Otto Folin.

Laboratory manual of biological chemistry with supplement online

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Hamilton Kuhn Professor of Biological Chemistry in Harvard Medical School









COPYRIGHT, 1916, 1919, BY




For this edition the manual has been very largely rewritten.
Much painstaking research is represented in the revised analytical
methods here given for the first time in book fqrm, and if the
directions are followed these methods give reliable results.

Much valuable assistance has been received from Dr. C. H.
Fiske in connection with this revision.




This manual of biological chemistry for medical students in
Harvard Medical School has been revised annually for the past
seven years, and it is believed now to meet our needs sufficiently
well to warrant publication.

For many years I have been interested in the development of
analytical methods applicable to metabolism investigations. The
most serviceable of my older methods and some of the newer
methods have been taught to our medical students ; these are de-
scribed in the main body of the manual. Others not heretofore
included have been incorporated in the supplement, so that nearly
all the newer methods devised in the department are now de-
scribed in this manual.

In connection with the revisions referred to above I am in-
debted for valuable help to W. R. Bloor, W. Denis, C. J. Farmer,
L. J. Morris, F. B. Kingsbury, F. S. Hammett, R. D. Bell, and
C. H. Fiske, as well as to my older friend, P. A. Shaffer.








V PROTEINS * ... 71


VII BLOOD . . 135



X BILE 145






Equivalent and Normal Solutions. Since the molecular weight
of sodium hydroxid (NaOH) is 40 and that of hydrochloric acid
(HC1) is 36.46, it follows that 40 g. of the former contain the
same number of molecules as 36.46 g. of the latter. If 40 g. of
sodium hydroxid and 36.46 g. of hydrochloric acid are each dis-
solved in pure water sufficient to make one liter of solution, each
liter will contain the same number of dissolved molecules.

It will take a little less than one liter of water to make a liter of
solution because the dissolved substance takes up some space. A nor-
mal sodium hydroxid solution contains four per cent, of sodium
hydroxid. By per cent, in the case of solutions is usually meant the
amount of substance present in 100 c.c. of solution.

Mixing equal volumes of two such solutions is, therefore, the
same as bringing together practically the same number of the
two kinds of molecules, and the result is -the instantaneous and
essentially complete transformation into sodium chlorid (and

X NaOH + X HC1 = X NaCl + X H 2 O

If either or both of the solutions should first be diluted with a
considerable bulk of, pure water, the result on mixing the two
would be the same, for the extra amount of water present takes
no part in the reaction (except to the extent of absorbing a part
of the heat set free).

The two solutions are equivalent. They also happen to be nor-
mal solutions. Tire hydrochloric acid is normal because it con-


tains i g. of active or replaceable hydrogen per liter' of solution,
and not because it contains the same number of grams of HC1
per liter as there are units in the molecular weight. The sodium
hydroxid solution is normal because it is equivalent to a solution
containing one gram of replaceable hydrogen per liter.

The molecular weight of sulphuric acid is 98. A sulphuric acid
solution containing exactly 98 g. per liter contains, therefore, the
same number of molecules per unit volume as the sodium hy-
droxid solution containing 40 g. per liter. But one molecule of
sulphuric acid requires two molecules of sodium hydroxid for the
formation of the neutral salt, sodium sulphate, because the sul-
phuric acid molecule has two replaceable hydrogen atoms. The
solutions are not equivalent, for the sulphuric acid contains 2 g.
active hydrogen per liter. It is exactly twice as strong as the
sodic hydrate solution; it is a 2 normal -solution.

On the basis of the above description of what constitutes a normal
solution, Calculate the number of grams per liter in tenth normal sul-
phuric acid (.iN H 2 SO 4 ), fifth normal hydrochloric acid (.2N HC1),
half normal oxalic acid (-5N C 2 H 2 O 4 , 2H 2 O), fourth normal acetic
acid OsN CH 3 COOH), half normal sodic hydrate (.5N NaOH),
twentieth normal barium hydrate (.05N Ba(OH) 2 ), fifth normal
ammonium hydrate (.2N NH 4 OH).

Atomic weights of some of the more important elements : Arsenic
(As) 74.96, Barium (Ba) 137.37, Bromin (Br) 79.92, Calcium (Ca)
40.09, Carbon (C) 12, Chlorin (Cl) 35.46, Copper (Cu) 63.57,
Hydrogen (H) 1.008, lodin (I) 126.92, Iron (Fe) 55.85, Lead (Pb)
207.1, Magnesium (Mg) 24.32, Manganese (Mn) 54.93, Mercury
(Hg) 200.6, Nitrogen (N) 14.01, Oxygen (O) 16, Phosphorus (P)
31.04, Potassium (K) 39.1, Sulphur (S) 32.07, Tin (Sn) 119,
Tungsten (Wo) 184, Uranium (U) 238.5, Zinc (Zn) 65.37.

The same description of normal solutions applies to other sub-
stances than acids and alkalis, as, for example, reducing and oxi-
dizing substances such as potassium permanganate, potassium bi-
chromate, iodin, cupric hydrate, stannous chlorid. A normal so-
lution is here one capable of liberating I g. of reducing hydrogen
(or of giving off exactly sufficient oxygen to oxidize one gram of
hydrogen) per liter. Potassium permanganate, for example, in
the presence of sulphuric acid and some easily oxidizable sub-
stance is decomposed as follows :

2KMnO 4 + 3H 2 SO 4 = K 2 SO 4 + 2MnSO 4 + 3H 2 O + 50.


As the two permanganate molecules liberate oxygen enough for
ten hydrogen atoms it takes only one fiftieth of the molecular
weight expressed in grams (3.161 g.) to make one liter of tenth
normal solution.

The calculation of what constitutes normal or equivalent solu-
tions of any reagent is not very difficult provided the equation
representing the chemical reaction involved is thoroughly clear.

To determine whether a given unknown solution is acid or
alkaline it is usually sufficient to dip a piece of delicate violet
colored litmus paper into it. (If the solution is acid the test
paper turns red ; if alkaline it turns blue.) Litmus, the substance
with which the paper has been impregnated, is a complex organic
product, and is one of the most familiar representatives of a
most useful class of organic compounds which are so sensitive
to acids or alkalis, or both, that they clearly and unmistakably in-
dicate the presence of free acid or alkali even when the amounts
present are so small as to be practically unweighable. By means
of such indicators and accurate measuring instruments (measur-
ing flasks, burets, and pipets), it becomes a simple matter to
determine (by titration) the relative concentration or equivalence
of acid and alkaline solutions. By their help it is possible to pre-
pare with very little labor normal or tenth normal solutions, even
of acids or alkalis which cannot be weighed on the balance, as for
example, hydrochloric acid and ammonia, both of which are gases.
Before this can be done we must, however, possess one normal
or standard solution prepared from some substance which can
be weighed.

Volumetric analysis consists of measuring the value of an un-
known solution in terms of another the value of which is known
(titration). The known solution is prepared directly or indirect-
ly by the help of the analytical balance, and the first step in any
kind of volumetric analysis is the preparation of the standard
solution by means of which the values of others are to be deter-

Every student who has had no experience in the use of the ana-
lytical balance must consult the instructor before proceeding. He
should also ask for instruction as to the proper use of measuring
flasks, pipets, and burets before using them. He must particularly
learn when the presence of unmeasured quantities of water does not
interfere with the accuracy of the work and when a single drop of


unmeasured water introduces a perceptible error. (See Button's
Volumetric Analysis, Part i "Instruments and Apparatus.")

All the common mineral acids and strong alkalis contain so
much water that it is in practice not feasible to weigh out with
sufficient accuracy the theoretical quantity required for a standard
solution of acid or alkali. The carbonates of sodium or calcium
(or the carbonates of sodium or potassium, obtained by ignition
of the corresponding oxalates) give exceedingly accurate results.
Oxalic acid is very serviceable as starting material for the prepa-
ration of standardized solutions of acids and alkalis if it is pure
and has lost none of its water of crystallization.

1. Calibrations. The volumetric ware now available is not
always accurate enough for the work of this course. Some
approximate calibrations are therefore necessary.

At least i pipet should be calibrated by weight. Clean a 20
c.c. pipet with '^cleaning fluid" and rinse. Weigh a clean and
empty but not necessarily dry 100 c.c. volumetric flask within
an accuracy of 2 mg. Fill the pipet, adjust the lower part of
the meniscus exactly at the mark, and transfer the contents to
the weighed flask allowing to drain for 15 seconds against the
inside of the flask. Weigh again to within an accuracy of 2 mg.

1 c.c. of water may be assumed to weigh 997 mg. The slight
fluctuations due to variations in room temperature may be neg-
lected. Calculate the correct volume of the pipet. Clean a glass
stoppered buret with cleaning fluid, rinse, fill to the mark with
distilled water, empty down to the 25 c.c. mark and let drain for

2 minutes. Then adjust the meniscus exactly to the 25 c.c. mark
and be sure that the sides of the buret above the water are
entirely free from drops of adhering water. Now run in 20 c.c.
from the calibrated pipet. Compare the reading obtained with
the correct value. If the divergence seems large consult the
instructor. Next calibrate the 5 c.c. pipet by transferring its
contents 5 times to the buret, beginning with the meniscus in the
latter at the 25 c.c. mark. Record each reading but note particu-
larly where the fourth comes. If the 5 c.c. pipet is inaccurate
make a new temporary mark on the stern and repeat. Inciden-
tally record the value of each 5 c.c. portion of the buret. The J&
10 c.c. and the 25 c.c. pipets can also be calibrated by the help^
of the buret.


For calibrating the larger pipets, calibrate first I dry 100 c.c.
volumetric flask with the most accurate 20 or 25 c.c. pipet, mak-
ing if necessary a new temporary mark on the stem of the flask.
When this flask is again dry use it for checking up .the values
of a 50 or a 100 c.c. pipet. The volumetric flasks larger than
100 c.c. need not be calibrated. If a flask is rinsed with a little
alcohol and left to drain over night it will usually be found to
be perfectly dry the following day.

2. Preparation of .5N Oxalic Acid (500 c.c.). The usefulness
of oxalic acid as a starting point for the preparation of standard
acids and alkalis is due entirely to the fact that it can be obtained
chemically pure and in condition suitable for direct weighing. Oxalic
acid is, however, not a strong enough acid to titrate well with all
the common indicators, and it is therefore not serviceable for acidi-
metric titrations in general. But by means of oxalic acid and with
phenolphthalein as indicator, standard solutions of a strong alkali,
like caustic soda, can be obtained, and by means of the latter stand-
ard solutions of the stronger mineral acids can then be prepared.
. The reason w4iy the strong acids and alkalis give more accurate
and reliable results is the fact that the salts which they form when
neutralized are no^ appreciably hydrolyzed by water into acid and
base, as are the corresponding salts of the weaker acids and bases.
The zone of neutrality to different indicators is therefore more
sharply defined, and corresponds more nearly to the point repre-
sented by the presence of exactly equivalent amounts of acid and

Weigh accurately (to the fourth decimal) a small, clean, and
dry beaker or large crucible. Then add to the weights on the
balance pan 15.7560 g., and add oxalic acid to the vessel on the
other side until exact equilibrium is reached. Dissolve in dis-
tilled water this oxalic acid without the loss of a single crystal.
The acid dissolves rather slowly. The solution is, therefore, best
made in a beaker by the aid of gentle heating with about 250
c.c. water. Transfer every drop of the solution to a measuring
flask (500 c.c.), carefully rinsing the last traces from the beaker
into the flask by means of successive small amounts of cold
distilled water. Cool the flask in running tap water until the
contents of the flask have reached the room temperature. (If
a thermometer is used it must be rinsed carefully before it is
removed from the flask.) Fill up with water until the lower side
of the "meniscus" is exactly even with the 500 c,c t mark. Stop-.


per the flask, and invert several times (30-40) so that the solu-
tion is thoroughly mixed. Transfer to a clean, dry bottle; label
and preserve.

Using a strong base like sodium hydroxid and a sensitive indi-
cator like phenolphthalein for the titration, it is possible to obtain
quite reliable and accurate results with oxalic acid. The volu-
metric determinations involved in metabolism studies and urine
analysis are, however, extensively based on titrating ammonia,
which is a very weak base. Phenolphthalein, because of its high
degree of sensitiveness to weak acids and its lack of sensitiveness
to weak bases, is useless in titrations of ammonia. The oxalic
acid and the phenolphthalein are therefore used only for the pur-
pose of securing a standard alkali solution.

3. Preparation of Standardized Sodium Hydroxid. The so-
dium hydroxid used for titrations must be as free as possible
from carbonates, because otherwise the solutions will not have
the same titrating value with all the common indicators. Sodium
hydroxid absorbs rapidly carbonic dioxid from the atmosphere
and should therefore not be exposed to the air more than is un-
avoidable. As the carbonates are insoluble in very strong sodium
hydroxid solutions, clear saturated solution should be used as
starting point for the preparation of standard solutions.

Transfer about 60 c.c. of clear saturated sodium hydroxid
solution to a large bottle and add 1,200 to 1,500 c.c. of water.
To determine the exact value of this solution it is only necessary
to find out how much of it is required for the neutralization of
a known volume of the half normal oxalic acid solution.

Rinse the 20 c.c. pipet with the oxalic acid solution and then
measure 20 c.c. into a beaker or flask. Add two drops of indi-
cator (i per cent, alcoholic solution of phenolphthalein).

Rinse a buret with the alkali, fill it and cover with a test tube.
After carefully adjusting the meniscus of the solution to the
zero point, run it into the oxalic acid solution more and more
cautiously toward the end until finally one single drop produces
a definite and stable end point. Note the volume of alkali re-
quired (within 0.05 c.c.). Repeat the titration until two succes-
sive ones give exactly the same value.

From the titration figure obtained calculate the normality of
the solution and how much of it must be taken for the prepara-
tion of i liter of tenth normal alkali.

As a check on the work determine the normality of an un-


known hydrochloric acid solution (furnished), using as indi-
cator (a) phenolphthalein, (b) alizarin red (2 drops of I per
cent, aqueous solution). A dated and signed report on the
unknown should be handed in before making the tenth normal
alkali. Label and preserve the standardized alkali solution.

4. Standardized Hydrochloric Acid. Concentrated hydro-
chloric acid is approximately a 10 N solution of HC1. With a
cylinder transfer 60 c.c. of strong hydrochloric acid to a large
bottle; add 1,500 c.c. of water and shake very thoroughly. It is
preferable but not absolutely necessary to let the shaken solution
stand over night before titrating.

Titrate this acid in the same way as the oxalic acid solution,
but using only alizarin red as indicator. Calculate the normality,
and how much of it must be taken for the preparation of I liter
of tenth normal acid. Label and preserve.

5. Tenth Normal Acid and Alkali. From the standardized
solutions of acid and alkali prepare I liter of tenth normal hydro-
chloric acid arid I liter of tenth. normal alkali. Titrate the acid
so prepared (20 c.c.) with the tenth normal alkali. The two
should be equivalent. Determine the normality of an unknown
acid with the tenth normal alkali. Hand in a dated and signed
report giving the value obtained for the unknown and giving also
the titration figures for the tenth normal acid.

Label and preserve the tenth normal solutions. The two alkali
solutions do not always keep their value unchanged because more
or less alkali is given off by the glass containers. The hydro-
chloric acids solutions keep indefinitely. If discrepancies are
found later between the acid and the alkali, the acid should be
taken as correct.

6. Strong and Weak Acids; the Use of Different Indicators.
(A) Titrate 25 c.c. tenth normal hydrochloric acid with the
tenth normal alkali, using as indicator (a) phenolphthalein (b)
methyl orange (c) alizarin red. Repeat the above mentioned
three titrations in the presence of 10 c.c. ammonium chlorid
solution (2 per cent.). Repeat the titration with each indicator
using in place of the hydrochloric acid (a) 25 c.c. .iN phosphoric
acid (b) 25 c.c. .iN lactic acid.

Record the titrations in tabular form :

Phenolphthalein: Methyl orange: Alizarin red:
c.c. End point.* c.c. End point.* c.c. End point.*

* Sharp, fair or indeterminate.



HC1 NH 4 C1 :

Oxalic acid :

(B) Dilute 10 c.c. tenth normal hydrochloric acid to 100 c.c.,
making an approximately o.oi N solution. (Measuring cylin-
ders are accurate enough for the dilutions referred to here.)

From this o.oi N solution prepare four 100 c.c. portions of
more dilute acids, viz. : o.ooi N ; 0.0004 N ; o.oooi N ; o.ooooi N.
Arrange in a row four test tubes, as nearly as possible of
the same size, and transfer to each one 5 c.c. of one of the four
dilute acid solutions. To the contents of each tube add one drop
(no more) of a 0.15 per cent, alcoholic solution of tetrabromo-
phenolsulfonephthalein ("bromphenol blue"), and compare the
colors. The approximate hydrogen concentrations of these solu-
tions are as follows :

C H p H *

o.ooiN... io- 3 3.0

O.OOO4N 4xio- 4 3.4

o.oooi N io- 4 4.0

o.ooooi N 10-* 5.0

Add the same amount of indicator to (a) 5 c.c. o.ooi N lactic
acid, (b) 5 c.c. o.ooi N acetic acid, (c) 5 c.c. o.ooi M mono-
potassium phosphate. Determine the approximate p H of each
of these three solutions by comparing their colors with the dilute
hydrochloric acid solutions. Although the total acid concentra-
tion is the same in the o.ooi N solutions of hydrochloric, lactic,
and acetic acids, and monopotassium phosphate (an acid salt),
the hydrogen ion concentration (and therefore the degree of
dissociation) is obviously different in each case. In any such
series of acid solutions of the same total concentration (o.ooi N
in this instance), the hydrogen ion concentration is less (and the
p H greater) the weaker the acid. The strength of an acid is
measured by its dissociation constant (k), a figure approximately
equal to the hydrogen ion concentration of a solution of the acid
that has been just half neutralized. The dissociation constants
of the three weak acids used in this experiment are given below,

*The hydrogen exponent (p a ) is the logarithm of the hydrogen ion
concentration (C H ) with the minus sign omitted,


along with those of other acids and bases of biological impor-

Hippuric acid 2.2x10-4 Uric acid 1.5x10-6

Acetoacetic acid . . . 1.5x10-4 Carbonic acid .... 3.0x10-7

Lactic acid 1.4x10-4 Primary phosphate 2.0x10-7

Acid oxalate ...... 3.0x10-5 Boric acid 6.6x10-10

j3 -Hy droxybutyric

acid 2.0x10-5

Acetic acid 1.8x10-5 Ammonia ...... .,. . 1.8x10-5

1. Acidity of Gastric Contents. The acidity of the normal
stomach contents is due almost wholly to hydrochloric acid. In
pure gastric juice, the concentration of hydrochloric acid is about
0.15 N, but the acidity of the material usually found in the
stomach is less, as a result of dilution and partial neutralization.
When, under abnormal conditions, the concentration of hydro-
chloric acid becomes very low, certain microorganisms are able
to grow in the stomach contents, producing lactic acid. It is
nevertheless an easy matter to distinguish between a relatively
low concentration of hydrochloric acid and a relatively high con-
centration of lactic acid, since the latter is a much weaker acid.

To 5 c.c. o.oi N hydrochloric acid in a test tube add just one
drop of a 0.4 per cent, alcoholic solution of thymolsulfonephtha-
lein. Add the same amount of indicator to (a) 5 c.c. o.ooi N
hydrochloric acid, and (b) 5 c.c. o.i N lactic acid. Compare the
colors. The hydrogen ion concentration of the o.i N lactic acid
should be less than that of the o.oi N hydrochloric acid.

8. Colorimetric Determination of Hydrogen Ion Concentration.

The hydrogen ion concentration of most biological fluids is
considerably less than in the solutions tested in the preceding
experiments, and dilute hydrochloric acid solutions cannot be
used here as standards, owing to the ease with which the p H is
changed by slight contamination. Instead, it is necessary to have
a series of standard buffer mixtures, whose p H is not readily

A suitable set of stock solutions from which to prepare such
standards is : 0.2 M monopotassium phosphate, 0.2 M acetic acid,
0.2 M boric acid (containing also 0.2 M potassium chlorid), and 0.2
M sodium hydroxid. The sodium hydroxid solution must be prac-


tically free from carbonate, and should not contain calcium or bari-
um. The compositions of the standard mixtures (diluted to 200 c.c.
in each case) are given in the table below. These mixtures, once
made up, can be relied upon for only about one week, but the stock
solutions from which they are prepared should keep indefinitely in
receptacles of resistance glass, except the sodium hydroxid solution,
which will gradually increase in strength unless kept in a paraffined

The determination is carried out as follows: With a pipet
transfer two c.c. of the unknown solution to a measuring cylin-
der and add water to make the total volume 20 c.c.f Mix, and
add one drop (no more) of phenol red solution. Compare the
color with the set of standards. f In case the color is beyond
the limits for phenol red on either side, repeat with the next
indicator in order (see table), until the unknown has been cor-
rectly matched against one of the standards. The p H reading
may be made to one-tenth unit by adding or subtracting o.i in
case the color lies definitely between two consecutive standards.

Determine, in the manner described, the hydrogen ion concen-
tration of two unknowns (supplied). The same method will
later be applied to urine.

Indicator: METHYL RED.* ,'

50 c.c. 0.2 M CH 3 COOH and 23.0 c.c. 0.2 M NaOH 4.6

50 c.c. 29.0 c.c. 4.8

50 c.cj 34.5 c.c. 5.0

50 c.c. 38.5 c.c. 5.2

50 c.c. J 42.5 c.c. 5.4

50 c.c. 45.0 c.c. 5.6
*Dimethylaminoazobenzene-o-carboxylic acid (0.4 per cent, alcoholic
solution; use one drop).


50 c.c. o.2MKH 2 PO 4 and 3.7 c.c. 0.2 M NaOH 5.8

50 c.c. fc ft 5-7 c.c. 6.0

50 c.c. " 8.6 c.c. 6.2

50 c.c. " 12.6 c.c. 6.4

50 c.c. " 17.8 c.c. 6.6
tDibromo-o-cresolsulfonephthalein (0.04 per cent, aqueous solution
of monosodium salt; use 2 drops).

*In this work, pipets should never be blown out, and water should

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