Lecture 1: Introduction and review

1
Lecture 1 Introduction and review

• Quiz 1
• Website http//www.esf.edu/chemistry/nomura/fch53
  0/
• Review of acid/base chemistry
• Universal features of cells on Earth
• Cell types Prokaryotes and Eukaryotes

2
Review of pH, acids, and bases

• pH is generally defined as the negative logarithm
  of the hydrogen ion activities (concentration)
  expressed over 14 orders of magnitude
• pH -log10 H
• The pH scale is a reciprocal relationship between
  H and OH-
• Because the pH scale is based on negative
  logarithms, low pH values represent the highest
  H and thus the lowest OH-
• At neutrality, pH 7, H OH-

3
Review of pH, acids, and bases

• Strong electrolytes dissociate completely in
  water
• Electrolytes are substances capable of generating
  ions in solution
• Increase the electrical conductivity of the
  solution
• The dissociation of a strong acid in water
• The equilibrium constant is
• H2O is constant in dilute aq. Solutions and is
  incorporated into the equilibrium constant.
  giving rise to a new term Ka-the acid
  dissociation constant KH2O, H3O is
  expressed as H

HCl H2O H3O Cl-
H3OCl-
K
H2OHCl
Because Ka is large for HCl, H in solution
HCl added to solution. Thus, a 1M HCl solution
has a pH of 0, a 1 mM HCl solution has a pH of 3,
and so on. Conversely, 0.1 M NaOH solution has a
pH of 13.
4
For a strong acid Ka will approach be large
because the nearly all of the protons will be
dissociated. The H at equilibrium is equal to
the initial concentration of the acid.
Calculate the pH of a 1M HCl solution
HCl H2O H3O Cl- 0.0004
at equilibrium 99.996 at equilibrium
Since we are at equilibrium, H3O is equal to the
initial concentration of acid.
H H3O HCl 1M
We know that pH is the -log of H, therefore
for 1M HCl at equilibrium
pH -log10 H
pH -log10 (1)
pH 0
5
pH
0 (100) 1.0 0.00000000000001 (10-14)
1 (10-1) 0.1 0.0000000000001 (10-13)
2 (10-2) 0.01 0.000000000001 (10-12)
3 (10-3) 0.001 0.00000000001 (10-11)
4 (10-4) 0.0001 0.0000000001 (10-10)
5 (10-5) 0.00001 0.000000001 (10-9)
6 (10-6) 0.000001 0.00000001 (10-8)
7 (10-7) 0.0000001 0.0000001 (10-7)
8 (10-8) 0.00000001 0.000001 (10-6)
9 (10-9) 0.000000001 0.00001 (10-5)
10 (10-10) 0.0000000001 0.0001 (10-4)
11 (10-11) 0.00000000001 0.001 (10-3)
12 (10-12) 0.000000000001 0.01 (10-2)
13 (10-13) 0.0000000000001 0.1 (10-1)
14 (10-14) 0.00000000000001 1.0 (100 )
6
Review of pH, acids, and bases

• Weak electrolytes only slightly dissociate in
  water
• Acetic acid, CH3COOH
• The dissociation of a weak acid in water
• The acid dissociation constant is
• Ka is also called the ionization constant,
  because Ka is small, most of the acetic acid is
  not ionized.

CH3COOH H2O H3O CH3COO-
HCH3COO-
Ka
1.74 X 10-5 M
CH3COOH
7
Acid dissociation constant

• The general ionization of an acid is as follows
• So the acid dissociation constant is as follows

There are many orders of magnitude spanned by Ka
values, so pKa is used instead
pKa - log10 Ka
The larger the value of the pKa, the smaller the
extent of dissociation. pKa lt2 is a strong acid
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Henderson-Hasselbalch Equation

• Describes the dissociation of a weak acid in the
  presence of its conjugate base
• The general ionization of a weak acid is as
  follows
• So the acid dissociation constant is as follows
• Rearranging this expression in terms of the
  parameter of interest H gives the following

HA H A-
HA-
Ka
HA
Ka
HA
H
A-
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Henderson-Hasselbalch Equation

• Take the log of both sides

log Ka log
logH
Change the signs and define pKa as -log Ka
pKa - log
pH
or
A-
pKa log
pH
HA
10
Titration curves and buffers

• Titration curves can be calculated by the
  Henderson-Hasselbalch equation
• As OH- is added to the reaction, it reacts
  completely with HA to form A-
• x the equivalents of OH- added and V represents
  the volume of the solution. If we let co
  represent HA equivalents initially present, then
• We can reincorporate this into the
  Henderson-Hasselbalch eqn.

A-