William Thomas Boone.

A complete course of volumetric analysis for middle and higher forms of schools online

. (page 2 of 11)
Online LibraryWilliam Thomas BooneA complete course of volumetric analysis for middle and higher forms of schools → online text (page 2 of 11)
Font size
QR-code for this ebook

away to drain dry immediately after use. Time
spent in washing apparatus is thereby minimized,
errors due to "dirty" apparatus are avoided, and
corrosion and staining of the vessels are prevented.



Exercise 6. Determine the end-point of the re-
action between a definite quantity of sulphuric acid
and barium chloride.

Directions. i. Specially dilute a few cubic centimetres
of the bench solutions of these reagents with four or five
times as much distilled water.

2. Boil up 10 to 15 cu. cm. of the acid, and drop in the
BaCl 2 cautiously as long- as fresh precipitate appears to form.


3. Boil for a minute or more, then allow the liquid to

4. Add further drops, one at a time, as long as each
produces more cloudiness.

5. Test your result. Decant some of the clear liquid (or
filter if necessary) and divide into two parts : (a) and (b).
Test (a) for free H 2 SO 4 with litmus, also by noting whether
milkiness appears on adding- more BaCl. 2 . Test (b) for
excess of BaCl. 2 by noting whether it becomes milky on
adding Na 2 CO 3 or more H 2 SO 4 . What do you find?

Exercise 7. Determine the end-point of the re-
action between some ferric chloride and sodium

Directions. i. Use very dilute solutions.

2. Boil up the FeCl 3 solution, and add to it the NaOH,
proceeding generally as in Ex. 6. Make notes on the
colour changes you observe, together with any peculiarity
connected with the precipitation of the ferric hydroxide,
Fe(OH) 3

3. Test your result. Filter, and stand the beaker on a
white surface. A coloured filtrate shows FeCl 3 to be still
present. To a portion add either K 4 FeC 6 N 6 or NH 4 NCS,
and account for what you observe.

A colourless filtrate may contain excess of NaOH. For
its detection use one drop of phenolphthalein. A pink
colour shows the presence of alkali.


Exercise 8. Find the end-point of the reaction
between a definite quantity of potassium perman-
ganate and sulphuretted hydrogen.'

Directions. i. Use some fairly strong solution of KMnO 4 ,
and dilute a fresh solution of H 2 S to twice its bulk.

2. Test each solution with litmus paper and note what
you observe. The highly-coloured KMnO 4 can be washed


off the paper. (If a burette is used for the KMnO 4 , it
should have a glass stopcock.)

3. Test the result of your determination by filtering- off
precipitate and noting whether the filtrate is (a) colourless,
and therefore contains no excess of KMnO 4 ; (b) odourless,
and consequently contains no free H 2 S.

4. Let pieces of red and blue litmus paper lie in the
filtrate for some time, and account for what you find.
Compare with observation in 2.

5. Repeat the exercise after acidifying the KMnO 4 with
about half its volume of dilute H 2 SO 4 . The filtrate may
be tested for excess of H 2 S with AgNO 3 . Ag 2 S is black;
Ag 2 SO 4 is white.

Question 2. What difference does the addition of
sulphuric acid make to the reaction? Write equations, ex-
pressing what has happened both with and without acid.

Exercise 9. Try to find the end-point of the re-
action between potassium dichromate and sulphur
dioxide solutions (sulphurous acid).

Directions. Test both with litmus paper, and state what
you find.

Question 3. Do you find it necessary to add sulphuric
acid in order to keep the liquid clear?

Question 4. Can you suggest a method of ensuring
completion of the reaction in Ex. 9 without having any
excess remaining of either original substance? State any
change in colour noted, and write an equation expressing
the reaction which has taken place.


Exercise 10. Find the end-point of the reaction
between some common salt and silver nitrate.

Directions. i. Use very dilute NaCl solution, to avoid
waste of AgNO 3 .

( C 906 ) 3


2. First, try to determine the end-point by noting- when
precipitation of AgCl ceases. The precipitate will clot on
shaking well, especially after adding- a little dilute HNO 3
or a few drops of CS 2 .

3. Test the result. Filter, and add each reagent to
separate portions of filtrate. There should be no fresh
precipitate in either case.

4. Find the effect of adding one drop of K 2 CrO 4 solution
to fresh portions of each reagent. Note. Both should be
quite neutral.

5. To that in which a red precipitate occurs, add NaCl,
and shake up. If no permanent change is noted, add more
chloride, and shake again.

6. From your observations, make use of K 2 CrO 4 in
order to make a more accurate determination of the
end-point. Remember that AgCl is somewhat soluble in
presence of NaCl.

Question 5. What is the red precipitate in Ex. 10?
Explain what happened on adding sodium chloride.

Question 6. What may be observed when, to a mixed
solution of sodium chloride and potassium dichromate,
silver nitrate is added in excess? What conclusion do
you draw from the permanence of the red precipitate?

Common Indicators


To be a good indicator, a substance should become
coloured when in the presence of traces of the sub-
stance to be detected. Its colour should be easily
distinguished, and any change in colour should be


sharp and capable of instant recognition. It is well
to become acquainted with more than one indicator
of a substance, since the colour of one may be affected
by the presence of other substances, and the same
colours cannot be distinguished equally well by dif-
ferent persons. Of the substances used to detect
acids and alkalis the commonest are litmus, methyl
orange, phenolphthalein, and turmeric. They are
all organic compounds of complex structure.

Exercise 11. Prepare a purple solution of litmus.

Directions. i. Crush up 10 grm. of granular litmus,
and stir it into 100 cu. cm. of hot water. Allow the
undissolved particles to settle, and decant
the extract.

2. Digest the residue with about the same
quantity of hot water for an hour. After
settling, add the two extracts and put by
till next day.

3. Decant, or, if necessary, filter off the
clear liquid, which will probably be blue.
Add specially diluted HNO 3 , a drop at a
time, till a delicate purple tint is got.

4. Keep the prepared solution in an open

Fig. 9 shows a suitable arrangement.
The cork forms a loose cap on the rim of
the bottle, and is pierced by a simple pipette;
i.e. a tube about six inches long nearly fused up at the
lower end. Adjust it to withdraw a suitable volume of

Preservation of the solution.

(i) To prevent growth of mould, add a few crystals of
phenol or a few drops of chloroform. Does phenol or
"carbolic acid" as it is commonly called show an acid
reaction with the litmus?

Fig. 9


(ii) To restore the colour, should it fade, expose the
liquid to the free air in an open dish.

Exercise 12. Examine the behaviour of litmus
with various substances.

Directions. i. Besides ordinary acids and alkalis, test
it with boric, phosphoric, and some organic acids. Also
with solutions of such salts as borax, disodium phosphate,
sodium carbonate, and sodium bicarbonate. Also with
lime-water, sulphuretted hydrogen solution, distilled water
that has been exposed in the laboratory, and some that
has just been made, and solution of carbon dioxide.

Note. In making a solution of carbon dioxide, see that no
acid used to generate the gas is allowed to pass over with it.
Bubble the gas slowly through water contained in a wash bottle,
and make the solution with the washed gas.

2. Boil some solution of carbon dioxide, to which litmus
has been added, as long as any change in colour occurs.

3. Make notes of your observations, and as far as pos-
sible account for what you have seen.

Exercise 13. Prepare a solution of methyl orange,
and examine its behaviour with acids, alkalis, &c.

Directions. i. Dissolve o. i grm. in 100 cu. cm. of dis-
tilled water, and keep in a bottle fitted with a cork carrying
a simple pipette. (See fig. 9.)

2. Find the effect of the solution on litmus paper.

3. Place one drop in each of two flasks containing about
50 cu. cm. of water and standing on a white surface.
Make the water in one of the flasks slightly acid; after-
wards make it alkaline. (This may be known by its effect
on litmus paper.) Compare the colours obtained with that
of the neutral liquid.

4. Try the effect on the indicator of such substances as
those named in Ex. 12, especially some carefully-prepared
solution of carbon dioxide.

Question 7. How do the indications of litmus and


methyl orange differ with respect to solution of carbon
dioxide? (We may assume the presence of carbonic acid
in this solution.)

Exercise 14. Prepare a solution of phenolphtha-
lein, and examine its behaviour with acids, alkalis,

Directions. i. Dissolve 0.5 rm. of the solid in 50 cu. cm.
of alcohol and keep in a corked bottle. The cork may
carry a simple pipette. (Fig. 9.)

2. Try the effect of one drop of the indicator on solu-
tions of the following: Mineral and organic acids, alkalis,
alkaline carbonates both normal and acid.

3. To a solution of acid sodium carbonate add one drop
of phenolphthalein, and heat. Cool, and then pass carbon
dioxide washed free from acid into the solution until a
decided colour change occurs.

4. To a solution of ammonia add a drop of the indicator.
Now dilute with water until the colour just disappears.
Can you detect the presence of the gas by its odour? Add
a few drops of purple or reddish litmus.

Question 8. In what general way does phenolphthalein
differ from the two indicators previously examined? How
could you make use of your observations of its behaviour
with acid sodium carbonate? Did you find the coloration
due to caustic soda permanent?

Exercise 15. Examine the behaviour of turmeric
with alkalis and acids.

Directions. i. Use turmeric paper; i.e. paper strips
which have been soaked in an alcoholic extract of the
colouring matter and dried. Dip some strips into very
dilute alkali, and dry them, noting the colour changes.

2. Try the effect of ordinary acids and alkalis upon
turmeric; also coloured liquids, such as vinegar, potassium
manganate, and permanganate; also saliva.


Question 9. Do turmeric and litmus give the same
indications with saliva?


Starch, especially when in solution, is an excellent
indicator for free iodine.

Potassium dichromate (K 2 Cr. 2 O 7 ) has already been
used as an internal indicator in the reaction between
a silver salt and a haloid. (Ex. 10.)

Where coloured liquids are concerned it is obviously
impossible to use indicators in the way previously
considered. For instance, the indications afforded by
a solution of litmus when added to common vinegar
can scarcely be relied upon to show when sufficient
alkali has been added to neutralize the acid of the
vinegar. A bit of litmus paper dropped into the
vinegar will do so, but not sharply. A better way is
to apply small drops of the acid liquid, withdrawn at
short intervals, either to a piece of the purple paper
laid on the bench, or to drops of purple litmus solu-
tion previously placed upon a white tile. By this
means not only can neutralization be detected, but also
its approach. Indicators when used in this way are
distinguished as " external", or as "drop reagents".
The end-points of many reactions can be detected by
their means, but some practice is necessary to obtain
results as accurate as with internal indicators. Two
external indicators in very common use are iodized
starch paper and potassium ferricyanide.

Exercise 16. Prepare a " starch solution" and
examine its behaviour with iodine.

Directions. i. Mix i grm. of potato starch, or, better,


" soluble" starch, with 10 cu. cm. of cold water, and pour
into it about 200 cu. cm. of boiling water. Should a milky
appearance persist, boil up for two minutes not longer,
or the starch will be converted to dextrose.

2. Filter off cloudy matter.

3. Add i cu. cm. of the clear solution to some water,
and then one drop of solution of iodine. Boil some of the
blue "iodide of starch" obtained, for a minute or two;
allow to cool, and state what happens.

4. To some dilute iodine solution add sodium thio-
sulphate (Na 2 S 2 O 3 ) drop by drop, till the colour just dis-
appears. Add i cu. cm. of starch solution. Was all the
iodine used up?

5. Find the effect of Na 2 S 2 O 3 on some of the iodide
of starch obtained in 3. (See note on Ex. 16, p. 151.)

Exercise 17. - Prepare some "iodized starch
paper", and with it determine when sufficient
chlorine water has been added to some sodium
arsenite solution to oxidize it to arsenate.

Directions. i. To prepare the paper. Dissolve about
0.5 grm. of pure KI in 100 cu. cm. of the starch solution
prepared in Ex. 16. Soak some clean white filter paper in
it, dry, and cut into strips. Preserve it (from chlorine and
other fumes) in a well-corked bottle.

2. To examine the effect of each reagent on the indicator.
(i) Moisten a strip with pure water; place a drop of each
solution upon it at different places, and note. N.B. With-
draw the solutions on separate, clean, thin glass rods,
(ii) See whether very dilute chlorine water produces the
same effect as a concentrated solution.

3. To determine the end-point. Drop the chlorine water
preferably from a burette into the arsenite solution
until a drop of the liquid withdrawn on a rod, and applied
to the starch paper, produces a faint blue colour. Shake
up after each addition of chlorine water. Take care that
the successive drops touch different parts of the paper.


4. Try the reverse process, viz. that of adding the
arsenite to chlorine water till the end-point is reached.
Which do you consider the better method? (See note on
Ex. 17, p. 151.)

Exercise 18. Oxidize some ferrous sulphate to
the ferric state by means of potassium dichromate.
Determine the end-point of the reaction with potas-
sium ferricyanide.

Directions. i. Try the effect of one of the substances
upon each of the others.

(i) Add K 2 Cr,O 7 solution to FeSO 4 solution.

(ii) Repeat (i) after acidifying the iron solution with
dilute H 2 SO 4 .

(iii) Add freshly - prepared solution of ferricyanide
(K 3 FeC 6 N 6 ) to solutions of (a) FeSO 4 ; (b) K 2 Cr 2 O 7 ; (c)
FeCl 3 or other ferric salt.

(iv) Add K 2 Cr 2 O 7 drop by drop to FeSO 4 solution pre-
viously acidified, and containing one or two drops of the
indicator. Does the blue colour disappear when K 2 Cr 2 O 7
is added in excess?

2. Determine the end-point.

(i) Place about a dozen separate drops of indicator on a
white tile or plate.

(ii) Acidify a fresh portion of FeSO 4 , and drop into it
K 2 Cr 2 O 7 until a bright-green solution is obtained. Shake
the mixture. Withdraw a drop on the end of a thin glass
rod, and mix it with one of the drops of K 3 FeC (; N 6 . A
blue colour shows ferrous salt to be still present. Add
more K 2 Cr 2 O 7 and test again, repeating until the blue
colour fails to appear.

Question 10. Could you detect the end-point or com-
pletion of oxidation of the ferrous salt without using the
indicator? Give your reasons.

Question ii. What is the blue precipitate obtained
with the indicator?


Question 12. What inference would you make if the
liquid produced in Ex. 18, 2, became yellowish green?
And what indication is there in the successive trials that
the end-point is being- approached?

N.B. Some notes on indicators will be found in Chapter XL

Standard or Volumetric Solutions

A standard or volumetric solution is made with
water and usually contains a known weight of some
definite chemical substance in i cu. cm. This weight
is termed the " strength of the solution ". Sometimes
the term refers to the weight of some other substance
with which i cu. cm. of the solution reacts.

The strength of volumetric solutions is low: usually
there is less than one molecular or formula weight in
grammes per litre. When exactly this quantity is
present, the solution is sometimes spoken of as a
" molecular" or " molar" solution. Most volumetric
solutions contain no more than one-tenth of this
weight of solute to the litre; as, while yielding results
as accurate as more concentrated solutions, they are
found more convenient to work with.

The strength of a " molar " solution of caustic soda
is given thus:

i cu. cm. = 0.04001 grm. NaOH,

because i cu. cm. contains one-thousandth part of
40.01 grm. of the alkali whose formula weight is
40.01. [O = 16, H = 1.008, Na = 23.]


It is, however, far more common to use solutions
of what is called "normal" strength, or those which
bear some simple ratio to this value, such as deci-
normal, twice normal, &c., and to express their

strength by the abbreviations N., , 2 N., &c.

A normal solution contains per litre either (i) a
gramme equivalent of its solute, or (ii) that weight
of it which can furnish in some special reaction the
gramme equivalent of one of its constituents. It may,
or may not, be molar.

Since the acidity of acids is due to hydrogen, normal
solutions of these compounds will contain 1.008 grm.
of the acid-forming hydrogen to the litre. As a con-
sequence, molar and normal solutions of monobasic
acids are the same; but normal solutions of dibasic
acids are semi-molar, and those of tri basic acids are
one-third molar strength.

Of hydrochloric acid (HC1), a molar solution con-
tains 36.468 grm. of hydrogen chloride per litre
36.468 being its molecular weight; and a normal
solution contains the same weight because it includes
the gramme equivalent (1.008) of acid-forming hydro-
gen per litre.

Of sulphuric acid (H 2 SO 4 ), a molar solution con-
tains 98.086 grm. of hydrogen sulphate per litre; but
a normal solution contains only half this weight or
49.043 grm., since the former solution contains two
gramme equivalents of acid-forming hydrogen.

With regard to alkalis, i 1. of normal solution
must contain that weight of alkali which can exactly
neutralize i 1. of normal acid; that is, it must contain
the gramme equivalent both of the metal and of
hydroxyl; for the metal replaces 1.008 grm. of hy-
drogen to form the salt, while the hydroxyl unites


with this hydrogen to form water. This will be seen
from the following equation :

HC1 + NaOH = NaCl + HOH

1.008 + 35.46 23 + 16 4- 1.008 23 + 35.46 1.008 + 16 + i. 008
36.468 40.008 58.46 18.016

It follows that equal volumes of normal acid and
alkali neutralize each other when mixed.

Normal caustic soda (NaOH) therefore must contain
40.008 grm. of sodium hydroxide, or 23 grm. of
sodium to the litre, as these weights are necessary to
displace the equivalent weight (1.008 grm.) of hy-
drogen from an acid. The equation shows that a
molar solution is of the same strength.

Similarly, it will be found that normal and molar
solutions of potassium hydroxide (KOH) and am-
monia (NH 3 ) respectively are of the same strength.

It must be remembered that acids may contain hy-
drogen which is quite independent of that conferring
acid properties. Only one-fourth of the hydrogen in
acetic acid (C 2 H 4 O 2 ) is so concerned, and its formula
is usually written CH 3 .COOH to express this fact.
Only the hydrogen in the carboxyl group (COOH)
is acid-forming. That of the methyl group (CH 3 ) is
not so; for the acidity of acetic acid is not lessened
when this hydrogen is wholly replaced by chlorine.
N. acetic acid therefore contains a mol of the acid to
the litre.

The following points should be well noted, as they
may aid you considerably to shorten your work when
using standard acids and alkalis:

1. Normal acid solution exactly neutralizes an equal
volume of normal alkali.

2. Normal solutions of acids are of equal " strength",
for they can all neutralize an equal volume of any


normal alkali. [Equal volumes of N. acids contain
the same weight of acid-forming hydrogen and equiva-
lent weights of acid.]

3. Normal solutions of alkalis are also of equal
strength, as they can all neutralize an equal volume
of normal acid. [Equal volumes of N. alkalis contain
the same weight of hydroxyl and equivalent weights
of metals.]

4. The statements under 2 and 3 equally apply
to decinormal and other solutions of corresponding

5. The gramme equivalent of an acid contains 1.008
grm. of that hydrogen replaceable by a metal.

6. The gramme equivalent of a base is that weight
which neutralizes a gramme equivalent of acid, and
can provide 17.008 grm. of hydroxyl to unite with the
displaced hydrogen.

Besides acids and alkalis, standard solutions of
"salts" are used. A mixture of i 1. of N. HC1 and i 1.
of N. NaOH yields 2 1. of sodium chloride solution,
containing only one gramme equivalent each of sodium
and chlorine. This solution is therefore seminormal.
In calcium chloride (CaCl. 2 ) there are two equivalents
of chlorine united with the calcium ; and in its reaction
with sulphuric acid two equivalents of hydrogen are
displaced by the metal.

CaCl, + H 2 SO 4 = CaSO 4 + 2 HC1.

A solution containing the gramme formula weight of
this salt to a litre is, therefore, of 2 N. strength.

Potassium Permanganate (KMnO 4 ). It may appear
to a beginner that a normal solution of this salt should
contain the gramme formula weight of the compound
to the litre, since this includes an equivalent weight
of potassium. This reagent, however, is not used for


the potassium it contains, but for the oxygen it can
part with;' and it is this consideration which must
guide us in determining the weight necessary to pre-
pare a solution of any required strength. A normal
solution, then, should contain that weight per litre
which will yield the gramme equivalent of oxygen
(8 grm.). An incautious student might assume that
a normal solution should contain one-eighth of the
gramme formula weight to a litre, since this weight
contains 64 grm. of oxygen. He would be wrong,
because all the oxygen of the compound is not parted
with in its reactions with other substances.

The formula of the salt used to be written K 2 Mn 2 O 8 ,
and we may conceive it to be made up of two oxides,
K 2 O and Mn. 2 O 7 , the former basic and the lattei
acidic. The heptoxide of manganese readily parts
with a large proportion of its oxygen to reducing
agents, and the products suggest that the salt breaks
up thus:

K 2 O, Mn 2 Qg = K 2 O + 2 MnO + 5 O
316.1 80

for 80 grm., or ten gramme equivalents of oxygen can
be obtained from 316 grm. of the permanganate. If
sulphuric acid is present, the precipitation of oxide
of manganese is prevented by its conversion into
sulphate. The break-up of the permanganate may
then be represented

2 KMnO 4 + 3 H 2 SO 4 = K 2 SO 4 + 2 MnSO 4

+ 3 H 2 + 50.


A litre of N. permanganate should contain only 31.61
grm. of the salt, because this weight can yield 8 grm.


of oxygen to a reducing agent. It is usual, however,
to make permanganate solution of decinormal strength
for reasons which will be given later.


Potassium permanganate.

i cu. cm. = 0.003161 grm. KMnO 4
i cu. cm. = 0.0008 grm. oxygen.

Potassium dichromate (K 2 Cr 2 O 7 ) is also largely
used for the same purpose as permanganate, as, like
the latter, it easily parts with a portion of its oxygen
to reducing agents. We may consider this compound
also to be made up of basic and acidic oxides, K 2 O
with two of CrO 3 . The latter readily give up a
portion of their oxygen and pass into the chromous

K 2 O, 2 CrOg = K 2 + Cr 2 3 + 30.


This agrees with the fact that 48 grm. of oxygen are
obtainable from 294.2 grm. of the salt. Evidently,
therefore, a normal solution should contain one-sixth
of this weight of the salt to the litre, viz. 49.03 grm.,
as this will yield the necessary 8 grm. (the gramme
equivalent) of oxygen. Acid is used with this re-
agent for the same purpose as with permanganate.

K 2 O, 2 CrO 3 + 8 HC1

= 2 KC1 + 2 CrCl 2 + 4 H 2 + 3 O.

The above examples will serve to illustrate the
principles by which we must be guided in making

2 4 5 6 7 8 9 10 11

Online LibraryWilliam Thomas BooneA complete course of volumetric analysis for middle and higher forms of schools → online text (page 2 of 11)