Lecture 2: Bonding, structure of water and acids/bases

1. Lecture 2: Bonding, structure of water and acids/bases Dr. NadeemAsad BIOL 313

2. Properties of water • WATER is the solvent of choice for biological systems • Constitutes 70-85% of cell weight, typically • Important as a solvent and a reactant in biochemical reactions • Helps regulate temperature since it is able to absorb large amounts of heat • Helps regulate intracellular pH • Used for transport – delivers nutrients and removes waste from cells • Water is a unique solvent whose properties areextremely important to biochemistry

4. Water and it’s importance

5. Water and it’s importance Water is a unique and universal solvent whose properties are extremely important to biochemistry

6. Nature of bonding in water

7. Angular geometry of water imp for living organisms • Water is polar- O2 with its unshared electrons carries a partial -ve charge e -0.66 and the H atoms a + ve charge e + 0.33 • Electrostatic interactions b/w dipoles of H2O imp

8. Nature of bonding in water • Water is a dipole due to it’s geometry and the difference in electronegativity between hydrogen and oxygen. Oxygen is more electronegative than hydrogen • Oxygen is sp3 hybridised, tetrahedral electron geometry, bent geometry

10. Polar nature of water • The polar nature and geometry of the water molecule allows water molecules to form hydrogen bonds with each other and with dissolved hydrophilic substances. • Hydrogen bonds between water molecules= electrostatic attraction between the oxygen atom of one water and the hydrogen of another • Water can also form hydrogen bonds with functional groups of hydrophilic (polar or ionic) biomolecules and organic compounds. •  hydrogen bond donors •  hydrogen bond acceptors

11. Polar nature of water

12. Hydrogen bonding in water

16. Non-Covalent Interactions • Relatively weak and reversible • 1-Hydrogen Bonds: Special dipole-dipole interaction (electronegative atom (e.g. O or N) interacts with H atom that is partially positive (i.e. attached to N, O, F) • Very important for protein and DNA structure.

17. Non-Covalent Interactions • 2. Van Der Waals interactionsa. Dispersion Forces (London Forces) (induced dipoles in non-polar molecules) b. Dipole-dipole forces

18. Non-Covalent Interactions • 3. Ionic bonding-Electrostatic interactions between two oppositely charged ions

19. Non-Covalent Interactions • 4. Hydrophobic Interactions/Hydrophobic Effect: Relations between water and hydrophobic molecules (low water-soluble molecules) • Nonpolar substances tend to aggregate in aqueous solution and exclude water molecules. Thermodynamically unfavorable to dissolve hydrophobic substance in water • Entropy-driven process. ΔG = ΔH – TΔS

20. Bond strength

23. Hydrophobic effect • The hydrophobic effect is the observed tendency of nonpolar substances to aggregate in aqueous solution and exclude water molecules. • Mostly entropy driven: Water more ordered at interface.Entropy of water molecules increases upon exclusion of non-polar species

24. Hydrophobic effect • Water molecules align themselves around non-polar molecule and lose freedom to form hydrogen bonds • Entropy lost in system results in thermodynamic barrier • Multiple molecules aggregate – increase the entropy of the system because fewer water molecules needed to surround the aggregate than to hydrate each dispersed molecule

25. Hydrophobic effect • Exclusion of nonpolar substances from aqueous solution = HYDROPHOBIC EFFECT • Important concept in biochemistry – governs protein folding (hydrophobic amino acids are on the interior of the protein), formation of membranes (hydrophobic lipid tails sequestered in bilayer) • Water soluble compounds are those in which the interactions between the solute and water are greater than those between solute molecules. i.e. salts, biological molecules that have polar or ionic groups (e.g. glucose, ethanol)

30. H bonding in drug design • 25% of drugs contain Fluorine! Why?Takes the place of H in a chemical structure. Very electronegative – thus electron withdrawing. Gives different properties to the group. • CF3 – Electron withdrawing, can decrease basicity of nearby amino groups – fewer positive charges – can penetrate cells better. • C – F bond can form hydrogen bonds and other dipole-dipole interactions, potentially increasing binding to target molecule in the body. May be more effective at lower concentrations.

31. Non-Covalent Interactions in proteins • Non-covalent interactions also form between two biomolecules (e.g. proteins & DNA) • The 3-dimensional structure of many biological molecules (eg. proteins) and macromolecular structures (eg. membranes, DNA) is determined by hydrogen bonding, hydrophobic interactions, ionic interactions and van der Waals interactions. • Hydrogen bonds are weak but their high abundance makes them important!

33. Key Concepts

35. The ionization of water = equilibrium expression in which the concentration of the parent substance = denominator • And dissociated products =numerator • K is the dissociation constant; • Because the concentration of the undissociated H2O ([H2O]) • is so much larger than the concentrations of its component ions, it can be • considered constant and incorporated into K to yield an expression for the ionization of water, Quantities in square brackets symbolize the molar concentrations of the indicated substances, which in many cases are only negligibly different from their activities

38. What is pH? • pH is a measure of Hydrogen ion concentration (acidity or alkalinity of a solution) • The first important concept is that the acidity of a solution reflects only the free hydrogen ions, not those still bound to anions

39. Blood pH • Blood is in contact with nearly every blood cell • Regulation of it’s pH is particularly critical • Normally, blood pH varies within a very narrow range (7.35-7.45 • If blood pH varies from these limits it maybe fatal . Living range: pH 7.0-7.6 • Acidosis-normal range-alkalosis

40. pH scale

41. Where do acids in the body come from? • Some acids can enter the body in food • Most are generated by: • Breakdown of proteins • Incomplete oxidation of fats or glucose • The loading and transport of carbon dioxide in the blood

43. Acids can be classified according to their relative strengths, that is, their abilities to transfer a proton to water. Table 2-4 are known as weak acids because they are only partially ionized in aqueous solution (K 1). Many of the so-called mineral acids, such as HClO4, HNO3, and HCl, are strong acids (K 1). Since strong acids rapidly transfer all their protons to H2O, the strongest acid that can stably exist in aqueous solutions is H3O. Likewise, there can be no stronger base in aqueous solutions than OH.

48. Acid-base balance • It is regulated in the body by: • The lungs • The kidneys • Systems in the blood known as chemical buffers

49. Buffering • Buffers resist abrupt and large swings in the pH of body fluids by: • releasing H+ (acting as acids) when the pH begins to rise • binding H+ (acting as bases) when the pH drops

50. Dissociation of acids • Acids are proton donors • Bases are proton acceptors • Acids that dissociate completely in solution are strong acids • Acids that dissociate incompletely are weak acids • Strong bases are more effective proton acceptors than are weak bases