a liquid as they convert to the gaseous state. The heat of vaporization of water is 540 cal/g. Density: Water has the highest density at 4°C due to the packing of small association groups. Hence, ice has less density and floats in water. Surface tension Water has a very high surface tension. Due to capillarity, water rises up in a glass tube of about 0.03 mm (fine diameter) to a height of 12 cm. This property of water enables it to move extensively through narrow cavities in the soil or in plant cell walls. Viscosity and tensile strength Water has high viscosity and high tensile strength. This property of water is also very significant for water movement through plants. Polarity and solvent property Water is a highly polar solvent. Its molecules have relatively high tendency to attract each other electrostatically. In comparison to most other common liquids, water is the best solvent. Most of the salts i.e., sodium chloride, calcium chloride, etc. are readily dissolved in water. Water has a very high dielectric constant i.e., 80. Transmission of visible light Water has high transmission of visible light waves and is useful for submerged aquatic plants in carrying out photosynthesis. Above physical properties of water make it an ideal medium for conducting various activities and processes of life. Here is the transcribed text from the image: The Buffers A solution of known pH which has some reserve power of neutralizing acids and bases so that it maintains a steady value of the pH is called a "buffer solution." Examples of buffered solution: A weak acid and a salt of it with a strong base. e.g., acetic acid (CH3COOH) and sodium acetate (CH3COONa) A weak base and a salt of it with a strong acid. e.g., NH4OH and NH4Cl A buffer solution has a definite pH value. Its pH value does not alter even on keeping for long or on dilution. Its pH value is slightly altered by the addition of a strong acid or base. The molar concentration of the buffer components by one liter (The buffer's capacity is directly proportional to the concentration of the buffer components.) The effectiveness of a buffer solution is the ratio of the concentration of the conjugate base to the concentration of the weak acid i.e., [weak acid] [conjugate base] ​ Laws of Thermodynamics There are three laws of thermodynamics, such as; First law: It states that when energy passes as work, as heat, or with matter, into or out of a system, the system's internal energy changes in accord with the law of conservation of energy. The first law of thermodynamics does not explain how a process occurs but is concerned with the initial and final state of the system and thus the difference between them. The mathematical expression is: ​ = change in energy of the system ​ = energy of a system at the beginning of the process ​ = energy of a system at the end of the process 𝑄 Q = heat absorbed or released by the system 𝑊 W = work done by the system 2nd Law: In a natural process thermodynamics favors the sum of the entropies of the interacting thermodynamic systems to increase. The total entropy of a system must increase for the process to occur spontaneously. 3rd Law: The entropy of a system approaches a constant value as the temperature approaches absolute zero. With the exception of non-crystalline solids (glassy), the entropy of a system at absolute zero is typically close to zero. Zeroth Law of Thermodynamics: If two systems are each in thermal equilibrium with a third system, they are in thermal equilibrium with each other. Endergonic and Exergonic Reactions: Endergonic Reaction: An endergonic reaction may also be called an unfavorable reaction or non-spontaneous reaction. The reaction requires more energy than it gives off. Endergonic reactions absorb energy from their surroundings. This means that one or more of the reactants are weaker than the chemical bond that is broken. The free energy of the system increases. The change in the standard Gibbs free energy ( ΔG) of an endergonic reaction is positive (ΔG>0). The change in entropy ( ΔS) decreases. Examples of endergonic reactions include endothermic reactions such as photosynthesis, the melting of ice into liquid water, ATP synthesis, and anabolic synthesis, etc. ΔH=+ve