Lecture 2 Water: The Medium of Life

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