VOLUMETRIC ANALYSIS
The processes of volumetric analysis are, in general, simpler than
those of gravimetric analysis and accordingly serve best as an
introduction to the practice of quantitative analysis. For their
execution there are required, first, an accurate balance with which
to weigh the material for analysis; second, graduated instruments in
which to measure the volume of the solutions employed; third, standard
solutions, that is, solutions the value of which is accurately known;
and fourth, indicators, which will furnish accurate evidence of the
point at which the desired reaction is completed. The nature of the
indicators employed will be explained in connection with the different
analyses.
The process whereby a !standard solution! is brought into reaction is
called !titration!, and the point at which the reaction is exactly
completed is called the !end-point!. The !indicator! should show the
!end-point! of the !titration!. The volume of the standard solution
used then furnishes the measure of the substance to be determined as
truly as if that substance had been separated and weighed.
The processes of volumetric analysis are easily classified, according
to their character, into:
I. NEUTRALIZATION METHODS; such, for example, as those of acidimetry
and alkalimetry.
II. OXIDATION PROCESSES; as exemplified in the determination of
ferrous iron by its oxidation with potassium bichromate.
III. PRECIPITATION METHODS; of which the titration for silver with
potassium thiocyanate solution is an illustration.
From a somewhat different standpoint the methods in each case may
be subdivided into (a) DIRECT METHODS, in which the substance to be
measured is directly determined by titration to an end-point with a
standard solution; and (b) INDIRECT METHODS, in which the substance
itself is not measured, but a quantity of reagent is added which is
known to be an excess with respect to a specific reaction, and the
unused excess determined by titration. Examples of the latter class
will be pointed out as they occur in the procedures.
MEASURING INSTRUMENTS
THE ANALYTICAL BALANCE
For a complete discussion of the physical principles underlying the
construction and use of balances, and the various methods of weighing,
the student is referred to larger manuals of Quantitative Analysis,
such as those of Fresenius, or Treadwell-Hall, and particularly to
the admirable discussion of this topic in Morse's !Exercises in
Quantitative Chemistry!.
The statements and rules of procedure which follow are sufficient
for the intelligent use of an analytical balance in connection with
processes prescribed in this introductory manual. It is, however,
imperative that the student should make himself familiar with these
essential features of the balance, and its use. He should fully
realize that the analytical balance is a delicate instrument which
will render excellent service under careful treatment, but such
treatment is an essential condition if its accuracy is to be depended
upon. He should also understand that no set of rules, however
complete, can do away with the necessity for a sense of personal
responsibility, since by carelessness he can render inaccurate not
only his own analyses, but those of all other students using the same
balance.
Before making any weighings the student should seat himself before a
balance and observe the following details of construction:
1. The balance case is mounted on three brass legs, which should
preferably rest in glass cups, backed with rubber to prevent slipping.
The front legs are adjustable as to height and are used to level the
balance case; the rear leg is of permanent length.
2. The front of the case may be raised to give access to the balance.
In some makes doors are provided also at the ends of the balance case.
3. The balance beam is mounted upon an upright in the center of the
case on the top of which is an inlaid agate plate. To the center of
the beam there is attached a steel or agate knife-edge on which the
beam oscillates when it rests on the agate plate.
4. The balance beam, extending to the right and left, is graduated
along its upper edge, usually on both sides, and has at its
extremities two agate or steel knife-edges from which are suspended
stirrups. Each of these stirrups has an agate plate which, when the
balance is in action, rests upon the corresponding knife-edge of the
beam. The balance pans are suspended from the stirrups.
5. A pointer is attached to the center of the beam, and as the beam
oscillates this pointer moves in front of a scale near the base of the
post.
6. At the base of the post, usually in the rear, is a spirit-level.
7. Within the upright is a mechanism, controlled by a knob at the
front of the balance case, which is so arranged as to raise the entire
beam slightly above the level at which the knife-edges are in contact
with the agate plates. When the balance is not in use the beam must
be supported by this device since, otherwise, the constant jarring
to which a balance is inevitably subjected, will soon dull the
knife-edges, and lessen the sensitiveness of the balance.
8. A small weight, or bob, is attached to the pointer (or sometimes
to the beam) by which the center of gravity of the beam and its
attachments may be regulated. The center of gravity must lie very
slightly below the level of the agate plates to secure the desired
sensitiveness of the balance. This is provided for when the balance is
set up and very rarely requires alteration. The student should never
attempt to change this adjustment.
9. Below the balance pans are two pan-arrests operated by a button
from the front of the case. These arrests exert a very slight upward
pressure upon the pans and minimize the displacement of the beam when
objects or weights are being placed upon the pans.
10. A movable rod, operated from one end of the balance case, extends
over the balance beam and carries a small wire weight, called a rider.
By means of this rod the rider can be placed upon any desired division
of the scale on the balance beam. Each numbered division on the beam
corresponds to one milligram, and the use of the rider obviates the
placing of very small fractional weights on the balance pan.
If a new rider is purchased, or an old one replaced, care must be
taken that its weight corresponds to the graduations on the beam of
the balance on which it is to be used. The weight of the rider in
milligrams must be equal to the number of large divisions (5, 6, 10,
or 12) between the central knife-edge and the knife-edge at the end of
the beam. It should be noted that on some balances the last division
bears no number. Each new rider should be tested against a 5 or
10-milligram weight.
In some of the most recent forms of the balance a chain device
replaces the smaller weights and the use of the rider as just
described.
Before using a balance, it is always best to test its adjustment. This
is absolutely necessary if the balance is used by several workers; it
is always a wise precaution under any conditions. For this purpose,
brush off the balance pans with a soft camel's hair brush. Then note
(1) whether the balance is level; (2) that the mechanism for raising
and lowering the beams works smoothly; (3) that the pan-arrests touch
the pans when the beam is lowered; and (4) that the needle swings
equal distances on either side of the zero-point when set in motion
without any load on the pans. If the latter condition is not
fulfilled, the balance should be adjusted by means of the adjusting
screw at the end of the beam unless the variation is not more than one
division on the scale; it is often better to make a proper allowance
for this small zero error than to disturb the balance by an attempt at
correction. Unless a student thoroughly understands the construction
of a balance he should never attempt to make adjustments, but should
apply to the instructor in charge.
The object to be weighed should be placed on the left-hand balance pan
and the weights upon the right-hand pan. Every substance which
could attack the metal of the balance pan should be weighed upon a
watch-glass, and all objects must be dry and cold. A warm body gives
rise to air currents which vitiate the accuracy of the weighing.
The weights should be applied in the order in which they occur in the
weight-box (not at haphazard), beginning with the largest weight which
is apparently required. After a weight has been placed upon the pan
the beam should be lowered upon its knife-edges, and, if necessary,
the pan-arrests depressed. The movement of the pointer will then
indicate whether the weight applied is too great or too small. When
the weight has been ascertained, by the successive addition of small
weights, to the nearest 5 or 10 milligrams, the weighing is completed
by the use of the rider. The correct weight is that which causes the
pointer to swing an equal number of divisions to the right and left
of the zero-point, when the pointer traverses not less than five
divisions on either side.
The balance case should always be closed during the final weighing,
while the rider is being used, to protect the pans from the effect of
air currents.
Before the final determination of an exact weight the beam should
always be lifted from the knife-edges and again lowered into place,
as it frequently happens that the scale pans are, in spite of the
pan-arrests, slightly twisted by the impact of the weights, the beam
being thereby virtually lengthened or shortened. Lifting the beam
restores the proper alignment.
The beam should never be set in motion by lowering it forcibly upon
the knife-edges, nor by touching the pans, but rather by lifting the
rider (unless the balance be provided with some of the newer devices
for the purpose), and the swing should be arrested only when the
needle approaches zero on the scale, otherwise the knife-edges become
dull. For the same reason the beam should never be left upon its
knife-edges, nor should weights be removed from or placed on the
pans without supporting the beam, except in the case of the small
fractional weights.
When the process of weighing has been completed, the weight should
be recorded in the notebook by first noting the vacant spaces in the
weight-box, and then checking the weight by again noting the weights
as they are removed from the pan. This practice will often detect and
avoid errors. It is obvious that the weights should always be returned
to their proper places in the box, and be handled only with pincers.
It should be borne in mind that if the mechanism of a balance is
deranged or if any substance is spilled upon the pans or in the
balance case, the damage should be reported at once. In many instances
serious harm can be averted by prompt action when delay might ruin the
balance.
Samples for analysis are commonly weighed in small tubes with cork
stoppers. Since the stoppers are likely to change in weight from
the varying amounts of moisture absorbed from the atmosphere, it is
necessary to confirm the recorded weight of a tube which has been
unused for some time before weighing out a new portion of substance
from it.
WEIGHTS
The sets of weights commonly used in analytical chemistry range from
20 grams to 5 milligrams. The weights from 20 grams to 1 gram are
usually of brass, lacquered or gold plated. The fractional weights
are of German silver, gold, platinum or aluminium. The rider is of
platinum or aluminium wire.
The sets of weights purchased from reputable dealers are usually
sufficiently accurate for analytical work. It is not necessary that
such a set should be strictly exact in comparison with the absolute
standard of weight, provided they are relatively correct among
themselves, and provided the same set of weights is used in all
weighings made during a given analysis. The analyst should assure
himself that the weights in a set previously unfamiliar to him are
relatively correct by a few simple tests. For example, he should make
sure that in his set two weights of the same denomination (i.e., two
10-gram weights, or the two 100-milligram weights) are actually equal
and interchangeable, or that the 500-milligram weight is equal to
the sum of the 200, 100, 100, 50, 20, 20 and 10-milligram weights
combined, and so on. If discrepancies of more than a few tenths of a
milligram (depending upon the total weight involved) are found, the
weights should be returned for correction. The rider should also be
compared with a 5 or 10-milligram weight.
In an instructional laboratory appreciable errors should be reported
to the instructor in charge for his consideration.
When the highest accuracy is desired, the weights may be calibrated
and corrections applied. A calibration procedure is described in a
paper by T.W. Richards, !J. Am. Chem. Soc.!, 22, 144, and in many
large text-books.
Weights are inevitably subject to corrosion if not properly protected
at all times, and are liable to damage unless handled with great care.
It is obvious that anything which alters the weight of a single piece
in an analytical set will introduce an error in every weighing made
in which that piece is used. This source of error is often extremely
obscure and difficult to detect. The only safeguard against such
errors is to be found in scrupulous care in handling and protection
on the part of the analyst, and an equal insistence that if several
analysts use the same set of weights, each shall realize his
responsibility for the work of others as well as his own.
BURETTES
A burette is made from a glass tube which is as uniformly cylindrical
as possible, and of such a bore that the divisions which are etched
upon its surface shall correspond closely to actual contents.
The tube is contracted at one extremity, and terminates in either a
glass stopcock and delivery-tube, or in such a manner that a piece of
rubber tubing may be firmly attached, connecting a delivery-tube of
glass. The rubber tubing is closed by means of a glass bead. Burettes
of the latter type will be referred to as "plain burettes."
The graduations are usually numbered in cubic centimeters, and the
latter are subdivided into tenths.
One burette of each type is desirable for the analytical procedures
which follow.
PREPARATION OF A BURETTE FOR USE
The inner surface of a burette must be thoroughly cleaned in order
that the liquid as drawn out may drain away completely, without
leaving drops upon the sides. This is best accomplished by treating
the inside of the burette with a warm solution of chromic acid in
concentrated sulphuric acid, applied as follows: If the burette is of
the "plain" type, first remove the rubber tip and force the lower
end of the burette into a medium-sized cork stopper. Nearly fill the
burette with the chromic acid solution, close the upper end with a
cork stopper and tip the burette backward and forward in such a way
as to bring the solution into contact with the entire inner surface.
Remove the stopper and pour the solution into a stock bottle to be
kept for further use, and rinse out the burette with water several
times. Unless the water then runs freely from the burette without
leaving drops adhering to the sides, the process must be repeated
(Note 1).
If the burette has a glass stopcock, this should be removed after
the cleaning and wiped, and also the inside of the ground joint. The
surface of the stopcock should then be smeared with a thin coating of
vaseline and replaced. It should be attached to the burette by means
of a wire, or elastic band, to lessen the danger of breakage.
Fill the burettes with distilled water, and allow the water to run out
through the stopcock or rubber tip until convinced that no air
bubbles are inclosed (Note 2). Fill the burette to a point above the
zero-point and draw off the water until the meniscus is just below
that mark. It is then ready for calibration.
[Note 1: The inner surface of the burette must be absolutely clean if
the liquid is to run off freely. Chromic acid in sulphuric acid is
usually found to be the best cleansing agent, but the mixture must be
warm and concentrated. The solution can be prepared by pouring over a
few crystals of potassium bichromate a little water and then adding
concentrated sulphuric acid.]
[Note 2: It is always necessary to insure the absence of air bubbles
in the tips or stopcocks. The treatment described above will usually
accomplish this, but, in the case of plain burettes it is sometimes
better to allow a little of the liquid to flow out of the tip while it
is bent upwards. Any air which may be entrapped then rises with the
liquid and escapes.
If air bubbles escape during subsequent calibration or titration, an
error is introduced which vitiates the results.]
READING OF A BURETTE
All liquids when placed in a burette form what is called a meniscus at
their upper surfaces. In the case of liquids such as water or
aqueous solutions this meniscus is concave, and when the liquids are
transparent accurate readings are best obtained by observing the
position on the graduated scales of the lowest point of the meniscus.
This can best be done as follows: Wrap around the burette a piece of
colored paper, the straight, smooth edges of which are held evenly
together with the colored side next to the burette (Note 1). Hold the
paper about two small divisions below the meniscus and raise or lower
the level of the eyes until the edge of the paper at the back of the
burette is just hidden from the eye by that in front (Note 2). Note
the position of the lowest point of the curve of the meniscus,
estimating the tenths of the small divisions, thus reading its
position to hundredths of a cubic centimeter.
[Note 1: The ends of the colored paper used as an aid to accurate
readings may be fastened together by means of a gummed label. The
paper may then remain on the burette and be ready for immediate use by
sliding it up or down, as required.]
[Note 2: To obtain an accurate reading the eye must be very nearly on
a level with the meniscus. This is secured by the use of the paper
as described. The student should observe by trial how a reading is
affected when the meniscus is viewed from above or below.
The eye soon becomes accustomed to estimating the tenths of the
divisions. If the paper is held as directed, two divisions below the
meniscus, one whole division is visible to correct the judgment. It is
not well to attempt to bring the meniscus exactly to a division mark
on the burette. Such readings are usually less accurate than those in
which the tenths of a division are estimated.]
CALIBRATION OF GLASS MEASURING DEVICES
If accuracy of results is to be attained, the correctness of all
measuring instruments must be tested. None of the apparatus offered
for sale can be implicitly relied upon except those more expensive
instruments which are accompanied by a certificate from the !National
Bureau of Standards! at Washington, or other equally authentic source.
The bore of burettes is subject to accidental variations, and since
the graduations are applied by machine without regard to such
variations of bore, local errors result.
The process of testing these instruments is called !calibration!.
It is usually accomplished by comparing the actual weight of water
contained in the instrument with its apparent volume.
There is, unfortunately, no uniform standard of volume which has been
adopted for general use in all laboratories. It has been variously
proposed to consider the volume of 1000 grams of water at 4°, 15.5°,
16°, 17.5°, and even 20°C., as a liter for practical purposes, and to
consider the cubic centimeter to be one one-thousandth of that volume.
The true liter is the volume of 1000 grams of water at 4°C.; but
this is obviously a lower temperature than that commonly found in
laboratories, and involves the constant use of corrections if taken as
a laboratory standard. Many laboratories use 15.5°C. (60° F.) as the
working standard. It is plain that any temperature which is deemed
most convenient might be chosen for a particular laboratory, but it
cannot be too strongly emphasized that all measuring instruments,
including burettes, pipettes, and flasks, should be calibrated at that
temperature in order that the contents of each burette, pipette,
etc., shall be comparable with that of every other instrument, thus
permitting general interchange and substitution. For example, it is
obvious that if it is desired to remove exactly 50 cc. from a solution
which has been diluted to 500 cc. in a graduated flask, the 50 cc.
flask or pipette used to remove the fractional portion must give
a correct reading at the same temperature as the 500 cc. flask.
Similarly, a burette used for the titration of the 50 cc. of solution
removed should be calibrated under the same conditions as the
measuring flasks or pipettes employed with it.
The student should also keep constantly in mind the fact that all
volumetric operations, to be exact, should be carried out as nearly at
a constant temperature as is practicable. The spot selected for
such work should therefore be subject to a minimum of temperature
variations, and should have as nearly the average temperature of
the laboratory as is possible. In all work, whether of calibration,
standardization, or analysis, the temperature of the liquids employed
must be taken into account, and if the temperature of these liquids
varies more than 3° or 4° from the standard temperature chosen for the
laboratory, corrections must be applied for errors due to expansion or
contraction, since volumes of a liquid measured at different times are
comparable only under like conditions as to temperature. Data to be
used for this purpose are given in the Appendix. Neglect of this
correction is frequently an avoidable source of error and annoyance in
otherwise excellent work. The temperature of all solutions at the time
of standardization should be recorded to facilitate the application of
temperature corrections, if such are necessary at any later time.
CALIBRATION OF THE BURETTES
Two burettes, one at least of which should have a glass stopper, are
required throughout the volumetric work. Both burettes should be
calibrated by the student to whom they are assigned.
PROCEDURE.--Weigh a 50 cc., flat-bottomed flask (preferably a
light-weight flask), which must be dry on the outside, to the nearest
centigram. Record the weight in the notebook. (See Appendix for
suggestions as to records.) Place the flask under the burette and draw
out into it about 10 cc. of water, removing any drop on the tip by
touching it against the inside of the neck of the flask. Do not
attempt to stop exactly at the 10 cc. mark, but do not vary more than
0.1 cc. from it. Note the time, and at the expiration of three minutes
(or longer) read the burette accurately, and record the reading in the
notebook (Note 1). Meanwhile weigh the flask and water to centigrams
and record its weight (Note 2). Draw off the liquid from 10 cc. to
about 20 cc. into the same flask without emptying it; weigh, and at
the expiration of three minutes take the reading, and so on throughout
the length of the burette. When it is completed, refill the burette
and check the first calibration.
The differences in readings represent the apparent volumes, the
differences in weights the true volumes. For example, if an apparent
volume of 10.05 cc. is found to weigh 10.03 grams, it may be assumed
with sufficient accuracy that the error in that 10 cc. amounts to
-0.02 cc., or -0.002 for each cubic centimeter (Note 3).
In the calculation of corrections the temperature of the water must be
taken into account, if this varies more than 4°C. from the laboratory
standard temperature, consulting the table of densities of water in
the Appendix.
From the final data, plot the corrections to be applied so that they
may be easily read for each cubic centimeter throughout the burette.
The total correction at each 10 cc. may also be written on the burette
with a diamond, or etching ink, for permanence of record.
[Note 1: A small quantity of liquid at first adheres to the side of
even a clean burette. This slowly unites with the main body of liquid,
but requires an appreciable time. Three minutes is a sufficient
interval, but not too long, and should be adopted in every instance
throughout the whole volumetric practice before final readings are
recorded.]
[Note 2: A comparatively rough balance, capable of weighing to
centigrams, is sufficiently accurate for use in calibrations, for a
moment's reflection will show that it would be useless to weigh the
water with an accuracy greater than that of the readings taken on
the burette. The latter cannot exceed 0.01 cc. in accuracy, which
corresponds to 0.01 gram.
The student should clearly understand that !all other weighings!,
except those for calibration, should be made accurately to 0.0001
gram, unless special directions are given to the contrary.
Corrections for temperature variations of less than 4°C. are
negligible, as they amount to less than 0.01 gram for each 10 grams of
water withdrawn.]
[Note 3: Should the error discovered in any interval of 10 cc. on the
burette exceed 0.10 cc., it is advisable to weigh small portions (even
1 cc.) to locate the position of the variation of bore in the
tube rather than to distribute the correction uniformly over the
corresponding 10 cc. The latter is the usual course for small
corrections, and it is convenient to calculate the correction
corresponding to each cubic centimeter and to record it in the form
of a table or calibration card, or to plot a curve representing the
values.
Burettes may also be calibrated by drawing off the liquid in
successive portions through a 5 cc. pipette which has been accurately
calibrated, as a substitute for weighing. If many burettes are to be
tested, this is a more rapid method.]
PIPETTES
A !pipette! may consist of a narrow tube, in the middle of which is
blown a bulb of a capacity a little less than that which it is desired
to measure by the pipette; or it may be a miniature burette, without
the stopcock or rubber tip at the lower extremity. In either case, the
flow of liquid is regulated by the pressure of the finger on the top,
which governs the admission of the air.
Pipettes are usually already graduated when purchased, but they
require calibration for accurate work.
CALIBRATION OF PIPETTES
PROCEDURE.--Clean the pipette. Draw distilled water into it by sucking
at the upper end until the water is well above the graduation mark.
Quickly place the forefinger over the top of the tube, thus preventing
the entrance of air and holding the water in the pipette. Cautiously
admit a little air by releasing the pressure of the finger, and allow
the level of the water to fall until the lowest point of the meniscus
is level with the graduation. Hold the water at that point by pressure
of the finger and then allow the water to run out from the pipette
into a small tared, or weighed, beaker or flask. After a definite time
interval, usually two to three minutes, touch the end of the pipette
against the side of the beaker or flask to remove any liquid adhering
to it (Note 1). The increase in weight of the flask in grams
represents the volume of the water in cubic centimeters delivered by
the pipette. Calculate the necessary correction.
[Note 1: A definite interval must be allowed for draining, and a
definite practice adopted with respect to the removal of the liquid
which collects at the end of the tube, if the pipette is designed to
deliver a specific volume when emptied. This liquid may be removed
at the end of a definite interval either by touching the side of the
vessel or by gently blowing out the last drops. Either practice, when
adopted, must be uniformly adhered to.]
FLASKS
!Graduated or measuring flasks! are similar to the ordinary
flat-bottomed flasks, but are provided with long, narrow necks in
order that slight variations in the position of the meniscus with
respect to the graduation shall represent a minimum volume of liquid.
The flasks must be of such a capacity that, when filled with the
specified volume, the liquid rises well into the neck.
GRADUATION OF FLASKS
It is a general custom to purchase the flasks ungraduated and to
graduate them for use under standard conditions selected for the
laboratory in question. They may be graduated for "contents" or
"delivery." When graduated for "contents" they contain a specified
volume when filled to the graduation at a specified temperature, and
require to be washed out in order to remove all of the solution from
the flask. Flasks graduated for "delivery" will deliver the specified
volume of a liquid without rinsing. A flask may, of course, be
graduated for both contents and delivery by placing two graduation
marks upon it.
PROCEDURE.--To calibrate a flask for !contents!, proceed as follows:
Clean the flask, using a chromic acid solution, and dry it carefully
outside and inside. Tare it accurately; pour water into the flask
until the weight of the latter counterbalances weights on the opposite
pan which equal in grams the number of cubic centimeters of water
which the flask is to contain. Remove any excess of water with the aid
of filter paper (Note 1). Take the flask from the balance, stopper
it, place it in a bath at the desired temperature, usually 15.5°
or 17.5°C., and after an hour mark on the neck with a diamond the
location of the lowest point of the meniscus (Note 2). The mark may
be etched upon the flask by hydrofluoric acid, or by the use of an
etching ink now commonly sold on the market.
To graduate a flask which is designed to !deliver! a specified volume,
proceed as follows: Clean the flask as usual and wipe all moisture
from the outside. Fill it with distilled water. Pour out the water
and allow the water to drain from the flask for three minutes.
Counterbalance the flask with weights to the nearest centigram.
Add weights corresponding in grams to the volume desired, and add
distilled water to counterbalance these weights. An excess of water,
or water adhering to the neck of the flask, may be removed by means of
a strip of clean filter paper. Stopper the flask, place it in a bath
at 15.5°C. or 17.5°C. and, after an hour, mark the location of the
lowest point of the meniscus, as described above.
[Note 1: The allowable error in counterbalancing the water and
weights varies with the volume of the flask. It should not exceed one
ten-thousandth of the weight of water.]
[Note 2: Other methods are employed which involve the use of
calibrated apparatus from which the desired volume of water may be run
into the dry flask and the position of the meniscus marked directly
upon it. For a description of a procedure which is most convenient
when many flasks are to be calibrated, the student is referred to the
!Am. Chem J.!, 16, 479.]
GENERAL DIRECTIONS FOR VOLUMETRIC ANALYSES
It cannot be too strongly emphasized that for the success of analyses
uniformity of practice must prevail throughout all volumetric work
with respect to those factors which can influence the accuracy of the
measurement of liquids. For example, whatever conditions are imposed
during the calibration of a burette, pipette, or flask (notably the
time allowed for draining), must also prevail whenever the flask or
burette is used.
The student should also be constantly watchful to insure parallel
conditions during both standardization and analyst with respect to the
final volume of liquid in which a titration takes place. The value
of a standard solution is only accurate under the conditions which
prevailed when it was standardized. It is plain that the standard
solutions must be scrupulously protected from concentration or
dilution, after their value has been established. Accordingly, great
care must be taken to thoroughly rinse out all burettes, flasks, etc.,
with the solutions which they are to contain, in order to remove all
traces of water or other liquid which could act as a diluent. It is
best to wash out a burette at least three times with small portions of
a solution, allowing each to run out through the tip before assuming
that the burette is in a condition to be filled and used. It is, of
course, possible to dry measuring instruments in a hot closet, but
this is tedious and unnecessary.
To the same end, all solutions should be kept stoppered and away from
direct sunlight or heat. The bottles should be shaken before use to
collect any liquid which may have distilled from the solution and
condensed on the sides.
The student is again reminded that variations in temperature of
volumetric solutions must be carefully noted, and care should always
be taken that no source of heat is sufficiently near the solutions to
raise the temperature during use.
Much time may be saved by estimating the approximate volume of a
standard solution which will be required for a titration (if the data
are obtainable) before beginning the operation. It is then possible to
run in rapidly approximately the required amount, after which it is
only necessary to determine the end-point slowly and with accuracy.
In such cases, however, the knowledge of the approximate amount to be
required should never be allowed to influence the judgment regarding
the actual end-point.
STANDARD SOLUTIONS
The strength or value of a solution for a specific reaction is
determined by a procedure called !Standardization!, in which the
solution is brought into reaction with a definite weight of a
substance of known purity. For example, a definite weight of pure
sodium carbonate may be dissolved in water, and the volume of a
solution of hydrochloric acid necessary to exactly neutralize the
carbonate accurately determined. From these data the strength or value
of the acid is known. It is then a !standard solution!.
NORMAL SOLUTIONS
Standard solutions may be made of a purely empirical strength dictated
solely by convenience of manipulation, or the concentration may
be chosen with reference to a system which is applicable to all
solutions, and based upon chemical equivalents. Such solutions are
called !Normal Solutions! and contain such an amount of the reacting
substance per liter as is equivalent in its chemical action to one
gram of hydrogen, or eight grams of oxygen. Solutions containing one
half, one tenth, or one one-hundredth of this quantity per liter are
called, respectively, half-normal, tenth-normal, or hundredth-normal
solutions.
Since normal solutions of various reagents are all referred to a
common standard, they have an advantage not possessed by empirical
solutions, namely, that they are exactly equivalent to each other.
Thus, a liter of a normal solution of an acid will exactly neutralize
a liter of a normal alkali solution, and a liter of a normal oxidizing
solution will exactly react with a liter of a normal reducing
solution, and so on.
Beside the advantage of uniformity, the use of normal solutions
simplifies the calculations of the results of analyses. This is
particularly true if, in connection with the normal solution, the
weight of substance for analysis is chosen with reference to the
atomic or molecular weight of the constituent to be determined. (See
problem 26.)
The preparation of an !exactly! normal, half-normal, or tenth-normal
solution requires considerable time and care. It is usually carried
out only when a large number of analyses are to be made, or when the
analyst has some other specific purpose in view. It is, however, a
comparatively easy matter to prepare standard solutions which differ
but slightly from the normal or half-normal solution, and these have
the advantage of practical equality; that is, two approximately
half-normal solutions are more convenient to work with than two which
are widely different in strength. It is, however, true that some of
the advantage which pertains to the use of normal solutions as regards
simplicity of calculations is lost when using these approximate
solutions.
The application of these general statements will be made clear in
connection with the use of normal solutions in the various types of
volumetric processes which follow.
I. NEUTRALIZATION METHODS
ALKALIMETRY AND ACIDIMETRY
GENERAL DISCUSSION
!Standard Acid Solutions! may be prepared from either hydrochloric,
sulphuric, or oxalic acid. Hydrochloric acid has the advantage of
forming soluble compounds with the alkaline earths, but its solutions
cannot be boiled without danger of loss of strength; sulphuric acid
solutions may be boiled without loss, but the acid forms insoluble
sulphates with three of the alkaline earths; oxalic acid can be
accurately weighed for the preparation of solutions, and its solutions
may be boiled without loss, but it forms insoluble oxalates with
three of the alkaline earths and cannot be used with certain of the
indicators.
!Standard Alkali Solutions! may be prepared from sodium or potassium
hydroxide, sodium carbonate, barium hydroxide, or ammonia. Of sodium
and potassium hydroxide, it may be said that they can be used with all
indicators, and their solutions may be boiled, but they absorb carbon
dioxide readily and attack the glass of bottles, thereby losing
strength; sodium carbonate may be weighed directly if its purity is
assured, but the presence of carbonic acid from the carbonate is a
disadvantage with many indicators; barium hydroxide solutions may
be prepared which are entirely free from carbon dioxide, and such
solutions immediately show by precipitation any contamination from
absorption, but the hydroxide is not freely soluble in water; ammonia
does not absorb carbon dioxide as readily as the caustic alkalies,
but its solutions cannot be boiled nor can they be used with all
indicators. The choice of a solution must depend upon the nature of
the work in hand.
A !normal acid solution! should contain in one liter that quantity of
the reagent which represents 1 gram of hydrogen replaceable by a base.
For example, the normal solution of hydrochloric acid (HCl) should
contain 36.46 grams of gaseous hydrogen chloride, since that amount
furnishes the requisite 1 gram of replaceable hydrogen. On the other
hand, the normal solution of sulphuric acid (H_{2}SO_{4}) should
contain only 49.03 grams, i.e., one half of its molecular weight in
grams.
A !normal alkali solution! should contain sufficient alkali in a liter
to replace 1 gram of hydrogen in an acid. This quantity is represented
by the molecular weight in grams (40.01) of sodium hydroxide (NaOH),
while a sodium carbonate solution (Na_{2}CO_{3}) should contain but
one half the molecular weight in grams (i.e., 53.0 grams) in a liter
of normal solution.
Half-normal or tenth-normal solutions are employed in most analyses
(except in the case of the less soluble barium hydroxide). Solutions
of the latter strength yield more accurate results when small
percentages of acid or alkali are to be determined.
INDICATORS
It has already been pointed out that the purpose of an indicator is to
mark (usually by a change of color) the point at which just enough of
the titrating solution has been added to complete the chemical change
which it is intended to bring about. In the neutralization processes
which are employed in the measurement of alkalies (!alkalimetry!)
or acids (!acidimetry!) the end-point of the reaction should, in
principle, be that of complete neutrality. Expressed in terms of ionic
reactions, it should be the point at which the H^{+} ions from an
acid[Note 1] unite with a corresponding number of OH^{-} ions from a
base to form water molecules, as in the equation
H^{+}, Cl^{-} + Na^{+}, OH^{-} --> Na^{+}, Cl^{-} + (H_{2}O).
It is not usually possible to realize this condition of exact
neutrality, but it is possible to approach it with sufficient
exactness for analytical purposes, since substances are known which,
in solution, undergo a sharp change of color as soon as even a minute
excess of H^{+} or OH^{-} ions are present. Some, as will be seen,
react sharply in the presence of H^{+} ions, and others with OH^{-}
ions. These substances employed as indicators are usually organic
compounds of complex structure and are closely allied to the dyestuffs
in character.
[Note 1: A knowledge on the part of the student of the ionic theory
as applied to aqueous solutions of electrolytes is assumed. A brief
outline of the more important applications of the theory is given in
the Appendix.]
BEHAVIOR OF ORGANIC INDICATORS
The indicators in most common use for acid and alkali titrations are
methyl orange, litmus, and phenolphthalein.
In the following discussion of the principles underlying the behavior
of the indicators as a class, methyl orange and phenolphthalein will
be taken as types. It has just been pointed out that indicators are
bodies of complicated structure. In the case of the two indicators
named, the changes which they undergo have been carefully studied by
Stieglitz (!J. Am. Chem. Soc.!, 25, 1112) and others, and it appears
that the changes involved are of two sorts: First, a rearrangement
of the atoms within the molecule, such as often occurs in organic
compounds; and, second, ionic changes. The intermolecular changes
cannot appropriately be discussed here, as they involve a somewhat
detailed knowledge of the classification and general behavior of
organic compounds; they will, therefore, be merely alluded to, and
only the ionic changes followed.
Methyl orange is a representative of the group of indicators which,
in aqueous solutions, behave as weak bases. The yellow color which it
imparts to solutions is ascribed to the presence of the undissociated
base. If an acid, such as HCl, is added to such a solution, the acid
reacts with the indicator (neutralizes it) and a salt is formed, as
indicated by the equation:
(M.o.)^{+}, OH^{-} + H^{+}, Cl^{-} --> (M.o.)^{+} Cl^{-} + (H_{2}O).
This salt ionizes into (M.o.)^{+} (using this abbreviation for the
positive complex) and Cl^{-}; but simultaneously with this ionization
there appears to be an internal rearrangement of the atoms which
results in the production of a cation which may be designated as
(M'.o'.)^{+}, and it is this which imparts a characteristic red color
to the solution. As these changes occur in the presence of even a
very small excess of acid (that is, of H^{+} ions), it serves as the
desired index of their presence in the solution. If, now, an alkali,
such as NaOH, is added to this reddened solution, the reverse
series of changes takes place. As soon as the free acid present is
neutralized, the slightest excess of sodium hydroxide, acting as
a strong base, sets free the weak, little-dissociated base of the
indicator, and at the moment of its formation it reverts, because of
the rearrangement of the atoms, to the yellow form:
OH^{-} + (M'.o'.)^{+} --> [M'.o'.OH] --> [M.o.OH].
Phenolphthalein, on the other hand, is a very weak, little-dissociated
acid, which is colorless in neutral aqueous solution or in the
presence of free H^{+} ions. When an alkali is added to such a
solution, even in slight excess, the anion of the salt which has
formed from the acid of the indicator undergoes a rearrangement of the
atoms, and a new ion, (Ph')^{+}, is formed, which imparts a pink color
to the solution:
H^{+}, (Ph)^{-} + Na^{+}, OH^{-} --> (H_{2}O) + Na^{+}, (Ph)^{-}
--> Na^{+}, (Ph')^{-}
The addition of the slightest excess of an acid to this solution, on
the other hand, occasions first the reversion to the colorless ion and
then the setting free of the undissociated acid of the indicator:
H^{+}, (Ph')^{-} --> H^{+}, (Ph)^{-} --> (HPh).
Of the common indicators methyl orange is the most sensitive toward
alkalies and phenolphthalein toward acids; the others occupy
intermediate positions. That methyl orange should be most sensitive
toward alkalies is evident from the following considerations: Methyl
orange is a weak base and, therefore, but little dissociated. It
should, then, be formed in the undissociated condition as soon as even
a slight excess of OH^{-} ions is present in the solution, and there
should be a prompt change from red to yellow as outlined above. On the
other hand, it should be an unsatisfactory indicator for use with weak
acids (acetic acid, for example) because the salts which it forms
with such acids are, like all salts of that type, hydrolyzed to a
considerable extent. This hydrolytic change is illustrated by the
equation:
(M.o.)^{+} C_{2}H_{3}O_{2}^{-} + H^{+}, OH^{-} --> [M.o.OH] + H^{+},
C_{2}H_{3}O_{2}^{-}.
Comparison of this equation with that on page 30 will make it plain
that hydrolysis is just the reverse of neutralization and must,
accordingly, interfere with it. Salts of methyl orange with weak acids
are so far hydrolyzed that the end-point is uncertain, and methyl
orange cannot be used in the titration of such acids, while with
the very weak acids, such as carbonic acid or hydrogen sulphide
(hydrosulphuric acid), the salts formed with methyl orange are, in
effect, completely hydrolyzed (i.e., no neutralization occurs), and
methyl orange is accordingly scarcely affected by these acids. This
explains its usefulness, as referred to later, for the titration of
strong acids, such as hydrochloric acid, even in the presence of
carbonates or sulphides in solution.
Phenolphthalein, on the other hand, should be, as it is, the best of
the common indicators for use with weak acids. For, since it is
itself a weak acid, it is very little dissociated, and its nearly
undissociated, colorless molecules are promptly formed as soon as
there is any free acid (that is, free H^{+} ions) in the solution.
This indicator cannot, however, be successfully used with weak bases,
even ammonium hydroxide; for, since it is weak acid, the salts
which it forms with weak alkalies are easily hydrolyzed, and as a
consequence of this hydrolysis the change of color is not sharp.
This indicator can, however, be successfully used with strong bases,
because the salts which it forms with such bases are much less
hydrolyzed and because the excess of OH^{-} ions from these bases also
diminishes the hydrolytic action of water.
This indicator is affected by even so weak an acid as carbonic acid,
which must be removed by boiling the solution before titration. It is
the indicator most generally employed for the titration of organic
acids.
In general, it may be stated that when a strong acid, such as
hydrochloric, sulphuric or nitric acid, is titrated against a strong
base, such as sodium hydroxide, potassium hydroxide, or barium
hydroxide, any of these indicators may be used, since very little
hydrolysis ensues. It has been noted above that the color change does
not occur exactly at theoretical neutrality, from which it follows
that no two indicators will show exactly the same end-point when acids
and alkalis are brought together. It is plain, therefore, that the
same indicator must be employed for both standardization and analysis,
and that, if this is done, accurate results are obtainable.
The following table (Note 1) illustrates the variations in the volume
of an alkali solution (tenth-normal sodium hydroxide) required to
produce an alkaline end-point when run into 10 cc. of tenth-normal
sulphuric acid, diluted with 50 cc. of water, using five drops of each
of the different indicator solutions.
====================================================================
| | | |
INDICATOR | N/10 | N/10 |COLOR IN ACID|COLOR IN ALKA-
| H_{2}SO_{4}| NaOH |SOLUTION |LINE SOLUTION
_______________|____________|__________|_____________|______________
| cc. | cc. | cc. |
Methyl orange | 10 | 9.90 | Red | Yellow
Lacmoid | 10 | 10.00 | Red | Blue
Litmus | 10 | 10.00 | Red | Blue
Rosalic acid | 10 | 10.07 | Yellow | Pink
Phenolphthalein| 10 | 10.10 | Colorless | Pink
====================================================================
It should also be stated that there are occasionally secondary
changes, other than those outlined above, which depend upon the
temperature and concentration of the solutions in which the indicators
are used. These changes may influence the sensitiveness of an
indicator. It is important, therefore, to take pains to use
approximately the same volume of solution when standardizing that is
likely to be employed in analysis; and when it is necessary, as is
often the case, to titrate the solution at boiling temperature, the
standardization should take place under the same conditions. It is
also obvious that since some acid or alkali is required to react with
the indicator itself, the amount of indicator used should be uniform
and not excessive. Usually a few drops of solution will suffice.
The foregoing statements with respect to the behavior of indicators
present the subject in its simplest terms. Many substances other than
those named may be employed, and they have been carefully studied to
determine the exact concentration of H^{+} ions at which the color
change of each occurs. It is thus possible to select an indicator
for a particular purpose with considerable accuracy. As data of this
nature do not belong in an introductory manual, reference is made to
the following papers or books in which a more extended treatment of
the subject may be found:
Washburn, E.W., Principles of Physical Chemistry (McGraw-Hill Book
Co.), (Second Edition, 1921), pp. 380-387.
Prideaux, E.B.R., The Theory and Use of Indicators (Constable & Co.,
Ltd.), (1917).
Salm, E., A Study of Indicators, !Z. physik. Chem.!, 57 (1906),
471-501.
Stieglitz, J., Theories of Indicators, !J. Am. Chem. Soc.!, 25 (1903),
1112-1127.
Noyes, A.A., Quantitative Applications of the Theory of Indicators to
Volumetric Analysis, !J. Am. Chem. Soc.!, 32 (1911), 815-861.
Bjerrum, N., General Discussion, !Z. Anal. Chem.!, 66 (1917), 13-28
and 81-95.
Ostwald, W., Colloid Chemistry of Indicators, !Z. Chem. Ind.
Kolloide!, 10 (1912), 132-146.
[Note 1: Glaser, !Indikatoren der Acidimetrie und Alkalimetrie!.
Wiesbaden, 1901.]
PREPARATION OF INDICATOR SOLUTIONS
A !methyl orange solution! for use as an indicator is commonly made by
dissolving 0.05-0.1 gram of the compound (also known as Orange III) in
a few cubic centimeters of alcohol and diluting with water to 100 cc.
A good grade of material should be secured. It can be successfully
used for the titration of hydrochloric, nitric, sulphuric, phosphoric,
and sulphurous acids, and is particularly useful in the determination
of bases, such as sodium, potassium, barium, calcium, and ammonium
hydroxides, and even many of the weak organic bases. It can also be
used for the determination, by titration with a standard solution of
a strong acid, of the salts of very weak acids, such as carbonates,
sulphides, arsenites, borates, and silicates, because the weak acids
which are liberated do not affect the indicator, and the reddening of
the solution does not take place until an excess of the strong acid
is added. It should be used in cold, not too dilute, solutions. Its
sensitiveness is lessened in the presence of considerable quantities
of the salts of the alkalies.
A !phenolphthalein solution! is prepared by dissolving 1 gram of the
pure compound in 100 cc. of 95 per cent alcohol. This indicator is
particularly valuable in the determination of weak acids, especially
organic acids. It cannot be used with weak bases, even ammonia. It
is affected by carbonic acid, which must, therefore, be removed by
boiling when other acids are to be measured. It can be used in hot
solutions. Some care is necessary to keep the volume of the solutions
to be titrated approximately uniform in standardization and in
analysis, and this volume should not in general exceed 125-150 cc. for
the best results, since the compounds formed by the indicator undergo
changes in very dilute solution which lessen its sensitiveness.
The preparation of a !solution of litmus! which is suitable for use
as an indicator involves the separation from the commercial litmus of
azolithmine, the true coloring principle. Soluble litmus tablets are
often obtainable, but the litmus as commonly supplied to the market is
mixed with calcium carbonate or sulphate and compressed into lumps. To
prepare a solution, these are powdered and treated two or three times
with alcohol, which dissolves out certain constituents which cause a
troublesome intermediate color if not removed. The alcohol is decanted
and drained off, after which the litmus is extracted with hot water
until exhausted. The solution is allowed to settle for some time, the
clear liquid siphoned off, concentrated to one-third its volume and
acetic acid added in slight excess. It is then concentrated to a
sirup, and a large excess of 95 per cent. alcohol added to it. This
precipitates the blue coloring matter, which is filtered off, washed
with alcohol, and finally dissolved in a small volume of water and
diluted until about three drops of the solution added to 50 cc. of
water just produce a distinct color. This solution must be kept in an
unstoppered bottle. It should be protected from dust by a loose plug
of absorbent cotton. If kept in a closed bottle it soon undergoes a
reduction and loses its color, which, however, is often restored by
exposure to the air.
Litmus can be employed successfully with the strong acids and bases,
and also with ammonium hydroxide, although the salts of the latter
influence the indicator unfavorably if present in considerable
concentration. It may be employed with some of the stronger organic
acids, but the use of phenolphthalein is to be preferred.
PREPARATION OF STANDARD SOLUTIONS
!Hydrochloric Acid and Sodium Hydroxide. Approximate Strength!, 0.5 N
PROCEDURE.--Measure out 40 cc. of concentrated, pure hydrochloric
acid into a clean liter bottle, and dilute with distilled water to an
approximate volume of 1000 cc. Shake the solution vigorously for a
full minute to insure uniformity. Be sure that the bottle is not too
full to permit of a thorough mixing, since lack of care at this point
will be the cause of much wasted time (Note 1).
Weigh out, upon a rough balance, 23 grams of sodium hydroxide (Note
2). Dissolve the hydroxide in water in a beaker. Pour the solution
into a liter bottle and dilute, as above, to approximately 1000 cc.
This bottle should preferably have a rubber stopper, as the hydroxide
solution attacks the glass of the ground joint of a glass stopper, and
may cement the stopper to the bottle. Shake the solution as described
above.
[Note 1: The original solutions are prepared of a strength greater
than 0.5 N, as they are more readily diluted than strengthened if
later adjustment is desired.
Too much care cannot be taken to insure perfect uniformity of
solutions before standardization, and thoroughness in this respect
will, as stated, often avoid much waste of time. A solution once
thoroughly mixed remains uniform.]
[Note 2: Commercial sodium hydroxide is usually impure and always
contains more or less carbonate; an allowance is therefore made for
this impurity by placing the weight taken at 23 grams per liter. If
the hydroxide is known to be pure, a lesser amount (say 21 grams) will
suffice.]
COMPARISON OF ACID AND ALKALI SOLUTIONS
PROCEDURE.--Rinse a previously calibrated burette three times with the
hydrochloric acid solution, using 10 cc. each time, and allowing the
liquid to run out through the tip to displace all water and air
from that part of the burette. Then fill the burette with the acid
solution. Carry out the same procedure with a second burette, using
the sodium hydroxide solution.
The acid solution may be placed in a plain or in a glass-stoppered
burette as may be more convenient, but the alkaline solution should
never be allowed to remain long in a glass-stoppered burette, as it
tends to cement the stopper to the burette, rendering it useless. It
is preferable to use a plain burette for this solution.
When the burettes are ready for use and all air bubbles displaced from
the tip (see Note 2, page 17) note the exact position of the liquid in
each, and record the readings in the notebook. (Consult page 188.) Run
out from the burette into a beaker about 40 cc. of the acid and add
two drops of a solution of methyl orange; dilute the acid to about
80 cc. and run out alkali solution from the other burette, stirring
constantly, until the pink has given place to a yellow. Wash down the
sides of the beaker with a little distilled water if the solution has
spattered upon them, return the beaker to the acid burette, and add
acid to restore the pink; continue these alternations until the point
is accurately fixed at which a single drop of either solutions served
to produce a distinct change of color. Select as the final end-point
the appearance of the faintest pink tinge which can be recognized, or
the disappearance of this tinge, leaving a pure yellow; but always
titrate to the same point (Note 1). If the titration has occupied more
than the three minutes required for draining the sides of the burette,
the final reading may be taken immediately and recorded in the
notebook.
Refill the burettes and repeat the titration. From the records of
calibration already obtained, correct the burette readings and make
corrections for temperature, if necessary. Obtain the ratio of the
sodium hydroxide solution to that of hydrochloric acid by dividing
the number of cubic centimeters of acid used by the number of cubic
centimeters of alkali required for neutralization. The check results
of the two titrations should not vary by more than two parts in one
thousand (Note 2). If the variation in results is greater than this,
refill the burettes and repeat the titration until satisfactory values
are obtained. Use a new page in the notebook for each titration.
Inaccurate values should not be erased or discarded. They should be
retained and marked "correct" or "incorrect," as indicated by the
final outcome of the titrations. This custom should be rigidly
followed in all analytical work.
[Note 1: The end-point should be chosen exactly at the point of
change; any darker tint is unsatisfactory, since it is impossible to
carry shades of color in the memory and to duplicate them from day to
day.]
[Note 2: While variation of two parts in one thousand in the values
obtained by an inexperienced analyst is not excessive, the idea must
be carefully avoided that this is a standard for accurate work to be
!generally applied!. In many cases, after experience is gained, the
allowable error is less than this proportion. In a few cases a
larger variation is permissible, but these are rare and can only
be recognized by an experienced analyst. It is essential that the
beginner should acquire at least the degree of accuracy indicated if
he is to become a successful analyst.]
STANDARDIZATION OF HYDROCHLORIC ACID
SELECTION AND PREPARATION OF STANDARD
The selection of the best substance to be used as a standard for acid
solutions has been the subject of much controversy. The work of Lunge
(!Ztschr. angew. Chem.! (1904), 8, 231), Ferguson (!J. Soc. Chem.
Ind.! (1905), 24, 784), and others, seems to indicate that the best
standard is sodium carbonate prepared from sodium bicarbonate by
heating the latter at temperature between 270° and 300°C. The
bicarbonate is easily prepared in a pure state, and at the
temperatures named the decomposition takes place according to the
equation
2HNaCO_{3} --> Na_{2}CO_{3} + H_{2}O + CO_{2}
and without loss of any carbon dioxide from the sodium carbonate, such
as may occur at higher temperatures. The process is carried out as
described below.
PROCEDURE.--Place in a porcelain crucible about 6 grams (roughly
weighed) of the purest sodium bicarbonate obtainable. Rest the
crucible upon a triangle of iron or copper wire so placed within a
large crucible that there is an open air space of about three eighths
of an inch between them. The larger crucible may be of iron, nickel or
porcelain, as may be most convenient. Insert the bulb of a thermometer
reading to 350°C. in the bicarbonate, supporting it with a clamp so
that the bulb does not rest on the bottom of the crucible. Heat
the outside crucible, using a rather small flame, and raise the
temperature of the bicarbonate fairly rapidly to 270°C. Then regulate
the heat in such a way that the temperature rises !slowly! to 300°C.
in the course of a half-hour. The bicarbonate should be frequently
stirred with a clean, dry, glass rod, and after stirring, should be
heaped up around the bulb of the thermometer in such a way as to cover
it. This will require attention during most of the heating, as the
temperature should not be permitted to rise above 310°C. for any
length of time. At the end of the half-hour remove the thermometer and
transfer the porcelain crucible, which now contains sodium carbonate,
to a desiccator. When it is cold, transfer the carbonate to a
stoppered weighing tube or weighing-bottle.
STANDARDIZATION
PROCEDURE.--Clean carefully the outside of a weighing-tube, or
weighing-bottle, containing the pure sodium carbonate, taking care
to handle it as little as possible after wiping. Weigh the tube
accurately to 0.0001 gram, and record the weight in the notebook. Hold
the tube over the top of a beaker (200-300 cc.) and cautiously remove
the stopper, making sure that no particles fall from it or from the
tube elsewhere than in the beaker. Pour out from the tube a portion
of the carbonate, replace the stopper and determine approximately how
much has been removed. Continue this procedure until 1.00 to 1.10
grams has been taken from the tube. Then weigh the tube accurately
and record the weight under the first weight in the notebook.
The difference in the two weights is the weight of the carbonate
transferred to the beaker. Proceed in the same way to transfer a
second portion of the carbonate from the tube to another beaker of
about the same size as the first. The beakers should be labeled and
plainly marked to correspond with the entries in the notebook.
Pour over the carbonate in each beaker about 80 cc. of water, stir
until solution is complete, and add two drops of methyl orange
solution. Fill the burettes with the standard acid and alkali
solutions, noting the initial readings of the burettes and temperature
of the solutions. Run in acid from the burette, stirring and avoiding
loss by effervescence, until the solution has become pink. Wash down
the sides of the beaker with a !little! water from a wash-bottle, and
then run in alkali from the other burette until the pink is replaced
by yellow; then finish the titration as described on page 37. Note the
readings of the burettes after the proper interval, and record them in
the notebook. Repeat the procedure, using the second portion of sodium
carbonate. Apply the necessary calibration corrections to the volumes
of the solutions used, and correct for temperature if necessary.
From the data obtained, calculate the volume of the hydrochloric
acid solution which is equivalent to the volume of sodium hydroxide
solution used in this titration. Subtract this volume from the volume
of hydrochloric acid. The difference represents the volume of acid
used to react with the sodium carbonate. Divide the weight of sodium
carbonate by this volume in cubic centimeters, thus obtaining the
weight of sodium carbonate equivalent to each cubic centimeter of the
acid.
From this weight it is possible to calculate the corresponding weight
of HCl in each cubic centimeter of the acid, and in turn the relation
of the acid to the normal.
If, however, it is recalled that normal solutions are equivalent to
each other, it will be seen that the same result may be more readily
reached by dividing the weight in grams of sodium carbonate per cubic
centimeter just found by titration by the weight which would be
contained in the same volume of a normal solution of sodium carbonate.
A normal solution of sodium carbonate contains 53.0 grams per liter,
or 0.0530 gram per cc. (see page 29). The relation of the acid
solution to the normal is, therefore, calculated by dividing the
weight of the carbonate to which each cubic centimeter of the acid is
equivalent by 0.0530. The standardization must be repeated until the
values obtained agree within, at most, two parts in one thousand.
When the standard of the acid solution has been determined, calculate,
from the known ratio of the two solutions, the relation of the sodium
hydroxide solution to a normal solution (Notes 1 and 2).
[Note 1: In the foregoing procedure the acid solution is standardized
and the alkali solution referred to this standard by calculation. It
is equally possible, if preferred, to standardize the alkali solution.
The standards in a common use for this purpose are purified
oxalic acid (H_{2}C_{2}O_{4}.2H_{2}O), potassium acid oxalate
(KHC_{2}O_{4}.H_{2}O or KHC_{2}O_{4}), potassium tetroxalate
(KHC_{2}O_{4}.H_{2}C_{2}O_{4}.2H_{2}O), or potassium acid tartrate
(KHC_{4}O_{6}), with the use of a suitable indicator. The oxalic acid
and the oxalates should be specially prepared to insure purity,
the main difficulty lying in the preservation of the water of
crystallization.
It should be noted that the acid oxalate and the acid tartrate each
contain one hydrogen atom replaceable by a base, while the tetroxalate
contains three such atoms and the oxalic acid two. Each of the two
salts first named behave, therefore, as monobasic acids, and the
tetroxalate as a tribasic acid.]
[Note 2: It is also possible to standardize a hydrochloric acid
solution by precipitating the chloride ions as silver chloride and
weighing the precipitate, as prescribed under the analysis of sodium
chloride to be described later. Sulphuric acid solutions may be
standardized by precipitation of the sulphate ions as barium sulphate
and weighing the ignited precipitate, but the results are not above
criticism on account of the difficulty in obtaining large precipitates
of barium sulphate which are uncontaminated by inclosures or are not
reduced on ignition.]
DETERMINATION OF THE TOTAL ALKALINE STRENGTH OF SODA ASH
Soda ash is crude sodium carbonate. If made by the ammonia process it
may contain also sodium chloride, sulphate, and hydroxide; when made
by the Le Blanc process it may contain sodium sulphide, silicate, and
aluminate, and other impurities. Some of these, notably the hydroxide,
combine with acids and contribute to the total alkaline strength,
but it is customary to calculate this strength in terms of sodium
carbonate; i.e., as though no other alkali were present.
PROCEDURE.--In order to secure a sample which shall represent the
average value of the ash, it is well to take at least 5 grams. As this
is too large a quantity for convenient titration, an aliquot portion
of the solution is measured off, representing one fifth of the entire
quantity. This is accomplished as follows: Weigh out on an analytical
balance two samples of soda ash of about 5 grams each into beakers
of about 500 cc. capacity. (The weighings need be made to centigrams
only.) Dissolve the ash in 75 cc. of water, warming gently, and filter
off the insoluble residue; wash the filter by filling it at least
three times with distilled water, and allowing it to drain, adding the
washings to the main filtrate. Cool the filtrate to approximately the
standard temperature of the laboratory, and transfer it to a 250 cc.
measuring flask, washing out the beaker thoroughly. Add distilled
water of laboratory temperature until the lowest point of the meniscus
is level with the graduation on the neck of the flask and remove any
drops of water that may be on the neck above the graduation by means
of a strip of filter paper; make the solution thoroughly uniform by
pouring it out into a dry beaker and back into the flask several
times. Measure off 50 cc. of the solution in a measuring flask, or
pipette, either of which before use should, unless they are dry on the
inside, be rinsed out with at least two small portions of the soda ash
solution to displace any water.
If a flask is used, fill it to the graduation with the soda ash
solution and remove any liquid from the neck above the graduation with
filter paper. Empty it into a beaker, and wash out the small flask,
unless it is graduated for !delivery!, using small quantities of
water, which are added to the liquid in the beaker. A second 50 cc.
portion from the main solution should be measured off into a second
beaker. Dilute the solutions in each beaker to 100 cc., add two drops
of a solution of methyl orange (Note 1) and titrate for the alkali
with the standard hydrochloric acid solution, using the alkali
solution to complete the titration as already prescribed.
From the volumes of acid and alkali employed, corrected for burette
errors and temperature changes, and the data derived from the
standardization, calculate the percentage of alkali present, assuming
it all to be present as sodium carbonate (Note 2).
[Note 1: The hydrochloric acid sets free carbonic acid which is
unstable and breaks down into water and carbon dioxide, most of which
escapes from the solution. Carbonic acid is a weak acid and, as such,
does not yield a sufficient concentration of H^{+} ions to cause the
indicator to change to a pink (see page 32).
The chemical changes involved may be summarized as follows:
2H^{+}, 2Cl^{-} + 2Na^{+}, CO_{3}^{--} --> 2Na^{+}, 2Cl^{-} +
[H_{2}CO_{3}] --> H_{2}O + CO_{2}]
[Note 2: A determination of the alkali present as hydroxide in soda
ash may be determined by precipitating the carbonate by the addition
of barium chloride, removing the barium carbonate by filtration, and
titrating the alkali in the filtrate.
The caustic alkali may also be determined by first using
phenolphthalein as an indicator, which will show by its change from
pink to colorless the point at which the caustic alkali has been
neutralized and the carbonate has been converted to bicarbonate, and
then adding methyl orange and completing the titration. The amount of
acid necessary to change the methyl orange to pink is a measure of one
half of the carbonate present. The results of the double titration
furnish the data necessary for the determination of the caustic alkali
and of the carbonate in the sample.]
DETERMINATION OF THE ACID STRENGTH OF OXALIC ACID
PROCEDURE.--Weigh out two portions of the acid of about 1 gram
each. Dissolve these in 50 cc. of warm water. Add two drops of
phenolphthalein solution, and run in alkali from the burette until the
solution is pink; add acid from the other burette until the pink is
just destroyed, and then add 0.3 cc. (not more) in excess. Heat the
solution to boiling for three minutes. If the pink returns during the
boiling, discharge it with acid and again add 0.3 cc. in excess and
repeat the boiling (Note 1). If the color does not then reappear, add
alkali until it does, and a !drop or two! of acid in excess and boil
again for one minute (Note 2). If no color reappears during this time,
complete the titration in the hot solution. The end-point should be
the faintest visible shade of color (or its disappearance), as the
same difficulty would exist here as with methyl orange if an attempt
were made to match shades of pink.
From the corrected volume of alkali required to react with the
oxalic acid, calculate the percentage of the crystallized acid
(H_{2}C_{2}O_{4}.2H_{2}O) in the sample (Note 3).
[Note 1: All commercial caustic soda such as that from which the
standard solution was made contains some sodium carbonate. This reacts
with the oxalic acid, setting free carbonic acid, which, in turn,
forms sodium bicarbonate with the remaining carbonate:
H_{2}CO_{3} + Na_{2}CO_{3} --> 2HNaCO_{3}.
This compound does not hydrolyze sufficiently to furnish enough OH^{-}
ions to cause phenolphthalein to remain pink; hence, the color of
the indicator is discharged in cold solutions at the point at which
bicarbonate is formed. If, however, the solution is heated to boiling,
the bicarbonate loses carbon dioxide and water, and reverts to sodium
carbonate, which causes the indicator to become again pink:
2HNaCO_{3} --> H_{2}O + CO_{2} + Na_{2}CO_{3}.
By adding successive portions of hydrochloric acid and boiling, the
carbonate is ultimately all brought into reaction.
The student should make sure that the difference in behavior of the
two indicators, methyl orange and phenolphthalein, is understood.]
[Note 2: Hydrochloric acid is volatilized from aqueous solutions,
except such as are very dilute. If the directions in the procedure
are strictly followed, no loss of acid need be feared, but the amount
added in excess should not be greater than 0.3-0.4 cc.]
[Note 3: Attention has already been called to the fact that the color
changes in the different indicators occur at varying concentrations
of H^{+} or OH^{-} ions. They do not indicate exact theoretical
neutrality, but a particular indicator always shows its color change
at a particular concentration of H^{+} or OH^{-} ions. The results
of titration with a given indicator are, therefore, comparable. As a
matter of fact, a small error is involved in the procedure as outlined
above. The comparison of the acid and alkali solutions was made, using
methyl orange as an indicator, while the titration of the oxalic acid
is made with the use of phenolphthalein. For our present purposes the
small error may be neglected but, if time permits, the student is
recommended to standardize the alkali solution against one of the
substances named in Note 1, page 41, and also to ascertain
the comparative value of the acid and alkali solutions, using
phenolphthalein as indicator throughout, and conducting the titrations
as described above. This will insure complete accuracy.]
II. OXIDATION PROCESSES
GENERAL DISCUSSION
In the oxidation processes of volumetric analysis standard solutions
of oxidizing agents and of reducing agents take the place of the acid
and alkali solutions of the neutralization processes already studied.
Just as an acid solution was the principal reagent in alkalimetry, and
the alkali solution used only to make certain of the end-point, the
solution of the oxidizing agent is the principal reagent for the
titration of substances exerting a reducing action. It is, in general,
true that oxidizable substances are determined by !direct! titration,
while oxidizing substances are determined by !indirect! titration.
The important oxidizing agents employed in volumetric solutions are
potassium bichromate, potassium permangenate, potassium ferricyanide,
iodine, ferric chloride, and sodium hypochlorite.
The important reducing agents which are used in the form of standard
solutions are ferrous sulphate (or ferrous ammonium sulphate), oxalic
acid, sodium thiosulphate, stannous chloride, arsenious acid, and
potassium cyanide. Other reducing agents, as sulphurous acid,
sulphureted hydrogen, and zinc (nascent hydrogen), may take part in
the processes, but not as standard solutions.
The most important combinations among the foregoing are: Potassium
bichromate and ferrous salts; potassium permanganate and ferrous
salts; potassium permanganate and oxalic acid, or its derivatives;
iodine and sodium thiosulphate; hypochlorites and arsenious acid.
BICHROMATE PROCESS FOR THE DETERMINATION OF IRON
Ferrous salts may be promptly and completely oxidized to ferric salts,
even in cold solution, by the addition of potassium bichromate,
provided sufficient acid is present to hold in solution the ferric and
chromic compounds which are formed.
The acid may be either hydrochloric or sulphuric, but the former is
usually preferred, since it is by far the best solvent for iron and
its compounds. The reaction in the presence of hydrochloric acid is as
follows:
6FeCl_{2} + K_{2}Cr_{2}O_{7} + 14HCl --> 6FeCl_{3} + 2CrCl_{3} + 2KCl
+ 7H_{2}O.
NORMAL SOLUTIONS OF OXIDIZING OR REDUCING AGENTS
It will be recalled that the system of normal solutions is based upon
the equivalence of the reagents which they contain to 8 grams of
oxygen or 1 gram of hydrogen. A normal solution of an oxidizing agent
should, therefore, contain that amount per liter which is equivalent
in oxidizing power to 8 grams of oxygen; a normal reducing solution
must be equivalent in reducing power to 1 gram of hydrogen. In order
to determine what the amount per liter will be it is necessary to know
how the reagents enter into reaction. The two solutions to be employed
in the process under consideration are those of potassium bichromate
and ferrous sulphate. The reaction between them, in the presence of an
excess of sulphuric acid, may be expressed as follows:
6FeSO_{4} + K_{2}Cr_{2}O_{7} + 7H_{2}SO_{4} --> 3Fe_{2}(SO_{4})_{3} +
K_{2}SO_{4} + Cr_{2}(SO_{4})_{3} + 7H_{2}O.
If the compounds of iron and chromium, with which alone we are now
concerned, be written in such a way as to show the oxides of these
elements in each, they would appear as follows: On the left-hand side
of the equation 6(FeO.SO_{3}) and K_{2}O.2CrO_{3}; on the right-hand
side, 3(Fe_{2}O_{3}.3SO_{3}) and Cr_{2}O_{3}.3SO_{3}. A careful
inspection shows that there are three less oxygen atoms associated
with chromium atoms on the right-hand side of the equation than on the
left-hand, but there are three more oxygen atoms associated with iron
atoms on the right than on the left. In other words, a molecule of
potassium bichromate has given up three atoms of oxygen for oxidation
purposes; i.e., a molecular weight in grams of the bichromate (294.2)
will furnish 3 X 16 or 48 grams of oxygen for oxidation purposes.
As this 48 grams is six times 8 grams, the basis of the system, the
normal solution of potassium bichromate should contain per liter one
sixth of 294.2 grams or 49.03 grams.
A further inspection of the dissected compounds above shows that six
molecules of FeO.SO_{3} were required to react with the three atoms of
oxygen from the bichromate. From the two equations
3H_{2} + 3O --> 3H_{2}O
6(FeO.SO_{3}) + 3O --> 3(Fe_{2}O_{3}.3SO_{3})
it is plain that one molecule of ferrous sulphate is equivalent to one
atom of hydrogen in reducing power; therefore one molecular weight in
grams of ferrous sulphate (151.9) is equivalent to 1 gram of
hydrogen. Since the ferrous sulphate crystalline form has the formula
FeSO_{4}.7H_{2}O, a normal reducing solution of this crystalline salt
should contain 277.9 grams per liter.
PREPARATION OF SOLUTIONS
!Approximate Strength 0.1 N!