Water, p h, and ionic equilibria

Biochemistry 2/e - Garrett & Grisham

Copyright © 1999 by Harcourt Brace & Company

CHAPTER 2Water, pH, and Ionic Equilibria

to accompany

Biochemistry, 2/e

by

Reginald Garrett and Charles Grisham

All rights reserved. Requests for permission to make copies of any part of the work should be mailed to: Permissions Department, Harcourt Brace & Company, 6277 Sea Harbor Drive, Orlando, Florida 32887-6777



Outline

• 2.1 Properties of Water

• 2.2 pH

• 2.3 Buffers

• 2.4 Water's Unique Role in the Fitness of the Environment



Properties of Water

• High b.p., m.p., heat of vaporization, surface tension

• Bent structure makes it polar

• Non-tetrahedral bond angles

• H-bond donor and acceptor

• Potential to form four H-bonds per water





Comparison of Ice and WaterIssues: H-bonds and Motion

• Ice: 4 H-bonds per water molecule

• Water: 2.3 H-bonds per water molecule

• Ice: H-bond lifetime - about 10 microsec

• Water: H-bond lifetime - about 10 psec

• (10 psec = 0.00000000001 sec)

• Thats "one times ten to the minus eleven second"!







Solvent Properties of Water

• Ions are always hydrated in water and carry around a "hydration shell"

• Water forms H-bonds with polar solutes

• Hydrophobic interactions - a "secret of life"





Hydrophobic Interactions

• A nonpolar solute "organizes" water

• The H-bond network of water reorganizes to accommodate the nonpolar solute

• This is an increase in "order" of water

• This is a decrease in ENTROPY





Amphiphilic Molecules

Also called "amphipathic"

• Refers to molecules that contain both polar and nonpolar groups

• Equivalently - to molecules that are attracted to both polar and nonpolar environments

• Good examples - fatty acids









Acid-base EquilibriaThe pH Scale

• A convenient means of writing small concentrations:

• pH = -log10 [H+]

• Sørensen (Denmark)

• If [H+] = 1 x 10 -7 M

• Then pH = 7









Dissociation of Weak Electrolytes

Consider a weak acid, HA• The acid dissociation constant is given

by:• HA H+ + A-

• Ka = [ H + ] [ A - ] ____________________

[HA]



The Henderson-Hasselbalch Equation

Know this! You'll use it constantly.

• For any acid HA, the relationship between the pKa, the concentrations existing at equilibrium and the solution pH is given by:

• pH = pKa + log10 [A¯ ]

¯¯¯¯¯¯¯¯¯¯

[HA]





Consider the Dissociation of Acetic Acid

Assume 0.1 eq base has been added to a fully protonated solution of acetic acid

• The Henderson-Hasselbalch equation can be used to calculate the pH of the solution:With 0.1 eq OH¯ added:

• pH = pKa + log10 [0.1 ]

¯¯¯¯¯¯¯¯¯¯

[0.9]

• pH = 4.76 + (-0.95)

• pH = 3.81



Consider the Dissociation of Acetic AcidAnother case....

• What happens if exactly 0.5 eq of base is added to a solution of the fully protonated acetic acid?

• With 0.5 eq OH¯ added:• pH = pKa + log10

[0.5 ] ¯¯¯¯¯¯¯¯¯¯

[0.5]

• pH = 4.76 + 0

• pH = 4.76 = pKa



Consider the Dissociation of Acetic Acid

A final case to consider....• What is the pH if 0.9 eq of base is added

to a solution of the fully protonated acid?• With 0.9 eq OH¯ added:

• pH = pKa + log10 [0.9 ]

¯¯¯¯¯¯¯¯¯¯

[0.1]

• pH = 4.76 + 0.95

• pH = 5.71







Buffers

• Buffers are solutions that resist changes in pH as acid and base are added

• Most buffers consist of a weak acid and its conjugate base

• Note in Figure 2.15 how the plot of pH versus base added is flat near the pKa

• Buffers can only be used reliably within a pH unit of their pKa









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Water, p h, and ionic equilibria

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