Physiological Solutions for Biochemical investigations
The body's fluids
are all electrolyte solutions,
meaning they contain ions (charged particles) such as Na+, K+,
and Cl- dissolved in water. Electrolyte solutions are necessary to
maintain water balance, pH of body fluids, nervous transmission, muscle
contraction, etc.
Physiologists
divide the solutions of the body into intracellular
fluid (ICF) which is the
solution found inside cells and extracellular
fluid (ECF) which is the solution in which cells are bathed. The ECF can be
further divided into a) plasma
volume and b) interstitial volume.
They both function to provide the cell with what it needs and protect it from
changes in the outer environment. Physiological solutions may be
different for different purposes. Some solutions contain higher solutes while
some contain less solute in the solution. There is certain terminology that is used
when trying to predict whether osmosis will occur. These terms are used when
comparing two solutions separated by a semi-permeable membrane.
·
A hypertonic solution has a greater
solute concentration than an adjacent solution after accounting for the
permeability of the membrane (after considering whether the solute will
diffuse). Cells get shrinkage in
hypertonic solutions.
·
A hypotonic solution has a lesser
solute concentration than the adjacent solution after accounting for the
permeability of the membrane. Cells get turgid and may burst in hypotonic solutions (e.g. RBC
get hemolyzed in distilled water)
·
An isotonic solution is when the two
solutions have the same solute concentration. Again, the permeability of the
membrane has to be accounted for first. There are no changes in cell activity
in isotonic solution. These types of solutions are best for biochemical investigations.
In biochemical
investigations physiological solutions play vital roles in the protection of
biomolecules under the study. Without use of physiological solutions, it is
very difficult to achieve the active form of cellular molecules. Besides
protecting biomolecules, it also protects whole cell. So in order to study
biomolecules, suitable physiological solutions must be used during extraction
and purification of such molecules. Some of the important physiological
solutions are normal saline and buffers. In case of biochemical investigations,
buffers are most commonly used.
Buffers
Proteins have a pH dependent charge and many of the properties of proteins
change with pH. Consequently, in working with proteins, it is important to
control the pH. This is achieved by the use of buffers, and so at the outset it
is important to have some insight into buffers, to know which buffer to use for
any particular purpose, and how to make up the buffer.
Buffers are solutions of weak acids or bases and their salt(s), which resist
changes in pH. Weak acids and bases are distinguished from strong acids and
bases by their incomplete dissociation. In the case of a weak acid the
dissociation is:-
HA ® H+ + A-
From above Equations i and ii are forms of the Henderson-Hasselbalch
equation, which can be written in a general form as:-
Pka = pH - log [basic species]/[acidic species]
Pka = pH - log [basic species]/[acidic species]
From which it can be seen that, when [basic species] = [acidic species],
then, 
It will be noticed that when pH = pKa, the solution resists changes in pH,
i.e. it functions best as a buffer in the range pH = pKa ± 0.5.
In acetate buffer, CH3COOH is the acidic species in this buffer and CH3COO-
is the basic species. It may be observed that a solution of acetic acid itself (CH3COOH)
will have a pH less than the pKa of acetic acid. Conversely, a solution
containing only sodium acetate will have a pH greater than the pKa of
acetic acid. It is important to understand this point in order to appreciate
how to make an acetate buffer using the approach described in figure 1.
Figure 1: Schematic titration curve of acetic acid
A tri-protic acid, such as phosphoric acid will yield a titration curve having
three inflexion points (Figure 2), corresponding to the three pKa values of
phosphoric acid.
Figure 2: Schematic titration curve of phosphoric
acid
For most biochemical purposes, pKa2 is of greatest interest, since it is closest
to the pH of the extracellular fluid of animals.
Note that:-
At pKa2, [NaH2PO4]
= [Na2HPO4]
At pH< pKa2 [NaH2PO4]
> [Na2HPO4]
At pH> pKa2 [NaH2PO4]
< [Na2HPO4]
Put another way, a solution __ NaH2PO4 will have a
pH less than pKa2 and a solution of Na2HPO4
will have a pH greater than pKa2. It is important to
understand this point in order to appreciate how to make a phosphate buffer
using the approach described below.
Making a buffer
A simpler method for preparing is as follows:-
1. Choose the buffer:
A buffer works best at its pKa, so the first step is to choose a buffer
with a pKa as close as possible to the desired pH.
2. Identify the buffering species:
As described above, a buffer consists of two components: a weak acid and
its salt or a weak base and its salt. The second step is thus to identify the
species which will constitute the buffer. For example, in the case of an
acetate buffer, the buffering species are CH3COOH and CH3COONa.
In a phosphate buffer at pKa2, the buffer species are NaH2PO4
and Na2HPO4.
3. Identify whether the buffer is made from an acid or a base:
The two buffer examples given above are made from acids, acetic acid or phosphoric
acid. In the case of phosphate buffer at pKa2, the acid is NaH2PO4.
An example of a buffer made from a base is Tris/Tris- HCl, which buffers best
at pH 8.1, the pKa of Tris.
4. Choose the species that gives no by-products when titrated:
Almost all buffers can be made up by weighing out one component, dissolving
in a volume just short of the final volume, titrating to the right pH, and making
up to volume. It is not necessary to make up separate solutions of the
two buffer constituents - the required salt can be generated in situ by
titrating the acid with an appropriate base - or vice versa in the case
of a buffer made from a base. [Remember: Titrate an acid “up” (i.e. with a
strong base) and titrate a base “down” (i.e. with a strong acid)].
Remember, acid +
base = salt + water
and, a
buffer = (acid + its salt ) or (base + its salt)
The term “its salt” is important. For example, if we wanted to make an
acetate buffer, it is easy to
identify that this buffer is made from acetic acid and its salt, say, sodium
acetate. But,
Q: Could the required mixture of CH3COOH and CH3COONa
be made by titrating a solution of CH3COONa to the correct pH with
HCl?
A: No! Because the reaction in this case is:-
CH3COONa + HCl ® CH3COOH + NaCl
and the resultant solution contains NaCl, which is an unwanted byproduct and
which is not a salt of acetic acid (i.e. it is not “its salt”).
On the other hand,
Q: Could the required mixture be made by titrating a solution of CH3COOH
with NaOH?
A: Yes! The reaction in this case is:-
CH3COOH +
NaOH ® CH3COONa + H2O
Which yields only the salt of acetic acid and water, i.e. there are no
byproducts. Similarly, in the case of a phosphate buffer, if one chooses Na2HPO4,
the pH of a solution of this salt will be higher than pKa, and this will
require titration with an acid. If one chooses HCl, the reaction will be:-
Na2HPO4
+ HCl ® NaH2PO4 + NaCl
Which yields NaCl as an unwanted by-product. (And if one chooses NaH2PO4,
this will change the phosphate molarity.) However, if one starts with NaH2PO4,
the pH of a solution of this salt will be lower than pKa, and this will
require titration with a base. If one chooses NaOH, the reaction will be:-
NaH2PO4
+ NaOH ® Na2HPO4 + H2O
Which yields only the desired salt (Na2HPO4) and
water.
For a Tris buffer, one should start with the free base and titrate this with
HCl to yield the salt of Tris, Tris-HCl.
5. Calculate the mass required to give the required molarity:
Having settled on the single buffer component to be weighed out, calculate
the mass required to give the required molarity, when finally made up to volume.
For example, the molarity of a phosphate buffer is determined by the molarity
of the phosphate moiety (-PO43-), which does not change
when NaH2PO4 is titrated to Na2HPO4.
If a liter of a 0.1 M buffer is required, then 0.1 moles of NaH2PO4
can be weighed out.
6. Add all other components titrate and make up to volume
Buffers often contain ingredients other than the two buffering species. For
ion-exchange elution the buffer might contain extra NaCl, and buffers often
contain preservatives such as NaN3 or chelating agents such as EDTA.
Except for NaN3, these should all be added before the titration. All
constituents should be dissolved in the same solution to just less than the
final volume, i.e. a volume must be left for the titration but the final
dilution after titration should be as small as possible. (The
Henderson-Hasselbalch equation predicts that the pH of a buffer should not
change with dilution, but this is only true over a small range, due to
non-ideal behavior of ions in solution.) Finally the solution is titrated to
the desired pH and made up to volume. NaN3 should be added after
titration as it liberates the toxic gas, HN3, when exposed to acid.
Manganese salts should also be added after adjustment of the pH as these may
form irreversibly insoluble salts at pH extremes.
Buffers of constant ionic strength
Besides pH, which influences the sign and magnitude of the charge on a
protein, proteins are also influenced by the specific ions present in solution
and by the solution ionic strength. In a buffer, the pH and the ionic strength
are related. The Henderson-Hasselbalch equation, for a buffer made from an
acid, is:- 
The ionic strength of the buffer is a function of the [salt]. Therefore, in
this case as the pH rises, the buffer ionic strength also rises. Ionic strength
is also a function of the molarity of the buffer.
For a buffer made from a weak base, the relevant form of the Henderson-Hasselbalch
equation is:- 
In this case, therefore, the ionic strength increases as the pH decreases and
the relationship.
Applications and uses:
1.
Provide constant pH: Physiological solutions such as buffers
are used to maintain the constant pH environment surrounding the cell mass
protecting pH sensitive molecules. Hence it is also used in cell culture and
microbial cultivation. E.g. It protect pH sensitive molecules such as enzymes, proteins, nucleic acids (DNA and
RNA) etc.
2.
Provide constant ionic strength: Buffer solutions also
maintain the constant ionic strength of solution or environment. The pH of the
buffers is the function of ionic strength of specific molecules (Phosphate
moiety in phosphate buffer) and they are not utilized by microbes so it
maintains constant ionic strength of solution and protects the denaturation of
biomolecules.
3.
Naturally present: Buffers are not only used in biochemical
investigations but also found in many biological systems like in soil, blood,
cytoplasm etc. The clay and humus particles act as buffer in soil while
bicarbonates and free amino acids are buffering components in blood.
4.
Preserve energy loss: By protecting denaturation of
biomolecules, it minimizes the energy loss for synthesis of same biomolecules.
5.
It protects cell organelles and vital organs in higher organisms: E. g.
uptake of physiological solutions prevents the loss of electrolytes and
dehydrations hence protecting kidney and liver.
6.
It is required for effective activity for biomolecules:
Phosphate
buffer for ® Proteins
Tris
buffer for ®
Nucleic acids (DNA and RNA)
Bi-carbonate
buffer ® Blood cells and Blood molecules
Acetate
buffer with glucose o r sucrose ® tissue homogenization-etc.
