Biochemistry Lecture: Water PDF

Biochemistry Lecture: Water  Water : The Solvent for Biochemical Reactions 2  OBJECTIVES: At the end of the session, students should be able to: 1. Describe properties of water and buffers 2. Compute the pH using Henderson...

Biochemistry Lecture: Water  Water : The Solvent for Biochemical Reactions 2  OBJECTIVES: At the end of the session, students should be able to: 1. Describe properties of water and buffers 2. Compute the pH using Henderson Hasselbach equation  Importance & functions: ⚫Constitutes 70-85% of cell weight ⚫It is a major component of many body fluids including blood, urine, and saliva. ⚫It is the solvent in which most biological transformations take place ⚫It is needed for transporting substances, across membranes, maintaining body temperature, dissolving waste products for excretion and producing digestive fluids 4  Molecular Structure of water ⚫Water molecule consists of two hydrogen atoms bonded to one oxygen atom by polar covalent bonds. ⚫Polar covalent bond – when electrons are unequally shared due to difference in electronegativity of atoms involved in bond ⚫water is a "polar" molecule. 5  Solvent Properties of Water Water is a good solvent ⚫The polar nature of water allows it to: a. Dissolve ionic compounds easily. ⮚ The positive side of water surrounds negatively charged molecules, and the negative side of water surrounds positively charged molecules. ⮚ Ion-dipole interaction- an attractive force between: A charged ion (cation or anion) b. Dissolve low-molecular-weight polar covalent compounds (e.g., C2H5OH and CH3COCH3) ⮚ Dipole-dipole interaction 6 Hydration Shells Surrounding Ions in Water 7  Solvent Properties of Water Ion–dipole and dipole–dipole interactions help ionic and polar compounds dissolve in water 8  Solvent Properties of Water ⚫However, water cannot dissolve nonpolar molecules (e.g. oils) because they lack dipoles. Hydrophobic - water-fearing ⮚ nonpolar molecules that do not dissolve in water Hydrophilic - means water-loving ⮚ Molecules that dissolve readily in water ⚫Amphipathic – A single molecule may have both polar (hydrophilic) and nonpolar (hydrophobic) portions, e.g. sodium palmitate (C16H31NaO2) a skincare and cosmetic ingredient used to saponify oils and fats to create soaps. 9  Amphipathic Molecule ⚫Example is a long-chain fatty acid having a polar carboxylic acid group and a long nonpolar hydrocarbon portion 10 Micelle Formation by Amphipathic Molecule ⚫Micelle – a spherical arrangement of organic molecules in water solution clustered so that their hydrophobic parts are buried inside the sphere and their hydrophilic parts are on the surface of the sphere, in contact with the water environment. 11 Examples of Amphipathic Biomolecules 12 Examples of Hydrophilic Biomolecules 13 Examples of Hydrophobic Biomolecules 14  Hydrogen Bonding ⚫Hydrogen Bonding can be considered a special case of dipole–dipole interaction. ⚫When hydrogen is covalently bonded to a very electronegative atom such as oxygen or nitrogen, it has a partial positive charge that can interact with an unshared (nonbonding) pair of electrons on another electronegative atom, most commonly O or N 15  Hydrogen Bonding ⚫Hydrogen bonding between polar groups and water. 16 Other Biologically Important Hydrogen bonds ⚫Hydrogen bonding is important in stabilization of 3- D structures of biological molecules such as: DNA, RNA, proteins. 17  Acids, Bases and pH ⚫ Ionization of water can be considered as proton transfer from one water molecule to another forming a hydronium ion (H3O+) and a hydroxide ion (OH-). ⚫Ionization process is reversible. ⚫For simplicity, the (H3O+) is often written as H+. H2O + H2O H3O+ + OH- 18  Acids, Bases and pH Ionization of water – reversible self –dissociation ⮚ Can be described by the following equilibrium: ⚫To express the extent of ionization quantitatively: ⚫Keq - equilibrium constant -defined as the ratio of the concentrations of the products and reactants ⮚ concentration are Molarity (M) = moles/L 19  Acids, Bases and pH Keq for pure water = 1.8 x 10- 16 M at 25ºC Concentration of pure H2O, [H2O] = 55.5M [H+][OH-] = 1.0 x 10- 14 M = Kw Kw is ion product constant for water ⚫Based from the chemical equation for dissociation, concentration of H+ and OH- in pure water must be equal : 20  Acids, Bases and pH. “potential of hydrogen" or "power of hydrogen" 21 pH of common substances 22  Examples Calculate the pH of the following solutions: 1. 0.035 M hydrochloric acid 2. 0.155 M sulfuric acid 3. 2.3 x 10-4 M NaOH 23  Acids, Bases and pH ⚫Acid – a molecule that behaves as a proton (hydrogen ion) donor ⚫Base – a molecule that behaves as a proton acceptor ⚫Water can act as an acid and a base -> amphoteric ⚫The number of hydrogen ions, [H+], present in a solution is a measure of the acidity of the solution ⚫The number of hydroxyl ions, [OH-], present in a solution is a measure of the alkalinity of the solution 27  Acids, Bases and pH ⚫Strong acids: ⮚ defined as a substance that has a greater tendency to lose its proton and completely dissociates (or ionizes) in water, such as HCl, HNO3 and H2SO4 ⚫ Weak acids: ⮚ a molecule that has a lesser tendency to lose its proton and does not readily dissociate in water, such as CH3COOH. 28  Acids, Bases and pH ⚫Acid Strength -defined as its tendency to release a proton (dissociate) ⮚ amount of hydrogen ion released when a given amount of acid is dissolved in water ⚫ dissociation for an acid: ⚫Conjugate Base = base formed by the removal of a proton from an acid 29  Acids, Bases and pH ⚫Ka = called the acid dissociation constant ⮚ a QUANTITATIVE measure of acid strength ⚫Smaller pKa -> stronger acid ⚫Larger pKa -> Weaker acid (stronger base) 30  Acids, Bases and pH ⚫Strong Acids: ⮚ Ka value is large, pKa is small ⮚ Ex. : nitric acid (HNO3), hydrochloric acid(HCl) ⮚ [H+] concentration is approx. equal to the concentration of acid in solution *In 4.0mM solution of HCl, [H+] = 4.0mM pH = –log [H+] = –log(4.0x103) pH = 2.4 31  Acids, Bases and pH ⚫Weak Acids: ⮚ Not completely dissociated in water -[H+] is lower than [HA] ⮚ Common in biological systems : -Acetic acid, phosphoric acid, carbonic acid, and lactic acid ⚫Henderson–Hasselbalch equation – equation that describes the behavior of weak acids in solution ⮚ Allows us to calculate the concentration of an acid and conjugate base at various pH 32 Henderson–Hasselbalch equation 33  Henderson–Hasselbalch equation From the equation, ⮚ when the concentrations of weak acid and its conjugate base are equal, the pH of the solution equals the pKa of the weak acid ⮚ when pH < pKa, the weak acid predominates ⮚ when pH > pKa, the conjugate base predominates 34 pH versus enzymatic activity 37  Buffers ⚫Buffers are solutions that resist changes in their pH when moderate amounts of acids or bases are added. ⚫Buffers are made with a weak acid and a soluble salt containing the its conjugate base or a weak base and a soluble salt containing its conjugate acid ⚫Examples of acid-base buffers are solutions containing ⮚ CH3COOH and CH3COONa ⮚ H2CO3 and NaHCO3 ⮚ NaH2PO4 and Na2HPO4 ⮚ NH3 and NH4Cl ⚫Strong acids will not buffer because it dissociates completely in water 38 Buffers 39  Mechanism of action of buffer acetic acid/ acetate system: 1. Addition of strong acid ⮚ strong acid reacts with acetate forming acetic acid (weaker acid) ⮚ minimal change in pH will occur. 40  Mechanism of action of buffer acetic acid/ acetate system: 2. Addition of strong base ⮚ strong base reacts with acetic acid forming acetate (conjugate base) ⮚ minimal change in pH will occur. 41  Buffer Capacity ⚫two factors that determine the effectiveness or capacity of buffer solution a. Molar concentration of the buffer components ⮚ the greater the concentration of the buffer components, the greater the buffer capacity b. Relative concentrations of the conjugate base and the weak acid ⮚ most effective buffer would be one with equal concentrations of acidic and basic components 42 Buffer Capacity 43 ⚫For carbonic acid, pKa = 6.37 : Effective buffering range = 5.77 - 7.37 44  Some Biological Buffer system ⚫Phosphate buffer system - principal buffer in cells ⮚ H2PO4 – (weak acid ) and HPO4-2 (conjugate base) pair ⚫Bicarbonate buffer system – important buffer in blood ⮚ carbonic acid (H2CO3) and bicarbonate (HCO3-) carbonic acid is formed from dissolved carbon dioxide and water ⮚ [HCO3−]/[H2CO3] ratio of 10 :1 is required for the pH of blood plasma to remain 7.40 ⮚ Acidosis is a serious condition where the pH of blood is below a pH of 7.35 ⮚ Alkalosis is a serious condition where the pH of blood is above a pH of 7.45 Blood pHs below 6.9 or above 7.9 are usually fatal if they last for more than a short time 45  Selecting a Buffer Criteria for selecting a buffer for a biochemical reaction 1. Suitable pKa for the buffer. 2. No interference with the reaction or with the assay. 3. Suitable ionic strength of the buffer. 4. No precipitation of reactants or products due to presence of the buffer. 5. Non-biological nature of the buffer. 46 …..thank you! 51