Exercise 5
Acid – base balance in organism, buffers, colloids, dialysis.
Theoretical part
1. Diffusion, osmosis and osmotic pressure
Diffusion is a process in which molecules, ions or colloidal particles migrate from a region of
high to a region of low concentration (see Figure 1.). The kinetic energy of diffusing
molecules keeps them in motion. Diffusion is of a great importance in biology – all the
nutritive substances are distributed to the cells of a living organism by diffusion and also the
waste products of metabolism are removed by diffusion. Diffusion process is also used to
distribute substances of physiological importance, such as water, oxygen, inorganic ions,
enzymes, hormones and vitamins.
Diffusion. Example where the concentration of an ion
is 100% on one side of the membrane (blue line) and
0% on the other. This imbalance is maintained because
the membrane is impermeable to that ion.
Here the ions have diffused across the membrane
because it has become permeable to that ion (dotted
black line) and the concentration on either side is
50/50. Equilibrium has been reached.
Figure 1. Diffusion.
Osmosis is the passage of water from a region of high water concentration through a semipermeable membrane to a region of low water concentration. Semi-permeable membranes are
very thin layers of material (cell membranes are semi-permeable) which allow some
molecules to pass through them but prevent other molecules from passing through.
Cell membranes will allow small molecules like oxygen, water, CO2, NH3, amino-acids, etc.
to pass through. Cell membranes will not allow larger molecules like sucrose, starch, protein,
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etc. to pass through. Similarly ions (like Na+, K+, Ca2+, H+) do not diffuse spontaneously
through the semipermeable membrane.
Figure 2. Osmosis.
The osmotic pressure is the pressure required to stop osmosis through a semipermeable
membrane between a solution and pure solvent. If a more concentrated sugar solution is
placed on one side of a living membrane than on the other side, both the sugar and the water
will diffuse through the membrane, but the water will diffuse faster and will pass in a
direction opposite to that of the sugar, a greater osmotic pressure will develop on the side of
the membrane having the solution of greater concentration. The situation will continue until
the concentration of sugar (= the osmotic pressures) becomes the same at each side of the
membrane.
Osmotic activity is expressed by equation:
 = TRiM
 - osmotic pressure [atm]
M – molarity [mol/l]
R – gas constant [0,083 l atm/mol K]
T – temperature [K]
i – van’t Hoff’s factor (for nonelectrolytes is equal 1; for ionizable substances is determined
by the number of particles formed by the ionization process).
Solutions having the same osmotic pressure (= osmolarity) are called isotonic solutions.
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Physiological solutions are isotonic (but not isoionic) to the tissues of the body.
A solution is said to be hypotonic to another solution when it has a lower osmolarity, and the
solution is hypertonic to another solution when it has a higher osmolarity.
A.
B.
C.
Figure 3. Diagram showing changes in blood cells due to osmosis. A) Erythrocyte in
hypotonic solution (e.g. distilled water) swells and bursts ( hemolysis). B) Erythrocyte in
isotonic solution (e.g. 0,9 % NaCl) shows no volume change. C) Erythrocyte in hypertonic
solution (2 % NaCl) shrinks and becomes crenated.
The concentration of salts in the blood or lymph of animals is approximately equivalent to
0,9 % sodium chloride. The living tissues undergo no change due to osmosis in a solution of
that strength, therefore it is called a physiological solution.
Active transport is the movement of a molecule across a membrane or another barrier that is
driven by energy other than stored in the concentration gradient or the electrochemical
gradient of the transported molecule. This type of transport requires usually the expenditure of
ATP and the help of specific transport proteins. In this way can even large molecules or ions
can be channeled through the membrane.
The sodium-potassium pump is a good example of active transport of molecules across a
membrane.
Figure 4. The sodium – potassium pump.
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● The membrane has protein (or enzyme) channels, or gaps, which forms a transmembrane
pump. These pumps use energy-storing molecules called adenosine triphosphate (ATP)
● ATP actively pumps 3 Na+ ions out of the cell, at the same time pumping 2 K+ into the cell.
● After a while, a ionic concentration gradient is generated across the membrane, whereby
more Na+ ions are outside and more K+ are inside
● Because of diffusion, the tendency is for Na+ ions to travel back to the inside, and vice
versa for K+ ions. There are nongated channels in the membrane that permit the passage of
some Na+ ions back into the neuron, and K+ ions out of the neuron (again, using diffusion to
achieve a concentration equilibrium), however, the membrane is not very permeable to Na+
ions. Hence many more K+ ions leave the cell than Na+ ions enter. This causes an excess of
negative charge in the cell.
● The K+ ions continue to leak out until there is an equilibrium reached between the
concentration gradient and the electric potential (i.e., the attraction of K+ positive ions back to
the negatively charged intracellular fluid)
2. Colloids
In the middle, between the suspensions (diameter of particles  1000 nm) and solutions (d 
10 nm) is the large group of mixtures called colloidal dispersions, or simply colloids, in which
a dispersed (solutelike) substance is distributed throughout a dispersing (solventlike)
substance. The dispersed colloidal particles are larger than a simple molecules but small
enough to remain distributed and not to settle out. A colloidal particle has a diameter between
1 and 1000 nm (10-9 m to 10-6 m) and may contain many atoms, ions or molecules.
A colloidal solution has two phases, the dispersed and the continuous (dispersing) phase.
The dispersed phase is the particular matter held in solution. The dispersing medium (solvent,
for example water) is the continuous phase. Colloids are classified according to whether
the dispersed and dispersing substances are gases, liquids or solids (Table 1.).
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Table 1. Types of colloids
Phase of Colloid
Gas
Gas
Gas
Liquid
Liquid
Liquid
Solid
Solid
Solid
Dispersing
Phase
Gas
Gas
Gas
Liquid
Liquid
Liquid
Solid
Solid
Solid
Dispersed Phase
Colloid Type
Example
Gas
Liquid
Solid
Gas
Liquid
Solid
Gas
Liquid
Solid
--Aerosol
Aerosol
Foam
Emulsion
Sol
Solid Foam
Solid Emulsion
Solid sol
none
Fog
Smoke
Whipped Cream
Milk
Paint
Marshmallow
Butter
Ruby Glass
Characteristic properties of colloids are:
1) colloids can be poured through an ordinary filter paper, but do not pass through any
semipermeable membrane (e.g. a cell membrane).
2) Tyndall effect – when a beam of light is passed through a colloidal solution, the pathway of
the light appears as a more highly illuminated region because the light ray are reflected by the
colloidal particles (like a beam of light is reflected by dust particles in the air).
3) Brownian motion – a characteristic movement in which the colloidal particles change speed
and direction erratically. This motion occurs as the colloidal particles are pushed this way and
that way by molecules of the dispersing medium. These collisions are primarily responsible
for keeping colloidal particles from settling.
4) colloids are coagulated by heat and by alcohol, that is they are converted into an insoluble
state by these agents.
Reversible coagulation of colloids is produced by an addition of a solvent, by electrolytes and
by heating.. A process that is reversed of coagulation is called peptization.
coagulation
sol  gelatine
peptization
Irreversible coagulation (denaturation) is produced by high temperatures, concentrated acids
and bases and heavy metal salts (e.g. PbSO4, CdSO4)
Colloids can be divided to:
 lyophilic (solvent-loving) colloid that has the attraction to its solvent. If the solvent is water,
the colloid is called a hydrophilic one. This kind of colloid (e.g. egg albumin in water) have
charge surfaces that interacts strongly with water.
 lyophobic (solvent-fearing, hydrophobic) colloid that does not have this attraction to its
solvent. Lyophobic colloids (e.g. AgCl or oil in water) form unstable solutions with their
solvents and are very likely to form large particles and precipitate from solution.
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Dialysis is called the separation of crystalloids (nonocolloids), e.g. glucose, NaCl, HCl, from
colloids by the passage of the crystalloids through a membrane that holds back colloids. It is
considered a practical process for the separation of crystalloids from colloids.
In medicine dialysis is a method of removing toxic substances (impurities or wastes) from the
blood when the kidneys are unable to do so. Dialysis is most frequently used for patients who
have kidney failure, but may also be used to quickly remove drugs or poisons in acute
situations. This technique can be life saving in people with acute or chronic kidney failure.
A dialyzator substitutes for renal’s activity or aids to purify the blood from metabolites and
toxic substances. The dialyzator consists of two vessels, the internal one has a semipermeable
cellophane membrane at the bottom. The colloidal solution contaminated by crystalloids (e.g.
blood) is placed in internal part, which is suspended in an external vessel of distilled water.
The crystalloids (like urea, creatine, derivatives of phenol) will pass through the cellophane
into distilled water and in this manner can be separated from the colloids.
H2O
colloid
H2O
3. Buffers can be defined as a weak acid or a weak base in the presence of its salts. Also
equimolar mixture of two salts are used for a buffer, the one with more hydrogen atoms in it is
considered to be the acid, whereas the salt with less hydrogen is considered to be the salt (see
Table below). Buffers can resist changes in pH when small amounts of strong acids or bases
are added.
Components of a buffer
Name of a buffer
a weak acid/ a salt of weak acid
acetate buffer
CH3COOH / CH3COONa
(acetic acid / sodium acetate)
bicarbonate buffer
a weak acid/ a salt of weak acid
H2CO3 / NaHCO3
(carbonic acid / sodium bicarbonate)
a weak base/ a salt of weak base
ammonia buffer
NH3 / NH4Cl
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(ammonia / ammonium chloride)
two salts: an acid/ a salt
phosphate buffer
KH2PO4/K2HPO4
(potassium dihydrogen phosphate /
potassium hydrogen phosphate)
The blood has three powerful buffer systems that protect against the introduction of acid or
base. These are (1) proteins in the plasma and hemoglobin in the red cells, (2) the phosphate
buffer and (3) bicarbonate buffer in the plasma.
The proteins are effective, because the blood contains a large amount of them. Proteins
behave as buffer because of their ability to neutralize either acid or base. Protein’s molecules
consist of amino acids that contain amine groups, which will neutralize acids, and carboxyl
groups, which will neutralize bases.
Bicarbonate acts as buffers against acidity by reacting with a strong acid to produce carbonic
acid (H2CO3), which is a weak acid. For example, if a strong acid HCl is added to a solution
containing sodium (potassium) bicarbonate, the following change occurs:
HCO3- + H3O+↔ H2CO3+ H2O
↓
H2O + CO2
In this reaction the strong acid (HCl) is replaced by H2CO3, which is so weak an acid that
there is very little change in the pH of the solution. In the same moment amount of H 2CO3 is
increasing, as much as amount of HCO3- was diminished.
The carbonic acid thus formed acts as a buffer against strong bases, neutralizing them and
producing water and the poorly ionized compound, sodium bicarbonate, which changes
the pH of the solution but slightly:
H2CO3+ OH-↔ HCO3- + H2O
A mixture of H2PO4- and HPO42- acts as a buffer as well. By analogy:
H2PO4- + OH-↔ HPO42- + H2O
HPO42- + H3O+↔ H2PO4- + H2O
And let us consider following reactions, that illustrate mechanisms of action of acetate and
ammonia buffer, respectively:
CH3COO- + H3O+↔ CH3COOH + H2O
CH3COOH + OH-↔ CH3COO- + H2O
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NH3aq + H3O+↔NH4+ + H2O
NH4+ + OH- ↔ NH3aq + H2O
The story of buffers in connection with the blood is the same for all the tissues of
the body. Without the buffers the human body could never withstand the acids produced in
normal metabolism, the excesses of acids and bases that are sometimes encountered as a result
of accidental intake of extra acid or base, or the abnormal amounts of acids or base resulting
from an unbalanced diet or disease.
pH of any buffer is defined by Henderson – Hasselbach equation:
Msalt
pH = pKa + log 
Macid
where:
pKa = -log Ka, and Ka – acid dissociation constant
Msalt – molar concentration of a salt
Macid – molar concentration of an acid
Expressions for pH of selected buffers:
 Carbonate buffer
[ HCO3-]
pH = pKa + log 
[ H2CO3]
 Phosphate buffer
pH  pK
 log

H 2 PO4
[ HPO42  ]
[ H 2 PO4 ]
 Ammonia buffer
pH  14  pOH  14  ( pK b  log
[ NH 4 Cl ]
)
[ NH 3  H 2 O]
where:
pKb = -log Kb, and Kb – base dissociation constant
or:
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pH  pK a  log
[ NH 3  H 2O]
[ NH 4Cl ]
where:
pKa = -log Ka, and Ka – acid dissociation constant
Buffer capacity is a measure of the ability of solution to resist pH change. Buffer capacity is
the number of moles of strong acid or strong base needed to change the pH of 1 Liter of buffer
solution by 1 pH unit.
The more concentrated the components of a buffer, the greater the buffer capacity. Since
the ratio of concentrations of the buffer components determines the pH, the less the ratio
changes. For a given addition of acid or base, the ratio changes less when buffer – component
concentrations are similar than when they are different. It follows that a buffer has the highest
capacity when its components are present at equal concentration, that is, when M salt/Macid = 1
, which gives pH=pKa.
Buffer capacity (A) is a ratio of acid or base added (to 1 liter of a buffer) to change pH:
A
where:
n
pH
A – buffer capacity
Δn – number of moles of added acid or base
ΔpH – pH change
Experimental
1. Preparation of buffers of known pH
2. Determination of pH of buffer solution using indicators
3. Determination of pH of buffer solution using pH-meter
4. Effect of dilution on buffer capacity
1. Preparation of buffers of known pH
Buffer solutions of define hydrogen-ion concentration may be made up from a series of stock
solutions in define proportions.
Prepare buffer solution using stock solutions of 1/15M Na2 HPO4 and 1/15M KH2PO4 or 0.2M
CH3COONa and 0.2M CH3COOH (see Table 2 and Table 3) according to assistant’s
instruction.
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Table 2. Phosphate buffer
Table 3. Acetate buffer
A-1/15M Na2 HPO4
A- 0.2M CH3COONa
B- 1/15M KH2PO4
B- 0.2M CH3COOH
nr
A[ml]
B[ml]
nr
A[ml]
B[ml]
1.
7.40
32.60
1.
7.20
32.80
2.
10.60
29.40
2.
10.60
29.40
3.
15.00
25.00
3.
14.80
25.20
4.
20.00
20.00
4.
19.60
20.40
5.
24.44
15.56
5.
24.00
16.00
6.
28.60
11.40
6.
28.20
11.60
7.
32.16
7.84
7.
31.60
8.40
2. Determination of pH of buffer solution using indicators
A) The first step in this procedure is to determine the approximate pH of the buffer solution
using indicate paper:
Treat small piece of indicator paper with 1 drop of buffer solution and compare the color
obtained with color scale. Such paper is convenient and sufficiently accurate for many
purposes. It is possible to attain a precision of 1-2 pH units in this method.
B) For attaining a precision of 0.5-0.1 pH and determine pH of buffer solution use indicator
solutions (see Table 4):
Using “serologic plate” treat a small portion (a few drops) of the buffer solution with 1 drop
of an indicator solution. Compare the color obtained with the alkaline color and acid color of
the indicator. If the color obtained is intermediate between the acid and alkaline color of the
indicator, the pH of the solution lies within the effective range of this indicator. If on the other
hand the solution shows either the full acid or full alkaline color with the indicator selected, it
is unsuitable and another indicator must be tried in a similar manner, until the pH will be
determine with 1-0.5 pH unit precision.
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Table 4. Colors and pH range of indicators
Indicator
Color at lower pH
Range of
Color at higher pH
color change
(pH)
Thymol Blue
pink-red
1.2–2.8
yellow
Töpfer’s reagent
red
2.8–4.6
yellow
Methyl Orange
red
3.1–4.4
yellow
Bromcresol Green
yellow
3.8–5.4
dark blue
Methyl Red
red
4.2–6.3
yellow
Litmus
red
5.0–8.0
blue
Bromthymol Blue
yellow
6.0–7.6
dark blue
Cresol Red
orange
7.2–8.8
purple
Neutral Red
purple-red
6.8–8.0
orange-brown
Phenol Red
yellow-orange
6.8–8.4
red-purple
Thymol Blue
yellow
8.0–9.6
purple-blue
Phenolphthalein
colorless
8.3–10.0
pink
Thymolphthalein
colorless
9.3–10.5
dark-blue
C) By use of the Henderson-Hasselbach equation calculate the pH of your buffer solution.
The values of Ka and pKa are:
KCH3COOH = 1.8×10-5 ( pKa = 4.74)
KH2PO4- = 6.2×10-8 (pKa = 7.21)
3. Determination of pH of buffer solution using pH-meter
Place buffer solution into a small beaker and introduce (carefully) both the glass and calomel
electrodes into a buffer. Read the pH. Before and after measure wash electrode (carefully)
with distilled water.
4. Effect of dilution on buffer capacity
A. Into a 4 clean test tubes measure appropriate 5 ml, 2.5 ml, 1 ml and 0.5 ml of buffer
solution and dilute all to 5 ml with distilled water.
B. Add to all test tubes 1-2 drops of appropriate indicator (pH of your buffer solution should
be near to effective pH range of indicator used, but not be within)
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C. Titrate all 4 solution with HCl or NaOH (depending on effective pH range). Note volume
of titrant used
D. Calculate buffer capacity for all 4 dilutions (according to Table 5)
Table 5. Sample presentation
nr
ml of
buffer
ml of
water
dilution
pH of
solution
(by pHmeter)
(indicator used)
color of solution
before
titration
after
titration
nr of ml of
ΔpH
used acid/ base
1.
2.
3.
4.
where:
ΔpH – pH change
Δn – number of moles of added acid or base (per 1 liter of a buffer)
A – buffer capacity
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Δn
A