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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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