William Thomas Boone.

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

. (page 8 of 11)
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nitrate to precipitate the whole of the chlorine present.
Calculate the percentage of each chloride in the mixture.

If the powder consisted wholly of KCI, the equation

KC1 + AgNO 3 = AgCl + KNO 3
74.56 169.89

shows that i grm. would require * 9' 9 or 2.2786 (nearly)
grm. of AgNO 3 . 74 ' s6

Similarly, if it consisted wholly of NaCl, i grm. would

require * 9- 9 or 2.9061 (nearly) grm. of AgNO 3 . It is

5 8 -5
evident, therefore, that the mixture requiring 2.812 grm.

must consist mainly of NaCl; and

Weight of NaCl = 2.812 - 2.2786 0.5344
Weight of KC1 2.9061 2.812 0.0941'

Per cent weight of NaCl = '5344 x IOO _ g^ p er cent


Note. In practice, however, good results can be obtained only
when the proportions are not widely different.

Question 77. In another experiment a mixture of
the same two salts was found to contain 55 per cent of
chlorine. Find the proportion of each chloride present.

Had it all been

NaCl there would have been 60.66 per cent Cl.
KCI ,, ,, 47.56

Since 55 per cent Cl has been found

Weight of NaCl . 55 - 47- 5 6 _. 7-44
Weight of KCI " 60.66-55 " 5.66'


The mixture, therefore, consists of 56.8 per cent NaCl
and 43.2 per cent KC1.

Question 78. i grm. of a mixture of sodium chloride
and sodium bromide was found to require 2 grm. of silver
nitrate for complete precipitation. Find the percentage
composition of the mixture.


Silver nitrate behaves somewhat differently with
these compounds. When added to potassium cyanide
solution silver cyanide does not appear as a permanent
precipitate as long as any excess of potassium cyanide
remains in the solution, but forms with it a complex
soluble salt, potassium argenticyanide, KAgC 2 N. 2 .
When no free potassium cyanide remains, further
addition of silver nitrate causes the break up of the
double cyanide, and a precipitate (AgCN) at once

AgNO 3 + KAgQN 2 = 2 AgCN + KNO 3 .

Standard silver nitrate can therefore be used to
estimate potassium cyanide if we bear in mind this
peculiarity. The appearance of precipitate marks the
point when the salts have been mixed in the propor-
tion of one equivalent of silver nitrate to two of
potassium cyanide, as expressed in the equation

AgNO 3 + 2 KCN = KAgC 2 N 2 + KNO & .
169.9 130-2

It will be noted that whereas in titrating halides
the end-point is reached when the precipitate ceases
to form, with cyanides this condition occurs when a
trace of permanent precipitate yznr/ appears. Another
point to be remembered is that metallic halides give

to 900 \ 9


no precipitate with silver nitrate as long as any free
potassium cyanide exists in the solution ; consequently
the appearance of precipitate, even should a halide be
present, marks the end-point of the same reaction.

Exercise 93. Find the percentage of KCN in some
of the solid salt by Liebig's method.

Directions. i. The equation given above shows that a
solution of 2 to 3 grm. of the solid in 100 cu. cm. should

work with '- AgNO 3 . Dissolve up this quantity to 100

cu. cm. and mix well.


2. Use AgNO 3 in burette.

3. Caution. Do not risk sucking up the poisonous solu-
tion into a pipette. Run out the required quantities from
a burette.

4. Stand the beaker upon a dark surface during titration,
to be able to detect the first traces of permanent precipitate.

Question 79. The presence of ammonia in the solu-
tion prevents the precipitation of silver cyanide or chloride.
Can you suggest a way of meeting this difficulty?


Exercise 94. Examine the behaviour of potassium
thiocyanate with silver nitrate in neutral and in
acid (nitric) solution; also when ferric and other
salts are present in addition.

Directions. i. Try simple solutions of (a] AgNO 3 ,
(b) FeCl 3 , (c) Am 2 Fe 2 (SO 4 ) 4 among others, both neutral
and acidified.

2. Add KCNS solution drop by drop (i) to a mixture of
solutions of AgNO 3 and iron alum until some striking
change appears; (ii) to another quantity of similar solution


after acidifying with a little HNO 3 . What effect has the
acid upon the colour due to the iron?

3. Repeat 2 after adding- some other salt in addition,
such as lead acetate, copper sulphate, &c.

The results should show, first, that silver thio-
cyanate is insoluble in water and dilute nitric acid ;
secondly, that in the reaction between silver salt and
thiocyanate in the presence of ferric salt, directly the
silver is precipitated a brown compound is formed
ferric potassium thiocyanate [Fe(CNS) 3 , 9 KCNS,

4 H 2 O], although some other metallic salts are present.
The ferric salt acts as an excellent indicator, as the
brown colour can usually be detected instantly.

For the estimation of silver, standard thiocyanate
offers several advantages over standard sodium
chloride. This is Volhard's method.

Exercise 95. Prepare 500 eu. cm. of ^ potassium

Directions. i. The reaction with AgNO 3 is given by
the equation

AgNO 3 + KCNS = AgCNS + KNO 3 ,

169.9 97- 2

which shows the equivalent quantities of the reacting salts.
Since the thiocyanate is deliquescent, we can only make a
first solution approximate to the required strength. About

5 grm. dissolved up to 500 cu. cm. should produce a solu-
tion slightly above *.

10 N.

2. Find the exact strength by titration with AgNO 3 .

Use the burette for KCNS. The indicator is a cold satu-
rated solution of iron alum, 3 cu. cm. of which should be
added to each 25 cu. cm. of silver solution, together witli
enough HNO 3 (usually about 5 cu. cm. is sufficient) to
destroy the colour. (See notes below, p. 116.)


3. The end-point is reached when the liquid remains
slightly brown after shaking" up. It should be viewed on
a white surface.

4. Calculate the strength of the KCNS solution on the

assumption that AgNO 3 is accurately . The weight of

KCNS in the average volume used = -2Z. of the weight
of AgNO 3 in 25 cu. cm. I '

5. Label the bottle containing the KCNS with the

strength found, or else reduce it to '-. Should it be

N I0

found that 25 cu. cm. ' AgNO 3 = 24.5 cu. cm. KCNS,

evidently 24.5 cu. cm. KCNS should be diluted to 25 cu. cm.
when i cu. cm. = 0.00972 grm. KCNS.

Notes. (i) Specially prepare the nitric acid by diluting con-
centrated acid with an equal volume of water and boiling till
quite colourless.

(ii) If iron alum is not available, boil up a few crystals of
FeSO 4 , 7 H 2 O in the prepared nitric acid.

Exercise 96. The metallic substance provided
consists of silver alloyed with a small proportion
of copper. Find the percentage of silver in it by
Volhard's method.

Directions. Dissolve about 0.5 grm. of the alloy in
HNO 3 ; boil for two or three minutes to remove any
nitrous acid produced; cool and make, up to 100 cu. cm.
Proceed as directed in Ex. 94. Ignore the presence of
copper nitrate.


In alkaline solutions this compound precipitates the
metals zinc, cadmium, nickel, &c., from solutions of
their salts according to the equation

Na 2 S + ZnSO, = ZnS + Na 2 SO 4 ,


so that if the strength of either solution is known we
can determine that of the other, providing the end-
point can be determined with sufficient accuracy.
Since the composition of ordinary commercial sodium
sulphide varies, a solution is usually prepared by
passing sulphuretted hydrogen into caustic soda
(2 NaOH .+ H 2 S = Na 2 S + 2 H.O), and its strength
found by titration againt a suitable standard solution.
This is usually one of zinc salt, since standard sul-
phide is especially used in the estimation of this
metal. As the completion of the reaction cannot be
observed directly (for the precipitated sulphide does
not settle quickly, neither does it show much tendency
to clot), the end-point is found by determining when
sodium sulphide is present in slight excess from its
effect upon some external indicator such as a bright
silver coin, or a drop of solution of lead or nickel
salt, all of which immediately darken. Sodium nitro-
prusside similarly becomes purple. Unfortunately
the same effects are produced upon the indicator by
the zinc sulphide itself; so care must be taken that
the suspended solid is removed before the test is
made. The method of doing this will be explained
in the exercise following. Beginners appear to get
better results with the lead indicator than with the

For practical purposes it is found convenient to
make the solution of sulphide so that i cu. cm. = o.oi
grm. zinc.

Exercise 97. Make some standard sodium sul-
phide solution of strength 1 cu. cm. = 0*01 grm.

Directions. i. Solution roughly of the required strength.
Dissolve about 4 grm. of 94-per-cent NaOH in a little


water, and pass H 2 S into it as long- as the gas is absorbed.
When saturated, the solution should smell strongly. Next,
add bench NaOH till, after shaking up, all odour of H. 2 S
has disappeared ; then dilute with water to i 1. and mix

2. Prepare the following solutions while the H 2 S is being
absorbed, and label each one.

(a) Standard Zinc Solution. Since 287.55 S rm - f zmc
sulphate crystals (ZnSO 4 , 7H 2 O) contain 65.37 grm. of
zinc, dissolve up 10.997 f rm - to i ! m graduated flask.
i cu. cm. of this solution will contain o.oi grm. zinc.

(b) Indicator: alkaline solution of lead. To a few cubic
centimetres of fairly strong solution of sodium tartrate (or
tartaric acid neutralized with NaOH) add solution of lead
acetate, then NaOH till the white precipitate of lead tar-
trate dissolves. Apply heat if necessary.

(c) Ammoniacal Solution. Mix together about 50 cu.
cm. bench ammonium carbonate with about 150 cu. cm.
bench AmOH.

3. Titration. Na.,S solution in burette.

(i) To each portion of zinc solution taken add the am-
moniacal solution a few drops at a time until the pre-
cipitate formed redissolves on shaking.

(ii) For the first rough test, run in the Na 2 S as long as
the precipitate (ZnS) appears to increase; after that about
i cu. cm. at a time until excess is indicated after being
well shaken.

(iii) Test for excess of Na 2 S as follows:

(a) Place one drop of indicator near the edge of a filter
paper. Withdraw a drop of the turbid liquid on another
glass rod, and, holding the paper vertically, apply the drop
about 5 mm. below the alkaline moist patch. The ZnS
withdrawn with the liquid will then remain where the rod
touched the paper, while the liquid will travel upwards to
the indicator. The formation of a brown line marks excess
of Na 2 S. The line is most readily seen when the paper is
held up to the light.


Or (b) lay a filter paper on a white glazed tile, and
damp it with the indicator; then place a drop of the turbid
liquid on another filter paper laid momentarily on the first
and slightly pressed down by the rod. A brown stain on
the lower paper marks the presence of Na 2 S.

4. Dilution. Probably less than 25 cu. cm. of Na 2 S
solution will be required for 25 cu. cm. of zinc solution.
Suppose 15 cu. cm. are found sufficient; evidently these on
dilution to 25 cu. cm. will provide a solution of which i
cu. cm. = o.oi grm. zinc.

Note. Since sodium sulphide solution decomposes on keeping,
its value must be redetermined before future use.

Exercise 98. Assuming 1 that the specimen of
sheet brass (English) consists only of copper and
zinc, find the percentage of zinc present.

Directions. i. Solution of Alloy. Dissolve about i grm.
to a clear blue solution. Use about 5 cu. cm. of concen-
trated HNO 3 . Afterwards add an equal volume of con-
centrated H 2 SO 4 , and heat carefully as long as brown fames
are expelled; but do not evaporate to dryness. Allow to
become cold, then add 100 cu. cm. to 150 cu. cm. of water.
Any white residue should completely dissolve, since English
brass contains practically no lead.

2. Removal of Copper. Add a little dilute HC1, and
precipitate this metal with H 2 S from a hot solution, and
filter off. Immediately wash the precipitate and filter
paper once with diluted bench HC1 and twice with H 2 S
water. Add the wash liquids to the filtrate, which should
be colourless.

[Treat the copper sulphide as follows in preparation for
Ex. 99. Wash it from the paper into a beaker without
loss of time ; any traces left must be dissolved off in dilute
HNO 3 . Similarly dissolve the main portion. When cold,
make up the solution to 100 cu. cm., and keep in a
stoppered bottle till required.]

3. Preparation of Zinc Solution. - Boil down the solution


to less than 100 cu. cm. All H 2 S should now be expelled.
Cool and make up to 100 cu. cm.

4. Titration -with standard Na 2 S. See Ex. 97. As this
solution is acid, rather more Am(OH) will be required to
produce and redissolve the precipitate. Much excess,
however, must be avoided. Note. Accurate results can
only be got by carrying out this titration exactly in the
same way as when the sulphide was standardized.


Parkes*s Method. Though not a precipitation
method, this is given here because in many respects
it resembles that just used for estimating zinc. The
standard cyanide is made up so that i cu. cm. = a
definite weight of copper, and does not necessarily
bear any simple relationship to N. strength. The
reacting solutions are both alkaline, but the precipi-
tation of copper hydroxide is prevented by the addition
of ammonia not in too great excess, however, or the
results are not concordant. To be trustworthy, the
estimation must be carried out in exactly the same
way as the standard solution is made. Although the
method, at least with a beginner, may not be quite
so accurate as that given in Ex. 84, it is of more
general application, and is commonly used by analytical

The addition of ammonia to a solution of copper
sulphate first produces a light-blue precipitate which
dissolves in excess to a dark-blue solution containing
CuSO 4 , 4 NH 3 , H 2 O. The addition of potassium
cyanide destroys this coloured compound, and pro-
duces several substances, among them being the
double cyanide CuCN, KCN. The quantity of cyan-


ide found necessary to effect this change may thus be
used to measure the weight of copper present.

Exercise 99. Prepare a standard solution of
potassium cyanide, so that 1 cu. cm. = 0-005 grm.

Directions. i. Rough Solution. Dissolve about 2Ogrm.
of the cyanide up to 500 cu. cm.

2. Standard Solution of Copper. Weigh out accurately
2.5 grin, of pure copper (foil or wire will do), and dissolve
in about 5 cu. cm. concentrated HNO 3 , mixed with an
equal volume of water. Evaporate down till most of the
acid has been expelled, taking care that there is no loss
by spirting. Transfer to a 5oo-cu.-cm. graduated flask
together with rinsings; cool, and make up to the mark,
i cu. cm. = 0.005 g r m- copper.

3. Standardization of the KCN. - KCN in burette. To
each quantity of copper solution taken (say 25 cu. cm.),
add NaOH till all Cu(OH) 2 is precipitated, then i or 2
cu. cm. of bench ammonia but always the same definite
volume to produce the deep-blue solution. Run in the
KCN till only a faint tint (bluish) remains. Dilute the
KCN solution to make it equivalent to that containing the
copper, i cu. cm. of KCN will now = 0.005 & rm - copper.
(Whatever average volume of KCN is found necessary to
decolorize 25 cu. cm. of copper solution must be diluted to
25 cu. cm., and other volumes in similar proportion.)

Exercise 100. Estimate the percentage weight
of copper in the sample of brass analysed in
Ex. 95.

Directions. Use the solution of copper put by, and'
titrate it exactly as given in Ex. 99. If this has not been>
kept, dissolve a fresh portion of the brass, precipitate the
copper as sulphide, and follow the treatment given in*
Ex. 98.



Basic salts of uranium contain the radical uranyl,
UO 2 , the formula for the acetate being UO,(CH 3 .COO)o,
2 H 2 O. This salt dissolves in water, and can be used
to precipitate phosphoric acid in presence of acetic
acid as uranyl hydrogen phosphate, U(XHPO 4 . The

y0. 2 (CH 3 .COO) 2 , 2H 2 + NaNH 4 HP0 4 , 4 H 2 O

424 6 209. 2

= UO. 2 HPO 4 + CH 2 .COONa + CH 3 .COONH 4 -f6H 2 O

indicates that 424.6 grm. of uranyl acetate = 209.2 grm.
of microcosmic salt (or 71 grm. of P 2 O 6 when united
with alkali). A slight excess of the former salt pro-
duces a brown coloration with potassium ferrocyanide.
But since the indications may be somewhat modified
by the presence of other salts, or, generally speaking,
by any difference in character from the solution origin-
ally used, it is obviously necessary that all estimations
must be carried out as nearly as possible under con-
ditions similar to those of standardization. Points to
be borne in mind are (i) that the purity and state of
hydration of ordinary commercial phosphates cannot
be relied upon ; (ii) that solutions standardized against
alkaline phosphate can only be used in estimating other
alkaline phosphates; for when phosphoric acid is com-
bined in such substances as bone ash or other similar
manure, mineral acid is required for its solution,
and salts are produced which may seriously interfere
with the indications of the ferrocyanide; (iii) that the
mineral acid so used must be neutralized by ammonia,
or it will prevent the precipitation of the uranyl phos-


phate; (iv) that any phosphate precipitated by the
ammonia may be redissolved in acetic acid in which
the uranyl compound is insoluble; and (v) for the
estimation of P 2 O 5 in bone ash, &c., the uranyl solu-
tion must be standardized against tricalcic phosphate,
whose percentage purity is usually found by a gravi-
metric method.

Exercise 101. Make a standard solution of uranyl
acetate, so that 1 cu. cm. = 0'005 grm. P. 2 5 .

Directions. i. Solution of approximate strength. Dis-
solve about 1 6 grm. of uranyl acetate crystals in water,
add 25 cu. cm. glacial acetic acid, and make up to 500
cu. cm. The acetic acid renders the solution less liable to
change under the influence of light. The theoretical weight
of the pure salt (14.94 f rm -) might be expected to give the
required solution; but allowance has to be made for the
interfering action of other salts; and the exact relationship
between the weights of uranyl salt and phosphate when
the end-point is reached (as shown by the indicator) must
be determined by a special test. When this is known the
uranyl solution may be used to estimate any similar
phosphate. Of alkaline phosphates that of most reliable
composition is KH.,PO 4 (W. B. Giles), although carefully
selected crystals of microcosmic salt are almost as good.

2. Solution of Microcosmic Salt. Since 418.4 grm.
contain 142.1 grm. of P 2 O 5 , 5 grm. of the oxide will be

contained in ^ ^ X <; 2frm. of salt. Dissolve ith of this

weight (3.681 grm.) of selected crystals up to 250 cu. cm.
i cu. cm. = 0.005 rm - *W

3. Titration. Uranyl acetate in burette. Boil each
quantity of NaNH 4 HPO 4 solution before adding the uranyl
solution. For first rough test run it in as long as the pre-
cipitate increases; then add i cu. cm. at a time till excess
is indicated.

External indicator', a fresh solution of K 4 FeCy 6 , several


drops of which are placed upon a white tile. When excess
of uranyl acetate is suspected, apply a drop withdrawn in
the usual way. Excess is shown by a faint brown colora-
tion due to uranium ferrocyanide.

4. Label the uranyl solution. Suppose it is found that
23.8 cu. cm. = 25 cu. cm. of the solution of phosphate,

evidently i cu. cm. = 0.005 X ~ D grm. P 2 O-, for use in
the estimation of alkali phosphates.

Exercise 102. By means of the standardized
uranyl solution, estimate the percentage weight of
Na,HP0 4 in a laboratory sample. How does it com-
pare with the theoretical value?

Directions. i. Weigh out 6.3 grm., and dissolve up to*
250 cu. cm. A little calculation will show that i cu. cm.
of this solution should contain 0.005 T rm ' ^2^5 if tne sa ^
is pure, and that 25 cu. cm. should require the same
volume of uranyl solution as used in the last exercise.

2. Precaution. Since ammonium salts are specially
active in modifying the colour produced by the K 4 FeC (5 N (5 ,
and ammonium salt was present in the phosphate used to
standardize the uranyl acetate, a little ammonium salt
must now be specially added, to make the conditions as
nearly alike as possible. To each quantity of Na 2 HPO 4
solution taken for titration, add i cu. cm. of bench AmOH>
then acetic acid to neutralize it.

3. Uranyl solution in burette, and follow the directions
given in Ex. 101.

Exercise 103. Assuming that the laboratory
sample of tricalcic phosphate is pure, find the per-
centage of P 2 5 in some bone ash.

Directions. i. Make a Standard Solution 0/"Ca a (PO 4 )o.
Weigh out accurately 2.73 grm. of dry Ca 3 (PO 4 )., into a
25O-cu.-cm. graduated flask, and dissolve in as little HC1
as possible. Add AmOH in sufficient quantity to neutralize


the free mineral acid, and precipitate the Ca 3 (PO 4 ) ; then
redissolve in acetic avoiding much excess in either case.
Add water to the mark. A little calculation will show
that if the Ca 3 (PO 4 ) 2 is pure, the solution will contain
0.005 g" rm ' fV^s per cubic centimetre.

2. Standardize the uranyl solution against the calcium
phosphate. Uranyl acetate in burette, and follow generally
the directions given in Ex. 101. Note. Since boiling is
apt to precipitate the Ca 3 (PO 4 ) 2 , except for the first rough
trial, do not apply heat till the precipitation of uranyl
phosphate is nearly completed. Label the solution i cu. cm.
= x grm. P 2 O 5 [in Ca 3 (PO 4 )J. If you have used pure salts,
both your standard solutions of phosphates [NaNH 4 HPO 4
and Ca 3 (PO 4 ) ; ,] contain 0.005 S rm ' ^2^5 P er cu ic centi-
metre. Do you find this borne out by the behaviour of
your uranyl solution?

3. Estimate the P 2 O 5 in the Bone Ash. Since this sub-
stance contains calcium (and magnesium) phosphate, but
no iron nor aluminium, make a solution of 2 to 3 grm. as
directed in i, and make the estimation, observing the pre-
cautions given in 2.

Miscellaneous Exercises

Exercise 104. Find the proportion of Na. 2 D in
a sample of borax crystals. How does it agree with
the formula of prismatic borax Na 2 B 4 Y , 10 H>0?

Directions. i. Dissolve up about 5 grm. to 250 cu. cm.,,

and titrate with H 2 SO 4 ,

2. Indicator: methyl orange. The liberated boric acid
does not interfere with this indicator.

Na :! R/X + H 2 S0 4 + 5 H,O = Na,SO 4 + 4 H 3 BO, r


Exercise 105. Estimate the percentage weight of
B 2 3 in the same specimen of borax.

Directions. i. Use same solution as for Ex. 104, and
titrate each portion taken with H 2 SO 4 till methyl orange
shows liquid to be just acid. Boil to expel any CO 2 pre-


sent; cool, and titrate back with NaOH till exact


neutrality is obtained, as shown by appearance of pure
yellow colour. The only free acid now present is boric
acid, which does not affect methyl orange.

2. Add at least half the volume of glycerin, one or two

drops of phenolphthalein, and titrate with - NaOH till
red colour appears.

Note. R. T. Thomson advocates the use of glycerin to make
the phenolphthalein sensitive to boric acid.

Boric acid reacts with NaOH as shown

NaOH + H 3 BO 3 = NaH 2 BO 3 + H 2 O,
so that 80 grm. NaOH = 70 grm. B 2 O 3 .

Exercise 106. Estimate the percentage of CaC0 3
in some impure calcite, assuming that no other
carbonate is present.

Directions. See Ex. 40. Neglect the slight amount of
insoluble matter.

Exercise 107. Determine the solubility of oxalic
acid in water at the temperature of the laboratory,
which should be stated.

Directions. i. Saturate about 50 cu. cm. of warm water
with pure powdered acid, and allow to cool to temperature
of laboratory. Some crystals should remain undissolved.

2. Make up 25 cu. cm. of the clear solution to 250 cu. cm. ,

and titrate with alkali or KMnO..
10 10

Note. i 1. should contain about 9.3 grm. of (COOH) 2 at 15

1 2 3 4 5 6 8 10 11

Online LibraryWilliam Thomas BooneA complete course of volumetric analysis for middle and higher forms of schools → online text (page 8 of 11)