lecture 2 water: the medium of life
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Lecture 2 Water: The Medium of Life. Mintel Office Hours Today 1:45 –3:00 Noyes 208. Fred Diehl – Univ. of Virginia. Fred. Physical Properties of Compounds. Explain the difference in these properties. A Water Molecule. Fig. 2-1, p. 29. Structure of Ice – Hydrogen Bonding. - PowerPoint PPT PresentationTRANSCRIPT
Lecture 2Water: The Medium of Life
Mintel Office Hours Today
1:45 –3:00
Noyes 208
Fred Diehl – Univ. of Virginia
Fred
Physical Properties of Compounds
Compound MP (oC) BP (oC) Hvap
(cal/g)Hfus
(cal/g)
H2O 0 100 540 80
H2S -83 60 132 16.7
NH3 -78 -33 327 84
Explain the difference in these properties.
Fig. 2-1, p. 29
A Water Molecule
Structure of Ice – Hydrogen Bonding
• Each molecule of water can be hydrogen-bonded to up to four other water molecules
Fig. 2-2, p. 29
Structure of Ice
Liquid Water
• Lacks the lattice-like structure of ice.
• H-bonds are not colinear with a line joining the centers of the atoms involved.
• Therefore the H-bonds are weaker and water is fluid.
• H-bonds are dynamically formed and broken.
Fig. 2-3a, p. 30
DynamicFormation ofH-bonds inliquid water.
Note the time scale.
Solvent Properties of Water
• Example – Solubility of sodium chloride
• Sodium and chloride ions are hydrated.
• Water molecules are oriented in an opposite direction about sodium and chloride ions, because the interaction is electrostatic.
Fig. 2-4, p. 31
Water’s Dielectric Constant
Solvent D
Water 78.5
Methanol 32.6
F = e1e2/Dr2
whereF is the attractive force between oppositely charged ions
e = charge on an ionr = distance between the ions
D = dielectric constant
Table 2-1, p. 31
Hydrophobic Interactions
• Clathrate (“iceberg”) structure forms surrounding hydrocarbon tails in an aqueous environment, as shown on the next slide.
Fig. 2-5, p. 32
Iceberg structure
Fig. 2-6, p. 32More iceberg cages
Note disruption ofcages when hydrocarbonscome together.
Fig. 2-7a, p. 33
An Amphipathic Molecule
Fig. 2-7b, p. 33
Micelle Formation
Soaps and Detergents
• Grease is dissolved in the hydrocarbon tails of a soap or a detergent.
• Then, when our coated hands are placed into water, micelles form and disperse down the drain with the grease trapped inside.
Fig. 2-8, p. 34
P = iRTm, where i = number of ions, R=gas constant, T=absolute temperature, m = molality(Chemical purists: I know the units don’t work out. See me for an explanation.)
Osmotic Pressure
Fig. 2-9, p. 34
Ionization of Water
p. 34
Formation of Hydronium Ions, H3O+
Fig. 2-10, p. 35
Hydration of aHydronium IonItself
Ion Product of Water• KW = 10 -14 = [H+] [OH-]
• In precisely neutral water, [H+] = [OH-] = 10-7M
Definition of pH
• pH = log [1/H+] = -log [H+]
• A logarithmic scale is more convenient for representing the large range of H+ concentrations encountered in biochemistry, just as the Richter scale is more useful for representing the large range of energy values for earthquakes.
• By extension, pK = log (1/K) - log K
Table 2-2, p. 36
Table 2-3, p. 36
Dissociation of Strong Acids
• Example: HCl
• Completely dissociated in solution
Dissociation of Weak Acids
• Example: Acetic Acid
• Incompletely dissociated in solution
Table 2-4, p. 39
Fig. 2-11, p. 39
TitrationofAceticAcid – AClosedSystem(Matter isnot ex-changedwith theenvironment.
Linear scale
Logarithmicscale
Fig. 2-11a, p. 39
Henderson-Hasselbalch Equation
• pH = pK + log ([A-]/[HA])
• Derivation of the equation
• Describes the shape of a titration curve in the neighborhood of the pK.
• In a molecule with several ionizable groups, there is one H-H equation for each group that titrates.
• The pK of a weak acid is that pH where HA is half-titrated.
Fig. 2-11b, p. 39
Titration Curvefor HAc
Fig. 2-12, p. 40
The Titration ofSome ImportantWeak Acids
Phosphoric Acid E quilibria
(1) H3PO4 = H+ + H2PO4- (pK = 2.15)
(2) H2PO4- = H+ + HPO4
2- (pK = 7.20)
(3) HPO42- = H+ + PO4
3- (pK = 12.4)
Fig. 2-13, p. 41
Buffers
• Definition – A mixture of a weak acid and its conjugate base.
• Function – Maintains cellular pH, and that of bodily fluids like plasma.
• Important intracellular buffers are the phosphate system and the histidine system.
• Buffer capacity is generally best within 1 pH unit of the pK.
Fig. 2-14, p. 41
Bicarbonate Buffer System
• Most important buffer system in blood plasma.• An open system – Exchanges matter with the
environment.
• C02 + H2O = CO2(d)(H2O) = H2CO3
• H2CO3 = H+ + HCO3-
• Henderson-Hasselbalch Equation
• pH = pK + log ([HCO3-]/ [H2CO3]
Bicarbonate Buffer System
• Normal valuespH = 7.4
[HCO3-] = 24 mM
[H2CO3] = 1.2 mM
[HCO3-]/ [H2CO3] = 20/1
pCO2 = 40 mmHg
Properties of Open SystemClosed System; HAc, pK=4.76
Base A- HA pH pH/Base0.5 0.5 4.76
0.05 0.55 0.45 4.84715 0.0871501760.1 0.6 0.4 4.936091 0.088941083
0.15 0.65 0.35 5.028845 0.0927540530.2 0.7 0.3 5.127977 0.099131473
0.25 0.75 0.25 5.237121 0.1091444690.3 0.8 0.2 5.36206 0.124938737
Open System; Bicarbonate, pK=6.1
HCO3- H2CO3 pH1.2 1.2 6.1
1.25 1.2 6.117729 0.0177287671.3 1.2 6.134762 0.017033339
1.35 1.2 6.151153 0.0163904161.4 1.2 6.166947 0.015794267
1.45 1.2 6.182187 0.0152399671.5 1.2 6.19691 0.014723257
24 1.2 7.40103
Moral
• In a closed system a buffer is poorer as one moves away from the pK.
• In an open system a buffer is better as one moves away from the pK.
• This is why the bicarbonate buffer system, with a pK = 6.1, is effective at normal plasma pH = 7.4.
Respiratory Acidosis
• Caused, for example, by breathing in and out of a paper bag.
• The partial pressure of carbon dioxide in the blood increases, and plasma pH accordingly falls.
Respiratory Alkalosis
• Caused, for example, by hyperventilating.
• More carbon dioxide is blown off by the lungs, and the plasma pH accordingly rises.
Problems
• Do Problems 1 and 2 in Problem Set 1, which is posted on the course web site.
Learning Goals
• Know the physical and chemical properties of water important to biological system.
• Understand the dissociation of weak electrolytes, and their importance in maintaining pH in cells and tissues.
• Understand the bicarbonate buffer system and its importance