Biophysical Chemistry 2023-2024 PDF

THE UNIVERSITY OF ZAMBIA SCHOOL OF MEDICINE DEPARTMENT OF PHYSIOLOGICAL SCIENCES PGY 2320 (MEDICAL BIOCHEMISTRY & GENETICS) 2023/2024 Biophysical Chemistry Dr. Lubinda Mukololo Email: [email protected] Objectives: Properties o...

THE UNIVERSITY OF ZAMBIA SCHOOL OF MEDICINE DEPARTMENT OF PHYSIOLOGICAL SCIENCES PGY 2320 (MEDICAL BIOCHEMISTRY & GENETICS) 2023/2024 Biophysical Chemistry Dr. Lubinda Mukololo Email: [email protected] Objectives: Properties of water as a solvent for biochemical reactions Hydrogen bonding Weak acids and bases pH and pKa determination Buffers and their role in biological systems Physiologically important buffer systems 2 Why is water such an important molecule in living systems? Serves as a solvent/medium for reactions Participates in many biological reactions, being a principal reactant in processes such as photosynthesis and the final product in aerobic energy metabolism Determines the shape of most biomolecules, since intra and intercellular environments for all organisms are aqueous Stabilises the temperature and pH of living organisms 3  Properties of Water 1. Polarity Covalent bonds (electron pair is shared) between oxygen and hydrogen atoms with a bond angle of 104.5o Oxygen atom is more electronegative than hydrogen atom --> electrons spend more time around oxygen atom than hydrogen atom --> result is a POLAR covalent bond. 4  Creates a permanent dipole in the molecule Electric dipole Can determine relative solubility of molecules “like dissolves like”. 5 2. Hydrogen bonds Due to polar covalent bonds --> attraction of water molecules for each other. Creates hydrogen bonds = attraction of one slightly positive hydrogen atom of one water molecule and one slightly negative oxygen atom of another water molecule. 6  Each water molecule can form hydrogen bonds with four other water molecules. The length of the bond is about twice that of a covalent bond. Weaker than covalent bonds (about 25x weaker). 7  Hydrogen bonds give water a high melting point. Density of water decreases as it cools; q water expands as it freezes q ice results from an open lattice of water molecules q less dense, but more ordered. Hydrogen bonds (shown as dashed lines) are formed between water molecules. 8  Hydrogen bonds contribute to water’s high specific heat (amount of heat needed to raise the temperature of 1 gm of a substance 1oC). This is due to the fact that hydrogen bonds must be broken to increase the kinetic energy (motion of molecules) and temperature of a substance. Water has a high heat of vaporization - large amount of heat is needed to evaporate water because hydrogen bonds must be broken to change water from liquid to gaseous state. 9 3. Universal solvent Water can interact with and dissolve other polar compounds and those that ionize (electrolytes) because they are hydrophilic. Do so by aligning themselves around the electrolytes to form solvation spheres - shell of water molecules around each ion. Solubility of organic molecules in water depends on polarity and the ability to form hydrogen bonds with water. 10 11  Functional groups on molecules that confer solubility: carboxylates protonated amines amino hydroxyl carbonyl As the number of polar groups increases in a molecule, so does its solubility in water. 12  4. Hydrophobic interactions Nonpolar molecules are not soluble in water because water molecules interact with each other rather than nonpolar molecules --> nonpolar molecules are excluded and associate with each other (known as the hydrophobic effect and the associated interactions are called hydrophobic interactions). Nonpolar molecules are hydrophobic. 13 Hydrophilic/ Hydrophobic Interactions 14 Hydrophobicity/ Micelles 15  Molecules such as detergents or surfactants are amphipathic (have both hydrophilic and hydrophobic portions). Usually have a hydrophobic chain of 12 carbon atoms plus an ionic or polar end. Soaps are alkali metal salts of long chain fatty acids; examples of detergents include: sodium palmitate, sodium dodecyl sulfate (synthetic detergent). 16  All form micelles (spheres in which hydrophilic heads are hydrated and hydrophobic tails face inward. Used to trap grease and oils inside to remove them. 17 Noncovalent Interactions in Biomolecules There are four major noncovalent forces involved in the structure and function of biomolecules: i. Hydrogen bonds q More important when they occur between and within molecules q stabilize structures such as proteins and nucleic acids. ii. Hydrophobic interactions q Very weak q Important in protein shape and membrane structure. 18. iii Charge-charge Interactions or Electrostatic Interactions (Ionic Bonds) q Occur between two oppositely charged particles. q Strongest noncovalent force that occurs over greater distances. q Can be weakened significantly by water molecules (can interfere with bonding). 19 iv. van der Waals forces qarise from attraction between transient dipoles generated by the rapid movement of electrons on all neutral atoms qCan be attractive or repulsive depending upon the distance of the two atoms. qMuch weaker than hydrogen bonds. qThe actual distance between atoms is the distance at which maximum attraction occurs. qDistances vary depending upon individual atoms. 20 5. Nucleophilic Nature of Water Chemicals that are electron-rich (nucleophiles) seek electron-deficient chemicals (electrophiles) Nucleophiles are negatively charged or have unshared pairs of electrons --> attack electrophiles during substitution or addition reactions. Examples of nucleophiles: oxygen, nitrogen, sulfur, carbon, water (weak). 21  Important in condensation reactions, where hydrolysis reactions are favored. e.g. protein ----> amino acids Condensation reactions usually use ATP and exclude water to make the reactions more favorable. In the cell, these reactions actually only occur in the presence of hydrolases 22  6. Ionization of water Pure water ionizes slightly to yield a hydrogen ion (proton) and a hydroxide ion Hence can act as an acid (proton donor) or base (proton acceptor). 23  6. Ionization of water Hydrogen ions formed in water are immediately hydrated to hydronium ions: 2H2O ---> H3O+ + OH-, but usually written H2O ---> H+ + OH- Protons exist in solution not only as H3O+, but also as multimers such as H5O2+ and H7O3+ ; but for simplicity, usually represented as H+. The actual probability of a hydrogen atom in pure water existing as a hydrogen ion is approximately 1.8 × 10−9. 24  Equilibrium constant for water: Keq = [H+][OH-] = 1.8 x 10-16M at 25oC [H2O] In pure water at 25⁰C, the concentration of water is 55.5 M; i.e. 1 liter of H2O is 1000 g, 1 mole of H2O is 18 g Can rearrange equation to the following: 25  1.8 x 10-16M(55.5 M) = [H+][OH-] = Kw (ion product of water) 1.0 x 10-14M2 = [H+][OH-] At equilibrium, [H+] = [OH-], so 1.0 x 10-14M2 = [H+]2 = Kw 1.0 x 10-7 = [H+] The concentration of H+ in a solution of 0.1M NaOH would therefore be calculated as: 26  pH pH is a measure of the concentration of H+ in a solution. It is given by: pH = - log [H+] pH 7 is basic or alkaline 1 change in pH units equals a 10-fold change in [H+] Note that: PKW= PH + POH = 14 27  Weak Acids and Bases Unlike strong acids and bases, weak acids and bases ionise only slightly in water (i.e. they do not ionise completely) Have characteristic dissociation constants Common in biological systems and play important roles in metabolism and regulation Acids may be defined as proton donors and bases as proton acceptors 28  A proton donor and its corresponding proton acceptor make up a conjugate acid-base pair Acetic acid (CH3COOH) and the e.g acetate anion (CH3COO-) constitute a conjugate acid-base pair Each acid has a characteristic tendency to lose its proton (H+) in an aqueous solution and the stronger the acid, the greater its tendency to lose its proton. 29  pKa The pKa of an acid is the pH at which the acid is half dissociated, when [A-]=[HA]. pKa is analogous to pH and is defined by the equation: The stronger the acid, the lower its pKa. 30  The Henderson-Hasselbalch Equation Describes the dissociation of a weak acid in the presence of its conjugate base. The ionization of a weak acid (HA) in water is given by: ……….equation 1 The tendency of a weak acid (HA) to lose a proton (H+) and form its conjugate base (A-) is defined by the equilibrium constant (Keq) for the above reversible reaction: ……….equation 2 where Ka is the dissociation constant of an acid. 31  Rearranging this expression in terms of the parameter of interest, (H+), we have: …………..equation 3 Taking logarithm of both sides gives: ……equation 4 Multiplying both sides by -1 and defining pKa = −logKa we have: 32  pH = p𝐾𝑎 − log [HA] − ……….equation 5 [A ] or [𝐀− ] pH = 𝐩𝑲𝒂 + 𝐥𝐨𝐠 [𝐇𝐀] , the Henderson -Hasselbalch equation The pH of a solution can be calculated from this equation if the molar proportion of A- to HA and the pKa of HA are known. Conversely, the pKa of an acid can be calculated if the molar proportion of A- to HA and the pH of the solution are known. Stronger acids, such as phosphoric and carbonic acids, have larger dissociation constants; weaker acids, such as monohydrogen phosphates have smaller dissociation constants 33  Example 1 What is the pH of a solution of 0.1 M acetic acid and 0.2 M acetate ion given that the pKa of acetic acid is 4.8? 34  Example 2 What is the pH of a solution whose hydroxide ion concentration is 4.0 × 10-4 mol/L? First define a quantity pOH that is equal to −log [OH-] and that may be derived from the definition of Kw, K w = [H+][OH-] = 10-14 Therefore, log [H+] + log [OH-] = log 10-14 or pH + pOH = 14 To solve the problem by this approach: 35 36  Buffers: Solutions of Weak Acids & Their Bases resist changes in pH. Buffers are aqueous systems that tend to resist changes in pH when small amounts of Acid (H+) or base (OH-) are added. A buffer system consists of a weak acid and its conjugate base (more common) or a weak base and its conjugate acid (less common). 37 Importance of buffers in the body Most enzymes that catalyze biochemical reactions are very sensitive to pH The intracellular and extracellular fluids of multicellular organisms have a near constant pH The body fluids must be protected against changes in pH all the time as acids and bases are continuously being produced by metabolic processes 38  The organism’s first line of defense against changes in internal pH is provided by buffer systems. Biological buffers include HCO3–, phosphate and proteins which accept or release protons to resist a change in pH eg – Proteins may contain amino acids with functional groups which may act as weak acids and bases – Phosphate groups of nucleotides function as weak acids – Etc 39 The phosphate buffer system It is a major intracellular buffering system Acts in the cytoplasm of all cells Consists of H2PO4- (dihydrogen phosphate ion) as proton donor and HPO42- (hydrogen phosphate ion) as proton acceptor The phosphate buffer system is maximally effective at a pH close to its pKa of 6.86 and thus tends to resist pH changes in the range between about 5.9 and 7.9 40  The bicarbonate buffer system Blood plasma is in part buffered by the bicarbonate buffer system consisting of carbonic acid (H2CO3) as proton donor and bicarbonate (HCO3-) as proton acceptor H2CO3 H+ + HCO3- This buffer system is more complex than other conjugate acid-base pairs because one of the components, carbonic acid is formed from dissolved (d) CO2 and water in a reversible reaction. 41 CO2(d) + H2O H2CO3 Dissolved CO2 is in equilibrium with CO2 of the gas phase CO2(g) CO2(d) 42  The pH of a bicarbonate buffer system depends on the concentration of H2CO3 and HCO3–, the proton donor and acceptor respectively. The concentration of H2CO3 depends on the concentration of dissolved CO2, which in turn depends on the concentration of CO2 in the gas phase (partial pressure of CO2) 43 44  Human blood plasma normally has a pH of about 7.4 In certain health conditions, pH can drop to dangerously lower levels resulting in cell damage or death. In other conditions, pH can increase to lethal levels. The catalytic activity of enzymes is particularly sensitive to variations in pH. 45  Some clinical disorders in the acid-base balance Respiratory acidosis: Caused by alveolar hypoventilation and accumulation of CO2 in body, occurs when respiration depth or rate decreases as seen in airway obstruction, neuromuscular disorders, diseases of the central nervous system or in chronic obstructive lung diseases like emphysema. Respiratory alkalosis: Arises from decreased alveolar PCO2 and is most commonly found as the result of hyperventilation due to anxiety. Also can be caused by salicylate poisoning, fever, artificial ventilation, and high altitude (since there is a decrease in total atmospheric pressure and therefore alveolar PCO2 ) 46  Metabolic acidosis: Caused by disorders in which excess lactic acid, acetoacetic acid or β- hydroxybutyric acid are produced, or ingestion of salicylates, ethylene glycol, or methyl alcohol all of which produce strong organic acids. Metabolic alkalosis: alkaline materials can not be synthesized from neutral starting materials, so metabolic alkalosis may be caused by intake of excess alkali (sodium bicarbonate) or abnormal loss of acid (prolonged vomiting). 47  Summary Water is a polar molecule with unique properties. Most metabolic machinery operates in an aqueous environment. Numerous weak, non-covalent interactions decisively influence the folding of macromolecules such as proteins and nucleic acids. 48  Summary cont. A mixture of a weak acid (or base) and its salt resists changes in pH caused by the addition of H+ or OH-. The mixture thus functions as a buffer. In cells and tissues, phosphate and bicarbonate buffer systems maintain intracellular and extracellular fluids at their optimum (physiological) pH, which is usually close to pH 7. Enzymes generally work optimally at this pH. 49