Chapter 8. Organic acids and Bases: acidity
constants and partitioning behavior
Many organic environmental chemicals are involved
in proton transfer reactions that result in charged
species
Is the partitioning behavior of charged species
different from neutral species??
A proton transfer only occurs if an acid (HA) called
here a proton donor in the classical sense
(Bronstead acid, 1924) reacts with a base (proton
acceptor)
HA
A- + H+
H++ B
HA++ B
BH+
BH+ + A-
A- is called the conjugate base of HA and BH+ is
the conjugate acid of B
The reactions is usually very fast and reversible and
hence can be treated as an equilibrium process.
1
For organic acids (HA)
H3O+ + A-
HA + H2O
For the solvent water
H+ + H2O
H3O+
the equilibrium constant K is
K= [H3O+ ]/{[H+] [H2O]}
Since water is the reference state
K is set to one and GO = O
Because
-RT ln K = GO
So
HA + H2O
H3O+ + A-
becomes
HA
H+ + A -
2
The equilibrium for
HA
H+ + A- is
'
 H+
[H+ ]  'A- [ A - ]
Kia 
'
 HA
[HA]
'
 H+
[H+ ]  'A- [ A - ]
ln Kia  ln
'
 HA
[HA]
= -rGo/RT
The prime on the “’s” indicates infinite dilution
Kia is commonly referred to as the acidity constant
Typically when analytical measurements are made
the hydrogen activity is measured
pH = -log {’H+ [H+]} while [HA] and [A-], are usually
in concentrations are in molar units,
this gives a mixed acidity constant K*ia that is
usually defined for a given aqueous medium
(ie. 0.05 - 0.01 M salt solution)
3
since
'
 H+
[H+ ]  'A- [ A - ]
Kia 
'
 HA
[HA]
'
+

[
H
]
[
A
]
H+
*
K ia 
; K ia
[HA]
'
 HA
 'A-
 K*ia
for low concentrations of [HA] and [A] in an
aqueous media K*ia = Kia
in treating low concentrations the convention is
pX = -log [X]
pH = -log {’H+ [H+]}
'
+

[
H
]
[
A
]
H+
*
log K ia  log
[HA]
log {[A-]/[HA]} = log Kia - log {’H+ [H+]}
4
log {[A-]/[HA]} = pH – pK*ia
(Henderson-Hasselbach
equation)
at what point does pH = pKia?
when [A-] = [HA]
Take the case where the pKia of an organic acid is
low , say the pKia = - 0.3
log K*  log
'
 H+
[H+ ] [ A - ]
ia
[HA]
pKia = -log [Kia]
If pKa is negative, log Ka is positive and this makes
H+ high in the definition of Ka
Examples of strong acids are trichloroacetic acid
and 2,4,6 trinitrophenol.
At ambient pH values (4-10) low pKa compounds will
be present in natural waters in their ionized
(dissociated form) as their anions.
Conversely, weaker acids are ones with relatively
high pKa values, and very weak acids (pKa = 9 - 12)
will exist in natural waters in their undissociated
form.
5
Many important organic acids have pKia(s) between
4 and 10;
Given this treatment it is useful to describe the
amount present in the acid (unionized or
dissociated) form and this will be denoted by a.
`
[HA]
a
- 
[HA] + [A ]
1
[A - ]
1+
[HA]
log {[A-]/[HA]} = pH - pKa
(Henderson-Hasselbach
equation)
a
1
pH-pK
ia
1+10
for example, 2-nitrophenol has a pKia of 7.17 at a pH
of 7. What is the fraction in the undissociated acid
form?
a 
1
1+10
7.0 -7.17
 0.597
6
To Review, what we have said is that from
log {[A-]/[HA]} = pH - pKa
and pH and pKa of an acid we can estimate
the [A-]/[HA] ratio, and the fraction in the acid form,
from a (Table 8.1 page 180, old book; see p 250 new book)
7
Example: Calculate the fraction of
pentachlorophenol (PCP) in the neutral form in a rain
drop at pH 4 and in lake water at pH 8.
From Table 8.1 (new and old book), PCP has a pKa
value of 4.75
a
1
pH-pKa
1+10
rain= 1/(1+10-0.75) = 0.849
lake= 1/(1+103.25) = 0.00056
in the rain drop, only 15% of the PCP is present as
the phenolate anion, whereas in the lake water,
virtually all of the PCP is ionized.
8
Organic Bases
By analogy with acids, the basic content (basicity) of
an organic base in water can be described as:
OH- + BH+ (BH+ is the conjugate acid of B)
B + H2O
Kib 
'
+
 'OH- [OH- ]  BH
[
BH
]
+
 B [B]
The reaction of a neutral base [B ] with water forms
the cation, BH+
Instead of working with a Kib dissociation
equilibrium constant, by convention, the acidity
constant is defined in terms of the conjugate acid of
B, or BH+.
BH+
Kia 
H+ + B
 H'+ [H+ ]  B' [B]
 BH [BH+ ]
9
If we multiply Kia and Kib
Kib 
'
+
 'OH- [OH- ]  BH
[
BH
]
+
 B [B]
(BH+ is the conjugate acid of B)
Kia 
 H'+ [H+ ]  B' [B]
 BH [BH+ ]
Kia Kib = Kw = ’H+ [H+] x ’OH- [OH-] = 1.01 x10-14
So for a given compound in water
pKia = pKw - pKib
Low pKia means a strong acid and the pKia of its
conjugate base is high, which means the conjugate
base it not in its ionized from or is weak.
The stronger the base (low pKib means the negative log of a
high Kb, which means more ionized base) the weaker its
conjugated acid (high pKa of the conjugate acid)
10
Hence a neutral base like methylamine with a conjugate
acid pKia of 10.66 will thus have a pKib of:
CH3–NH2 + H20 CH3-NH3++ OHpKa = pKw – pKb;
10.66 = 14 - pKb
and will exist in natural waters in its dissociated (ionized)
or cation state
11
We will be interested in compounds which have pKia
values in the 3-11 range
The above structures suggest compounds with phenolic,
carboxylic, amino , nitrogen in the rings groups, and aromatic
thiol groups should be considered.
For example depending on the structure pKa can differ by an
order of magnitude. So understanding how the structure
influences pKia is important.
12
It is also possible to have diportic acidic compounds and
di-basic compounds.
From measurements it is possible to see how different
species of a parent molecule vary with pH, depending on
their pKa
13
Speciation in Natural Waters
Given the pKia of a compound, to what extent does it
distribute in natural waters.
The pH in natural water is determined largely by
inorganic acids and bases (H2CO3  HCO3-  CO3-2).
These species act as H+ buffers; ie. a small amt. of
acid or base will not effect the overall distribution.
To illustrate, let’s assume that for a given acid-base
pair with a pKia of 7, is at equal concentrations in
water at 10-3moles/L
log {[A-]/[HA]} = pH – pKia (Henderson-Hasselbach
equation)
so,
pH = pKa + log {10-3/10-3} = 7
If we add 10-5 moles of additional strong acid the
total acid will be: HA + 10-5 new acid = 1.01x10-3;
A-, however reacts with the additional acid, to give
HA and its concentration goes from
1x10-3 to 0.99x10-3
hence: pH = pKia + log {0.99x10-3/1.01x10-3} = 6.991;
so, the addition of small amts. of a strong acid will
not overly affect the pH of the natural water.
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Chemical Structure and Acidity Constants
Since Kia is a partitioning constant we can directly
write:
ln Kia = -rGo/RT
the question is what is rGo ???
Is it not the chemical potential between the products
and initial reactants ??
for
HA


H+ + A-
-rGo = oHA-oA-BpKa = (oHA-oA-B-)/{RT 2.303}
so we must ask the question, what will change
the standard state chemical potential, so that it
alters its tendency to react in the case of the acid
HA to is conjugate base A-.
The engineers… would say its chemical
structure???
15
If there are substituents that makes it easier for
HA to loose a H+, this would influence oHA
Look at the following structures and their pKia
values,
where HA  A- + H+
CH3CH2 CH2-COOH
pKia = 4.81
CH3 CH CH2-COOH
Cl
PKia = 4.05
CH2CH2 CH2-COOH
Cl
pKia = 4.52
CH3CH2 CH-COOH
Cl
pKia = 2.86
chlorine is very electronegative and has a high
electron with drawing potential. The effect is called
a negative inductive effect. The resulting A- anion
from the above acids are stabilized by Cl
withdrawing the negative charge into the structure.
Notice as the chlorine gets closer to the carboxylic
group, the effect increases.
16
Delocalization electron effects: Compounds that
have alternate double bonds ( electrons) can have
inductive effects transmitted over longer distances
than just single bonded carbon molecules
Conjugated structure
Resonance structure
H3C- C=C-C=CH2
H H H
H3C-C-C-C-CH2
H H H
We would also expect resonance to be important
with aromatic ring structures (Figure 8.3 p 168, new book,
p 258)
17
Effects of position of a nitro substituent (Figure 8.4 p 169,
p 259 new book)
Proximity effects; (top) hydrogen bonding and
(bottom) steric interactions(Figure 8.5 p 170)
pKia= 7.28
18
19
The Hammett Correlation
In 1940 Hammett recognized for substituted benzoic
acids the effects of substituent groups on the
dissociation of the acid group
COOH
COO-
+H+
R
R
rGo= rGoH +  rGoi
since rGo = -RTln Ka
-RTln Ka = -RTln KaH - RTln Ka,i
Effect on the free energy change from dissociation
could be represented as the sum of the free energy
change by the unsubstituted benzoic acid and the
contributions from the various R groups.
Hammett then says that the effects of the
substituent group i, on the free energy can be
represented as:Goi = -2.303 RTi; where i =ln Ka,i
20
Goi = -2.303 RTi; where i = ln Ka,i
going back to:
-RTln Ka = -RTln Ka,H - RTln Ka,i
log Ka= log KaH +i
so,
log (Ka / KaH )= I
and
pKa = pKaH - I
21
Table 8.4 Hammet Constants
Phenols and
anilines
22
Hammet also observes that when considering other
compound classes, like phenyl acetic acids the 
values developed for benzoic acid can be used
log (Ka / KaH) = i
Figure 8.7, p 174
23
Exceptions: As a function of direct resonance, para
phenol substituents decrease the pKa more than
would be predicted by just the para  values.
Apparent values for phenols have been
determined as per Table 8.4
24
25
26
The air-water equilibrium
The distribution coefficient between the air and
water of an organic acid can be described as:
Daw (HA,A-) = [HA]air/{[HA]water+ [A-]water}
Multiplying by [HA]water/[HA]water
[HA]air
[HA] water
D aw 
x
[HA] water  [A] water [HA] water
[HA] water
[HA]air
D aw 
x
[HA] water  [A] water [HA] water
Daw (HA,A-) = a Kiaw
a
1
pH-pKa
1+10
Go through example on page 269 of new book
27
Conversion of Nicotine in tobacco smoke to its
volatile and Available Free-Base Form through the
Action of Gaseous Ammonia (Jim Pankow, et al, ES&T, 32,
2428-2433, 1997)
N
pka1 =3.12
pKa2 = 8.02
CH3
N
For pKa1
Nicotine-H+  nicotine + H+
At a pH of 5.5 for normal tobacco smoke
Pankow shows that the fraction fb in the free base
form =
 fb 
1
-pH+pK
1+10
-2pH+pK
a2 + 10
+ pK
a1
a2

at a pH of 5.5, fb= 0.002
28
at a pH of 8, fb= 0.5
So increasing the pH increases the amount of free
base nicotine
The amount of nicotine that is partitioning if it were
all in the free base form would be:
Kp = nicotinepart/(nicotinegasxTSP)
To account for acid-base equilibrium Kp is multiplied
by fb
Pankow writes: Nicotine in the particle phase that is
not in the free base form is not available for gasparticle partitioning. By raising the pH with
ammonium diphosphate as a “flavor enhancer”
more nicotine can come out of the particle phase
and be available for gas phase transfer to the lungs.
He also suggests that nicotine on high pH particles
deposited in the lungs will also diffuse off the particles
directly to the alveolar lung lining.
He then goes on to say: “the role of ammonia in tobacco
smoke is analogous to what happens when alkaloid
cocaine is “free based” or used in the “crack” form. Freebased cocaine and crack are typically prepared by
reacting “street” cocaine (usually the HCL salt of cocaine)
with aqueous ammonia and ammonium bicarbonate” the
29
basified cocaine is more lipid soluble and gives a “high”
when smoked that is comparable to intravenous injection.
30
Influence of organic acid-base equilibrium on
partitioning.
The water solubility of an ionized organic acid or
salt, is generally several orders of magnitude higher
then the neutral species, HA, (Csatw,HA).
The sum of the ionized and HA species is Ctot, and it
has a saturated concentration of (Csatw,tot).
Csatw,totx ia = Csatw,HA; so Csatw,tot = Csatw,HAia
|
31
32
For octanol water partitioning things are not as
straight forward because the organic and the ionized
organic can partition into both the water and the
octanol phase.
[HA] octanol,tot
D ow(HA,A-) 
[HA] water  [A] water
33