Water, p h, and ionic equilibria

Biochemistry 2/e - Garrett & Grisham
CHAPTER 2
Water, pH, and Ionic Equilibria
to accompany
Biochemistry, 2/e
Copyright © 1999 by Harcourt Brace & Company
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 Water
Issues: 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 Equilibria
The 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:
[0.1 ]
¯¯¯¯¯¯¯¯¯¯
[0.9]
• pH = 4.76 + (-0.95)
• pH = 3.81

Acetic Acid
Another 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:
[0.5 ]
¯¯¯¯¯¯¯¯¯¯
[0.5]
• pH = 4.76 + 0
• pH = 4.76 = pKa

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:
[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

Water, p h, and ionic equilibria

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