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Chapter 3: Outline-1

Chapter 3: Outline-1. Molecular Nature of Water Noncovalent Bonding Ionic interactions Hydrogen Bonds van der Waals Forces Thermal Properties of Water Solvent Properties of Water Hydrophilic, hydrophobic, and amphipathic molecules Osmotic pressure. Chapter 3: Outline-2.

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Chapter 3: Outline-1

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  1. Chapter 3: Outline-1 • Molecular Nature of Water • Noncovalent Bonding • Ionic interactions Hydrogen Bonds • van der Waals Forces • Thermal Properties of Water • Solvent Properties of Water • Hydrophilic, hydrophobic, and amphipathic molecules • Osmotic pressure

  2. Chapter 3: Outline-2 • Ionization of Water • Acids, bases, and pH • Buffers • Physiological buffers

  3. Water • Solvent for all chemical reactions. • Transports chemicals from place to place. • Helps to maintain constant body temperature. • Part of digestive fluids. • Dissolves excretion products.

  4. 3.1 Molecular Structure of Water • The oxygen in water is sp3 hybridized. Hydrogens are bonded to two of the orbitals. Consequently the water molecule is bent. The H-O-H angle is 104.5o.

  5. Water • Water is a polar molecule. • A polar molecule is one in which one end is partially positive and the other partially negative. • This polarity results from unequal sharing of electrons in the bonds and the specific geometry of the molecule.

  6. Water d- d+ d+ • Water molecule with bond ( ) and net • ( ) dipoles.

  7. Water • Water has an abnormally high boiling point due to intermolecular hydrogen bonding. H bonding is a weak attraction between an electronegative atom in one molecule and an H (on an O or N) in another.

  8. 3.2 Noncovalent Bonding • Ionic interactions • Hydrogen bonding • Van der Waals forces • Dipole-dipole • Dipole-induced dipole • Induced dipole-induced dipole

  9. Typical “Bond” Strengths

  10. Ionic Interactions Salt bridge • Ionic interactions occur between charged atoms or groups. • In proteins, side chains sometimes form ionic salt bridges, particularly in the absence of water which normally hydrates ions.

  11. Hydrogen Bonding • Water molecules hydrogen bond with one another. Four hydrogen bonding attractions are possible per molecule: • two through the • hydrogens and two • through the nonbonding electron pairs.

  12. Van der Waals Attractions • a. Dipole-dipole • b. Dipole-induced dipole • c. Induced dipole-induced dipole

  13. Hydrophobic interactions • Nonpolar molecules tend to coalesce into droplets in water. The repulsions between the water molecules and the nonpolar molecules cause this phenomenon. • The water molecules form a “cage” around the small hydrophobic droplets.

  14. 3.3 Thermal Properties • Hydrogen bonding keeps water in the liquid phase between 0 oC and 100 oC. • Liquid water has a high: • Heat of vaporization-energy to vaporize one mole of liquid at 1 atm • Heat capacity-energy to change the temperataure by 1 oC • Water plays an important role in thermal regulation in living organisms.

  15. 3.4 Water-solvent Properties • Water dissolves chemicals that have an affinity for it, ie. hydrophilic (water loving) materials. • -many ionic compounds • -polar organic compounds • These compounds are soluble in water due to three kinds of noncovalent interactions: • ion-dipole 2. dipole-dipole • 3. hydrogen bonding

  16. Ion-dipole Interactions • Ions are hydrated by water molecules. The water molecules orient so the opposite charge end points to the ion to partially neutralize charge. The shell of water molecules is a solvation sphere.

  17. Dipole-dipole Interactions Dipole-dipole interactions • The polar water molecule interacts with an O or N or an H on an O or N on an organic molecule.

  18. Hydrogen Bonding • A hydrogen attached to an O or N becomes very polarized and highly partial plus. This partial positive charge interacts with the nonbonding electrons on another O or N giving rise to the very powerful hydrogen bond. hydrogen bond shown in yellow

  19. Nonpolar Molecules • Nonpolar molecules have no polar bonds or the bond dipoles cancel due to molecular geometry. • These molecules do not form good attractions with the water molecule. They are insoluble and are said to be hydrophobic (water hating). • eg.: CH3CH2CH2CH2CH2CH3, hexane

  20. Nonpolar Molecules-2 • Water forms hydrogen-bonded cagelike structures around hydrophobic molecules, forcing them out of solution.

  21. Amphipathic Molecules • Amphipathic molecules contain both polar and nonpolar groups. • Ionized fatty acids are amphipathic. The carboxylate group is water soluble and the long carbon chain is not. • Amphipathic molecules tend to form micelles, colloidal aggregates with the charged “head” facing outward to the water and the nonpolar “tail” part inside.

  22. A Micelle

  23. Osmotic Pressure-2 • Osmosis is a spontaneous process in which solvent molecules pass through a semipermeable membrane from a solution of lower solute concentration to a solution of higher solute concentration. • Osmotic pressure is the pressure required to stop osmosis.

  24. Osmotic Pressure-3 • Osmotic pressure (p) is measured in an osmometer.

  25. Osmotic Pressure-4 p = iMRT i = van’t Hoff factor (% as ions) M = molarity (mol/L for dilute solns) R = 0.082 L atm/ mol K T = Kelvin temperature 1 M NaCl is 90% ionized and 10% ion pairs. i = 0.9 + 0.9 + 0.1 = 1.9 Osmolarity (osm/L) = iM

  26. Osmotic Pressure-5 • Because cells have a higher ion concentration than the surrounding fluids, they tend to pick up water through the semipermeable cell membrane. • The cell is said to be hypertonic relative to the surrounding fluid and will burst (hemolyze) if osomotic control is not effected.

  27. Osmotic Pressure-6 • Cells placed in a hypotonic solution will lose water and shrink (crenate). • If cells are placed in an isotonic solution (conc same on both sides of membrane) there is no net passage of water.

  28. 3.5 Ionization of Water • Water dissociates. (self-ionizes) • H2O + H2O = H3O+ + OH- Kw = Ka [H2O]2 = [H3O+ ][OH-]

  29. Water Ionization-2 • The conditions for the water dissociation equilibrium must hold under all situations at 25o. • Kw= [H3O+][OH-]=1 x 10-14 • In neutral water, • [H3O+ ] = [OH-] = 1 x 10-7 M

  30. Water: A/B Properties • When external acids or bases are added to water, the ion product ([H3O+ ][OH-] ) must equal Kw. • The effect of added acids or bases is best understood using the Lowry-Bronsted theory of acids and bases.

  31. Water: A/B Properties-2 • Lowry-Bronsted • acid = proton donor • HA + H2O = H3O+ + A- • A B CA CB • C: conjugate (product) A/B

  32. Water: A/B Properties-3 • Lowry-Bronsted • base = proton acceptor • RNH2 + H2O = OH- + RNH3+ • B A CB CA

  33. Measuring Acidity • Added acids increase the concentration of hydronium ion and bases the concentration of hydroxide ion. • In acid solutions [H3O+] > 1 x 10-7 M • [OH-] < 1 x 10-7 M • In basic solutions [OH-] > 1 x 10-7 M • [H3O+] < 1 x 10-7 M • pH scale measures acidity without using exponential numbers.

  34. pH Scale • Define: pH = -log(10)[H3O+] • 0---------------7---------------14 • acidic basic • [H3O+]=1 x 10-7 M, pH = ? • 7.0

  35. pH Scale-2 • [H3O+]=1 x 10-5 M, pH = ? • 5 (acidic) • [H3O+]=1 x 10-10 M, pH = ? • 10 (basic) • What if preexponential number is not 1?

  36. pH Scale-3 • [H3O+]=2.6 x 10-5 M, pH = ? • 4.59 (acidic) • [H3O+]=6.3 x 10-9 M, pH = ? • 8.20 (basic) • [H3O+]=7.8 x 10-3 M, pH = ? • 2.11 (acidic)

  37. pH Scale-4 • pH to [H3O+]? • inverse log of negative pH • orange juice, pH 3.5. [H3O+]=? • [H3O+] = 3.2 x 10-4 M • urine, pH 6.2. [H3O+]=? • [H3O+] = 6.3 x 10-7 M

  38. Strength of Acids • Strength of an acid is measured by the percent which reacts with water to form hydronium ions. • Strong acids (and bases) ionize close to 100%. • eg. HCl, HBr, HNO3, H2SO4

  39. Strength of Acids-2 • Weak acids (or bases) ionize typically in the 1-5% range . • eg. CH3COCOOH, pyruvic acid • CH3CHOHCOOH, lactic acid • CH3COOH, acetic acid

  40. Strength of Acids-3 • Strength of an acid is also measured by its Ka or pKa values. • HA + H2O = H3O+ + A- Larger Ka and smaller pKa values indicate stronger acids.

  41. Strength of Acids-4 • Ka pKa • CH3COCOOH 3.2x10-3 2.5 • CH3CHOHCOOH 1.4x10-4 3.9 • CH3COOH 1.8x10-5 4.8 • Larger Ka and smaller pKa values indicate stronger acids.

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