Buffers
and pH
|
Strong acids | Weak acids
| Application in the cell |
An understanding of what H ion concentration
means and how H ion concentration changes is very important in
biochemical processes and in following clinical progression of
certain disease processes. The concentrations are usually expressed
as one mole/liter (1M or gram molecular weight of a compound in
1000 gm of solvent; in case of water 1000 gm=1000 ml or one liter)
of water); one mmole/liter (1 mM); one mmole/liter
(1mM) and so on. In Biochemistry, concentrations
are always expressed in this way.
Strong
Acids
If
you put 1 mmole of HCl in a liter of water its concentration would
be 1 mM. Since HCl is a strong acid, it is completely dissociated
in water:
HCl -----> H+ + Cl-.
Note that H+
(also called a proton) and Cl- ions (called anions or conjugate base)
are surrounded by water molecules and form hydrogen bonds with
a number of water molecules. Because of this essentially irreversible
dissociation, the concentrations of protons and Cl-
ions would be equal to 1 mM. The hydrogen ion concentration is
therefore, 1 mM or 10-3 M.
This way of representation H+ is
cumbersome and has been changed to a logarithmic scale called
pH:
pH=
- log {H+} where {H+} is represented in
gram equivalent per liter. --- (eq.1)
Since 1 mM
HCl provides 1 mequivalent of H+ and one mequivalent
of Cl ions, the pH of 1mM HCl would be:
pH=
-log {10-3m Equ/lt} = - (-) log {103} =
3 log10 =3.
Thus
pH values can be expressed as whole numbers and fractions. For
"strong' acids like H2SO4 complete dissociation
in water would yield two equivalents of H+ and one
equivalent of SO4 ions. Therefore, 1 mM H2SO4
would provide 2 m equivalents of H+ and one equivalent
of SO4 ions. The pH of 1 mM H2SO4
solution would be:
pH = - log {2x10-3 m equ}
=2.7
For
acids which do not completely dissociate into H ions and anions,
one would have to calculate their H+ concentrations
from their known dissociation constants (Ka).
The section on Ka will be discussed later.
Weak
Acids
Weak
acids will dissociate in solution, but they do so less than 100%.
So water dissociates to H+ and OH- ions
to a very small extent. This can be represented as:
H2O <------------>
H+ + OH-.
This
is represented by a dissociation constant, Ka.
Therefore
Ka
= {H}{OH}/H2O. -------------(eq.
2)
The
concentration of water is very high ( 55.5 M; 1000/18 gm mol weight
of water) compared to H or OH concentrations. Therefore both sides
of equation are multiplied by 55.5 M to give a new expression:
Ka
x 55.5M = Kw = {H}{OH}. The value of Kw = 10 -14
M2 ---------(eq. 3)
Since
in water the concentrations of H and OH ions are equal, one can
write:
{H+}{OH} = {H+}2 = 10-14
M2 or {H+}
= 10--7 M
Therefore pH of water
would be - log {10-7M} = 7.
A
pH of 7 is called neutral
because the concentration of H and OH ions are equal,
so the solution is neither acidic or alkaline.
Equation
3 also indicates that:
pH +
pOH =14; thus if pH=7 then
pOH =7.
When a solution becomes more acidic i.e. if pH is 6, the the
pOH would be 14-6 =8. This means that when {H+} is 10-6M,
the {OH} is 10-8M.
Weak acids like acetic acid,
lactic acid do not completely dissociate into {H+} and conjugate
base. Therefore, the concentration of {H+} would have to be calculated
knowing the total concentrations of acid and conjugate base.
Application
in the Cell
Amino acids, like other organic molecules,
can be characterized in terms of their chemical and physical properties.
One of the most important of these is their ability to undergo
proton dissociation or association. The amino group of an
amino acid is a (moderately) weak base. The carboxyl group
is a weak acid.
The extent to which a carboxyl group
of a specific amino acid undergoes proton dissociation is expressed
by an acid-dissociation constant, Ka, which describes
the following process occurring, from left to right:
AH + H2O
===== H3O+ +
A-
[H3O+] [A-]
Ka = --------------------
[AH] [H2O]
and, pH = pKa
+ log [A-]/[HA] ......... Henderson-Hasselbach
equation
The dissociation of a proton from the protonated
form of the amino group of an amino acid can also be described
in terms of Ka,
as follows:
RNH3+
+ H2O ===== RNH2
+ H3O+
[RNH2] [ H3O+]
Ka = ------------------
[RNH3+] [H2O]
In order to avoid the use of exponents, the extent
of proton dissociation can be expressed in terms of pKa.
Just as pH is the negative logarithm of hydrogen ion concentration,
pKa is the negative logarithm of the acid dissociation constant.
The following are useful in
calculations relating to buffering:
ml x N =
milliequivalents, or
liters x N = equivalents,
and
pH = pKa + log ([conjugate base]
/ [conjugate acid])
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