ammonic chloride in water, and diluting to 100 c.c. Ten c.c. of this are
taken and diluted to 1 litre. One c.c. contains 0.01 milligram of
ammonia (NH_{3}).
In working, the solution containing the ammonia is diluted to a definite
volume, and to such an extent that 50 c.c. of it shall not contain more
than 0.02 or 0.03 milligram of ammonia. Fifty c.c. of it are transferred
to a Nessler glass and mixed with 2 c.c. of Nessler's reagent. The
colour is noted, and an estimate made as to the amount of ammonia it
indicates. A measured quantity of the standard ammonia, judged to
contain about as much ammonia as that in the assay, is then put into
another Nessler glass. It is diluted to 50 c.c. with water, and mixed
with 2 c.c. of "Nessler." After standing a minute or two, the colours in
the two glasses are compared. If the tints are equal, the assay is
finished; but if the standard is weaker or stronger than the assay,
another standard, containing more or less ammonia, as the case may be,
must be prepared and compared with the assay. Two such experiments will
generally be sufficient; but, if not, a third must be made. The addition
of more standard ammonia to the solution to which the "Nessler" has
already been added does not give a satisfactory result.
When the ammonia in 50 c.c. has been determined, that in the whole
solution is ascertained by a suitable multiplication. By 10, for
example, if the bulk was 500 c.c., or by 20 if it was a litre.
Distilled water is used throughout. It must be free from ammonia; and is
best prepared by distilling an ammonia-free spring water.
FOOTNOTES:
[90] Al_{2}Cl_{6} + 3Na_{2}S_{2}O_{3} + 3H_{2}O = Al_{2}(HO)_{6} + 6NaCl
+ 3S + 3SO_{2}
[91] 3BeO,Al_{2}O_{3},6SiO_{2}
[92] CaC_{2}O_{4} = CaCO_{3}+CO.
[93] Resolved into two with a powerful spectroscope.
[94] Ammonium compounds are frequently produced when dissolving metals
in nitric acid; or when nitrates are heated in the presence of the
metals.
PART III--NON-METALS.
CHAPTER XV.
OXYGEN AND OXIDES.--THE HALOGENS.
OXYGEN.
Oxygen occurs in nature in the free state, forming 23 per cent. by
weight, or 21 per cent. by volume of the atmosphere; but, since it is a
gas, its presence is easily overlooked and its importance
underestimated. Except in the examination of furnace-gases, &c., the
assayer is not often called upon to determine its quantity, but it forms
one of his most useful reagents, and there are many cases where he
cannot afford to disregard its presence. It occurs not only in the air,
but also dissolved in water; ordinary waters containing on an average
0.00085 per cent. by weight, or 0.85 parts per 100,000.
Chemically, it is characterised by its power of combining, especially at
high temperatures, with the other elements, forming an important class
of compounds called oxides. This combination, when rapid, is accompanied
by the evolution of light and heat; hence oxygen is generally called the
supporter of combustion. This property is taken advantage of in the
operation of calcining, scorifying, cupelling, &c. The importance of a
free access of air in all such work is seen when it is remembered that 1
litre of air contains 0.2975 gram of oxygen, and this quantity will only
oxidise 0.1115 gram of carbon, 0.2975 gram of sulphur, or 3.849 grams of
lead.
Oxidation takes place at the ordinary temperature with many substances.
Examples of such action are seen in the weathering of pyrites, rusting
of iron, and (in the assay office) the weakening of solutions of many
reducing agents.
For methods of determining the percentage of oxygen in gases, for
technical purposes, the student is referred to Winkler & Lunge's
"Technical Gas Analysis."
OXIDES.
Oxides are abundant in nature, almost all the commonly occurring bodies
being oxidised. Water (H_{2}O) contains 88.8 per cent. of oxygen;
silica, lime, alumina, magnesia, and the other earths are oxides, and
the oxides of the heavier metals are in many cases important ores; as,
for example, cassiterite (SnO_{2}), hæmatite (Fe_{2}O_{3}), magnetite
(Fe_{3}O_{4}), and pyrolusite (MnO_{2}). In fact, the last-named mineral
owes its value to the excess of oxygen it contains, and may be regarded
as an ore of oxygen rather than of manganese.
Most of the metals, when heated to redness in contact with air, lose
their metallic lustre and become coated with, or (if the heating be
prolonged) altogether converted into, oxide. This oxide was formerly
termed a "calx," and has long been known to weigh more than the metal
from which it was obtained. For example, one part by weight of tin
becomes, on calcining, 1.271 parts of oxide (putty powder). The student
will do well to try the following experiments:--Take 20 grams of tin and
heat them in a muffle on a scorifier, scraping back the dross as it
forms, and continuing the operation until the whole of the metal is
burnt to a white powder and ceases to increase in weight.[95] Take care
to avoid loss, and, when cold, weigh the oxide formed. The oxide should
weigh 25.42 grams, which increase in weight is due to the oxygen
absorbed from the air and combined with the metal. It can be calculated
from this experiment (if there has been no loss) that oxide of tin
contains 21.33 per cent. of oxygen and 78.67 per cent. of tin. Oxidation
is performed with greater convenience by wet methods, using reagents,
such as nitric acid, which contain a large proportion of oxygen loosely
held. Such reagents are termed oxidising agents. Besides nitric acid,
permanganate of potash, bichromate of potash, and peroxide of hydrogen
are largely used for this purpose. One c.c. of nitric acid contains as
much oxygen as 2.56 litres of air, and the greater part of this is
available for oxidising purposes. Try the following experiment:--Take 2
grams of tin and cover in a weighed Berlin dish with 20 c.c. of dilute
nitric acid, heat till decomposed, evaporate to dryness, ignite, and
weigh. The 2 grams of tin should yield 2.542 grams of oxide. The
increase in weight will be proportionally the same as in the previous
experiment by calcination, and is due to oxygen, which in this case has
been derived from the nitric acid.
The percentage of oxygen in this oxide of tin (or in any of the oxides
of the heavier metals) may be directly determined by heating such oxides
in a current of hydrogen, and collecting and weighing the water formed.
It is found by experiment that 88.86 parts by weight of oxygen,
combining with 11.14 parts of hydrogen, form 100 parts of water; so that
from the weight of water formed it is easy to calculate the amount of
oxygen the oxide contained.
[Illustration: FIG. 62.]
Take 1 gram of the dried and powdered oxide and place it in a warm dry
combustion tube. Place the tube in a furnace, and connect at one end
with a hydrogen apparatus provided with a sulphuric acid bulb for drying
the gas, and at the other with a weighed sulphuric acid tube for
collecting the water formed. The apparatus required is shown in fig. 62.
Pass hydrogen through the apparatus, and, when the air has been cleared
out, light the furnace. Continue the heat and current of hydrogen for
half an hour (or longer, if necessary). Allow to cool. Draw a current of
dry air through the weighed tube. Weigh. The increase in weight gives
the amount of water formed, and this, multiplied by 0.8886, gives the
weight of the oxygen. The percentage of oxygen thus determined should be
compared with that got by the oxidation of the metal. It will be
practically the same. The following results can be taken as examples:--
Twenty grams of tin, calcined as described, gave 25.37 grams of oxide.
Two grams of tin, oxidised with nitric acid and ignited, gave 2.551
grams of oxide.
One gram of the oxide of tin, on reduction in a current of hydrogen,
gave 0.2360 gram of water (equivalent to 0.2098 gram of oxygen), and
left 0.7900 gram of metal.
Ten grams of ferrous sulphate gave, on strong ignition, 2.898 grams of
ferric oxide (Fe_{2}O_{3})[96] instead of 2.877.
The student should similarly determine the percentage of oxygen in
oxides of copper and iron. The former oxide may be prepared by
dissolving 5 grams of copper in 50 c.c. of dilute nitric acid,
evaporating to dryness, and strongly igniting the residue. The oxide of
iron may be made by weighing up 10 grams of powdered ferrous sulphate
(= to 2.014 grams of iron) and heating, at first gently, to drive off the
water, and then at a red heat, until completely decomposed. The weight
of oxide, in each case, should be determined; and the percentage of
oxygen calculated. Compare the figures arrived at with those calculated
from the formula of the oxides, CuO and Fe_{2}O_{3}.
It would be found in a more extended series of experiments that the same
metal will, under certain conditions, form two or more oxides differing
among themselves in the amount of oxygen they contain. These oxides are
distinguished from one another by such names as "higher" and "lower
oxides," "peroxides," "protoxides," "dioxides," &c.
The oxides may be conveniently classified under three heads:--
(1) _Those that are reduced to metal by heat alone_, such as the oxides
of mercury, silver, platinum, gold, &c.;
(2) _Those which are reduced by hydrogen at a red heat_, which includes
the oxides of the heavy metals;
(3) _Those which are not reduced by these means_, good examples of which
are silica, alumina, the alkalies, and the alkaline earths.
Another important classification is into acid, basic and neutral oxides.
The oxides of the non-metallic elements, such as sulphur, carbon,
phosphorus, &c., are, as a rule, acid; and the more oxygen they contain,
the more distinctly acid they are. The oxides of the metals are nearly
all basic; and, as a rule, the less oxygen they contain, the more
distinctly basic they are.
The basic oxides, which are soluble in acids, give rise to the formation
of salts when dissolved therein. During the solution, water is formed,
but no gas is evolved. The oxide dissolved in each case neutralizes an
equivalent of the acid used for solution.[97] The basic properties of
many of these can be taken advantage of for their determination. This is
done in the case of soda, potash, lime, &c., by finding the quantity of
acid required to neutralize a given weight of the substance.
There are some oxides which, under certain conditions, are acid to one
substance (a stronger base) and basic to another (a stronger acid). For
example, the oxides of lead and of tin, as also alumina, dissolve in
caustic soda, acting as acids; whilst, on the other hand, they combine
with sulphuric or hydrochloric acid, playing the part of bases.
The oxides known as "earths," when ignited, are many of them insoluble
in acids, although easily dissolved before ignition.
It is common in complete analyses of minerals to meet with cases in
which the sum total of the elements found falls short of the amount of
ore taken; and here oxygen must be looked for. For example, this occurs
in the case of a mixture of pyrites with oxide of iron, or in a mixture
of sulphides and sulphates. The state in which the elements are present,
and the percentage (say of sulphides and sulphates) can in many cases be
determined; but this is not always required. When the difference between
the sum total and the elements found is small, it is reported as "oxygen
and loss." When, however, it is considerable, the oxygen may be reported
as such; and its amount be either determined directly in the way already
described, or calculated from the best determination that can be made of
the relative amounts of oxides, sulphides, sulphates, &c., present. Such
cases require a careful qualitative analysis to find out that the
substance is present; and then the separation of each constituent is
made as strictly as possible. These remarks apply especially to ores of
the heavy metals. The separation of the constituents is effected with
suitable solvents applied in proper order. The soluble sulphates, for
example, are extracted with water; the oxides by the dilute acids or
alkalies in which they are known to be soluble. The oxygen in the
sulphates and oxides thus obtained is estimated by determining the
sulphur and metals in the solutions, and calculating the amount of
oxygen with which they combine. The metals of the earths and alkalies
are almost invariably present as oxides, and are reported as such;
except it is known that they are present in some other form, such as
fluoride or chloride. Thus, silica, alumina, lime, water, &c., appear in
an analysis; even in those cases where "oxygen and loss" is also
mentioned. As an example of such a report, take the following analysis
of Spanish pyrites:--
Sulphur 49.00
Iron 43.55
Copper 3.20
Arsenic 0.47
Lead 0.93
Zinc 0.35
Lime 0.10
Silica, &c. 0.63
Water 0.70
Oxygen and loss 1.07
-----
100.00
The following example will illustrate the mode of calculating and
reporting. A mineral, occurring as blue crystals soluble in water, and
found on testing to be a mixed sulphate of iron and copper, gave on
analysis the following results:--
Water 44.51 per cent.
Sulphuric oxide 28.82 "
Copper 8.44 "
Ferrous iron 11.81 "
Ferric iron 0.38 "
Zinc 0.28 "
-----
94.24
There is here a deficiency of 5.76 per cent. due to oxygen. Nothing else
could be found, and it is known that in the sulphates the metals exist
as oxides. By multiplying the weight of the copper by 1.252, the weight
of copper oxide (CuO) will be ascertained; in this case it equals 10.57
per cent. The ferrous iron multiplied by 1.286 will give the ferrous
oxide (FeO); in this case 15.19 per cent. The ferric iron multiplied by
1.428 will give the ferric oxide (Fe_{2}O_{3}); in this case 0.54 per
cent. The zinc multiplied by 1.246 will give the zinc oxide (ZnO); in
this case it equals 0.35 per cent. The analysis will be reported as--
Water 44.51
Sulphuric oxide 28.82
Copper oxide 10.57 equal to copper 8.44%
Ferrous oxide 15.19
Ferric oxide 0.54
Zinc oxide 0.35
-----
99.98
The following (A) is an analysis of a sample of South American copper
ore, which will serve as a further illustration. The analysis showed the
presence of 6.89 per cent. of ferrous oxide, and some oxide of copper.
The analysis (B) is that of an ore from the same mine after an imperfect
roasting. It will be seen that the carbonates have been converted into
sulphates. If the total sulphur simply had been determined, and the
sulphate overlooked, the "oxygen and loss" would have been 5.65 per
cent., an amount which would obviously require an explanation.
A. B.
Water 0.25 0.59
Organic matter 0.54 --
Sulphur 29.50 21.33
Copper 10.92 9.80 {Copper 9.57
{Copper oxide 0.28
Iron 32.09 39.73 {Iron 34.32
{Ferric oxide 7.73
Lead 0.35 0.12
Zinc 0.86 0.69
Cobalt 0.06 0.11
Lime 5.25 7.69
Magnesia 2.33 2.55
Sulphuric oxide 1.00 5.30
Carbon dioxide 8.87 --
"Insoluble silicates" 5.12 8.38
Oxygen and loss 2.86 2.47
----- Potash 0.15
100.00 Soda 1.09
-----
100.00
WATER.
Water occurs in minerals in two forms, free and combined. The term
"moisture" ought, strictly, to be limited to the first, although, as has
already been explained, it is more convenient in assaying to apply the
term to all water which is driven off on drying at 100° C. The combined
water is really a part of the mineral itself, although it may be driven
off at a high temperature, which varies with the base. In some cases a
prolonged red heat is required; whilst with crystallised salts it is
sometimes given off at the ordinary temperatures. This latter
phenomenon, known as efflorescence, is mostly confined to artificial
salts.
The determination of the combined water may often be made by simply
igniting the substance from which the moisture has been removed. The
quantity of water may be determined, either indirectly by the loss, or
directly by collecting it in a calcium chloride tube, and weighing. In
some cases, in which the loss on ignition does not give simply the
proportion of combined water, it can be seen from the analysis to what
else the loss is due; and, after a proper deduction, the amount of water
can be estimated. For example, 1 gram of crystallised iron sulphate was
found to contain on analysis 0.2877 gram of sulphuric oxide; and on
igniting another gram, 0.2877 gram of ferric oxide was left. As the salt
is known to be made up of ferrous oxide, sulphuric oxide, and combined
water, the combined water can be thus calculated: 0.2877 gram of ferric
oxide is equal to 0.2589 gram of ferrous oxide,[98] and consequently,
the loss on ignition has been diminished by 0.0288 gram, which is the
weight of oxygen absorbed by the ferrous oxide during calcining. The
loss on ignition was 0.7123 gram, to which must be added 0.0288 gram;
hence 0.7411 gram is the weight of the combined sulphuric oxide and
water present. Deducting the weight of sulphuric oxide found, 0.2877
gram, there is left for combined water 0.4534 gram. The composition of 1
gram of the dry salt is then:--
Water 0.4534
Sulphuric oxide 0.2877
Ferrous oxide 0.2589
------
1.0000
The following is another example:--A sample of malachite lost on
ignition 28.47 per cent., leaving a residue which was found on analysis
to be made up of oxide of copper (equal to 70.16 per cent. on the
mineral), and silica and oxide of iron (equal to 1.37 per cent.). Carbon
dioxide and water (but nothing else) was found to be present, and the
carbon dioxide amounted to 19.64 per cent.; deducting this from the loss
on ignition, we have 8.82 as the percentage of water present. The
analysis was then reported as follows:--
Cupric oxide 70.16 equal to 56.0% copper.
Silica and ferric oxide 1.37
Carbon dioxide 19.64
Water 8.82
-----
99.99
[Illustration: FIG. 63.]
~Direct Determination of Combined Water.~--Transfer about 3 grams of the
substance to a piece of combustion tube (8 or 10 inches long), attached
(as in fig. 63) at one end to a ~U~-tube containing sulphuric acid, and
at the other end to a calcium chloride tube. The last is weighed
previous to the determination. The tube should be warmed to ensure
complete dryness, and must be free from a misty appearance. Aspirate a
current of air through the apparatus, heat the mineral by means of a
Bunsen burner, cautiously at first, and afterwards to redness (if
necessary). The water is driven off and condenses in the calcium
chloride tube, which is afterwards cooled and weighed. The increase in
weight is due to the water. If the substance gives off acid products on
heating, it is previously mixed with some dry oxide of lead or pure
calcined magnesia.
EXAMINATION OF WATERS.
The assayer is occasionally called on to test water for the purpose of
ascertaining the nature and quantity of the salts contained in it, and
whether it is or is not fit for technical and drinking purposes.
In mineral districts the water is generally of exceptional character,
being more or less charged, not only with earthy salts, but also
frequently with those of the metals. Distilled water is only used by
assayers in certain exceptional cases, so that by many it would be
classed among the rarer oxides. Water of ordinary purity will do for
most purposes, but the nature and quantity of the impurities must be
known.
The following determinations are of chief importance:--
~Total Solids at 100° C.~--Where simply the amount is required, take 100
c.c. and evaporate on the water-bath in a weighed dish; then dry in the
water-oven, and weigh.
~Total Solids Ignited.~--The above residue is very gently ignited
(keeping the heat well below redness), and again weighed. A larger loss
than 4 or 5 parts per 100,000 on the water requires an explanation.
~Chlorine.~--Take 100 c.c. of the water in a porcelain dish, add 2 c.c.
of a 5 per cent. solution of neutral potassic chromate, and titrate with
a neutral standard solution of nitrate of silver, made by dissolving
4.789 grams of crystallised silver nitrate in distilled water, and
diluting to 1 litre. The addition of the nitrate of silver is continued
until the yellow of the solution assumes a reddish tint. The reaction is
very sharp. Each c.c. of nitrate of silver used is equal to 1 part by
weight of chlorine in 100,000 of water. At inland places this rarely
amounts to more than 1 in 100,000; but near the sea it may amount to 3
or 5. More than this requires explanation, and generally indicates
sewage pollution.
~Nitric Pentoxide (N_{2}O_{5}).~--It is more generally reported under
the heading, "nitrogen as nitrates." Take 250 c.c. of the water and
evaporate to 2 or 3 c.c.; acidulate with a few drops of dilute sulphuric
acid, and transfer to a nitrometer (using strong sulphuric acid to wash
in the last traces). The sulphuric acid must be added to at least twice
the bulk of the liquid. Shake up with mercury. The mercury rapidly
flours, and nitric oxide is given off (if any nitrate is present). The
volume of the nitric oxide (corrected to normal temperature and
pressure), multiplied by 0.25, gives the parts of nitrogen per 100,000;
or, multiplied by 0.965, will give the nitric pentoxide in parts per
100,000. In well and spring waters the nitrogen may amount to 0.3 or 0.4
parts per 100,000; or in richly cultivated districts 0.7 or 0.8 parts
per 100,000. An excess of nitrates is a suspicious feature, and is
generally due to previous contamination.
~Ammonia.~--Take 500 c.c. of the water and place them in a retort
connected with a Liebig's condenser. Add a drop or two of a solution of
carbonate of soda and distil over 100 c.c.; collect another 50 c.c.
separately. Determine the ammonia in the distillate colorimetrically
(with Nessler's solution, as described under _Ammonia_) and compare with
a standard solution of ammonic chloride containing 0.0315 gram of
ammonic chloride in 1 litre of water. One c.c. contains 0.01 milligram
of ammonia. The second distillate will show little, if any, ammonia in
ordinary cases. The amounts found in both distillates are added
together, and expressed in parts per 100,000.
Waters (other than rain and tank waters) which contain more than 0.003
per 100,000 are suspicious.
~Organic Matter.~--The organic matter cannot be determined directly; but
for ordinary purposes it may be measured by the amount of permanganate
of potassium which it reduces, or by the amount of ammonia which it
evolves on boiling with an alkaline permanganate of potassium solution.
A. _Albuminoid Ammonia._--To the residue left after distilling the
ammonia add 50 c.c. of a solution made by dissolving 200 grams of potash
and 8 grams of potassium permanganate in 1100 c.c. of water, and rapidly
boiling till the volume is reduced to 1 litre (this should be kept in a
well stoppered bottle, and be occasionally tested to see that it is free
from ammonia). Continue the distillation, collecting 50 c.c. at a time,
until the distillate is free from ammonia. Three or four fractions are
generally sufficient. Determine the ammonia colorimetrically as before.
If the total albuminoid ammonia does not exceed 0.005 in 100,000, the
water may be regarded as clean as regards organic matter; if it amounts
to more than 0.015, it is dirty.
B. _Oxygen Consumed._--A standard solution of permanganate of potash is
made by dissolving 0.395 gram of the salt in water and diluting to 1
litre. Each c.c. equals 0.1 milligram of available oxygen. The following
are also required:--1. A solution of sodium hyposulphite containing 1
gram of the salt (Na_{2}S_{2}O_{3}.5H_{2}O) in 1 litre of water. 2.
Dilute sulphuric acid, made by adding one part of the acid to three of
water, and titrating with the permanganate solution till a faint pink
persists after warming for several hours. 3. Starch paste. 4. Potassium
iodide solution.
Take 250 c.c. of the water in a stoppered bottle, add 10 c.c. of
sulphuric acid and 10 c.c. of the permanganate, and allow to stand in a
warm place for four hours. Then add a few drops of the solution of
potassium iodide, and titrate the liberated iodine with "hypo," using
starch paste towards the end as an indicator. To standardise the
hyposulphite, take 250 c.c. of water and 10 c.c. of sulphuric acid, and
a few drops of potassium iodide; then run in 10 c.c. of the
"permanganate" solution, and again titrate; about 30 c.c. of the "hypo"
will be used. The difference in the two titrations, divided by the last
and multiplied by 10, will give the c.c. of permanganate solution used
in oxidising the organic matter in the 250 c.c. of water. Each c.c.
represents 0.04 parts of oxygen in 100,000.
~Metals.~--These may for the most part be estimated colorimetrically.
~Lead.~--Take 100 c.c. of the water in a Nessler tube, and add 10 c.c.
of sulphuretted hydrogen water, and compare the tint, if any, against a
standard lead solution, as described under _Colorimetric Lead_. Report
in parts per 100,000.
~Copper.~--Proceed as with the last-mentioned metal; but, if lead is
also present, boil down 500 c.c. to about 50 c.c., then add ammonia,
filter, and estimate the copper in the blue solution, as described under
_Colorimetric Copper_.
~Iron.~--Take 50 c.c., or a smaller quantity (if necessary), dilute up
to the mark with distilled water, and determine with potassium
sulphocyanate, as described under _Colorimetric Iron_.
~Zinc.~--Zinc is the only other metal likely to be present; and, since
it cannot be determined colorimetrically, it must be separately
estimated during the examination of the "total solids."
~Examination of "Total Solids."~--Evaporate 500 c.c. to dryness with a
drop or two of hydrochloric acid. Take up with hydrochloric acid,
filter, ignite, and weigh the residue as "silica." To the filtrate add a
little ammonic chloride and ammonia, boil and filter, ignite, and weigh
the precipitate as "oxide of iron and alumina." Collect the filtrate in
a small flask, add a few drops of ammonium sulphide or pass sulphuretted
hydrogen, cork the flask, and allow to stand overnight; filter, wash,
and determine the zinc gravimetrically as oxide of zinc. If copper or
lead were present, they should have been previously removed with
sulphuretted hydrogen in the acid solution. To the filtrate add ammonic
oxalate and ammonia, boil for some time, allow to stand, filter, wash,
ignite, and weigh as "lime." Evaporate the filtrate with nitric acid,
and ignite. Take up with a few drops of dilute hydrochloric acid, add
baric hydrate in excess, evaporate, and extract with water. The residue
contains the magnesia; boil with dilute sulphuric acid, filter,
precipitate it with phosphate of soda and ammonia, and weigh as
pyrophosphate. The aqueous extract contains the alkalies with the excess
of barium. Add sulphuric acid in slight excess, filter, evaporate, and
ignite strongly. The residue consists of the sulphates of the alkalies
(which are separately determined, as described under _Potash_).
~Sulphuric Oxide (SO_{3}).~--Take 200 c.c. and boil to a small bulk with
a little hydrochloric acid, filter (if necessary), add baric chloride
solution in slight excess to the hot solution, filter, ignite, and weigh
as baric sulphate.
~Carbon Dioxide (free).~--Carbon dioxide exists in waters in two forms,
free and combined. The latter generally occurs as bicarbonate, although
on analysis it is more convenient to consider it as carbonate, and to
count the excess of carbon dioxide with the free. The method is as
follows:--To determine the free carbon dioxide, take 100 c.c. of the
water, place them in a flask with 3 c.c. of a strong solution of calcium
chloride and 2 c.c. of a solution of ammonic chloride, next add 50 c.c.
of lime-water. The strength of the lime-water must be known. Make up to
200 c.c. with distilled water, stop the flask, and allow the precipitate
to settle. Take out 100 c.c. of the clear solution with a pipette, and
titrate with the standard solution of acid.[99] The number of c.c.
required, multiplied by two, and deducted from that required for the 50
c.c. of lime-water, and then multiplied by 0.0045, will give the carbon
dioxide present other than as normal carbonates.
~Carbon Dioxide combined~ as normal carbonate.--100 c.c. of the water
are tinted with phenacetolin or lacmoid; then heated to near boiling,
and titrated with standard acid. The number of c.c. used, multiplied by
0.0045, will give the weight in grams of the combined carbon dioxide.
~Free Acid.~--In some waters (especially those from mining districts)
there will be no carbonates. On the contrary, there may be free mineral
acid or acid salts. In these cases it is necessary to determine the
amount of acid (other than carbon dioxide) present in excess of that
required to form normal salts. This is done in the following way:--Make
an ammoniacal copper solution by taking 13 grams of copper sulphate
(CuSO_{4}.5H_{2}O), dissolving in water, adding solution of ammonia
until the precipitate first formed has nearly dissolved, and diluting to
1 litre. Allow to settle, and decant off the clear liquid. The strength
of this solution is determined by titrating against 10 or 20 c.c. of the
standard solution of sulphuric acid (100 c.c. = 1 gram H_{2}SO_{4}). The
finishing point is reached as soon as the solution becomes turbid from
precipitated cupric hydrate. At first, as each drop falls into the acid
solution, the ammonia and cupric hydrate combine with the free acid to
form ammonic and cupric sulphates; but as soon as the free acid is used
up, the ammonia in the next drop not only precipitates an equivalent of
cupric hydrate from the solution, but also throws down that carried by
itself. This method is applicable in the presence of metallic sulphates
_other than ferric_. The standardising and titration should be made
under the same conditions. Since sulphuric acid and sulphates are
predominant in waters of this kind, it is most convenient to report the
acidity of the water as equivalent to so much sulphuric acid.
~Dissolved Oxygen.~--For the gasometric method of analysing for
dissolved oxygen, and for the Schützenberger's volumetric method, the
student is referred to Sutton's "Volumetric Analysis." The following is
an easy method of estimating the free oxygen in a water:--Take 20 c.c.
of a stannous chloride solution (about 20 grams of the salt with 10 c.c.
of hydrochloric acid to the litre); add 10 c.c. of hydrochloric acid,
and titrate in an atmosphere of carbon dioxide with standard
permanganate of potassium solution (made by dissolving 1.975 gram of the
salt in 1 litre of water: 1 c.c. equals 0.5 milligram of oxygen). A
similar titration is made with the addition of 100 c.c. of the water to
be tested. Less permanganate will be required in the second titration,
according to the amount of oxygen in the water; and the difference,
multiplied by 0.5, will give the weight of the oxygen in milligrams.
Small quantities of nitrates do not interfere.
In REPORTING the results of the analysis, it is customary to combine the
acids and bases found on some such principle as the following:--The
sulphuric oxide is calculated as combined with the potash, and reported
as potassic sulphate (K_{2}SO_{4}); the balance of the sulphuric oxide
is then apportioned to the soda, and reported as sulphate of soda
(Na_{2}SO_{4}); if any is still left, it is reported as calcium sulphate
(CaSO_{4}), and after that as magnesic sulphate (MgSO_{4}). When the
sulphuric oxide has been satisfied, the chlorine is distributed, taking
the bases in the same order, then the nitric pentoxide, and lastly the
carbon dioxide. But any method for thus combining the bases and acids
must be arbitrary and inaccurate. It is extremely improbable that any
simple statement can represent the manner in which the bases and acids
are distributed whilst in solution; and, since different chemists are
not agreed as to any one system, it is better to give up the attempt,
and simply state the results of the analysis. This has only one
inconvenience. The bases are represented as oxides; and, since some of
them are present as chlorides, the sum total of the analysis will be in
excess of the actual amount present by the weight of the oxygen
equivalent to the chlorine present as chloride. The following is an
example of such a statement:--
Parts per 100,000.
Total solids, dried at 100° C. 28.73
Chlorine 1.70
Nitrogen as nitrate 0.03
Ammonia 0.001
Albuminoid ammonia 0.004
"Oxygen consumed" in 4 hours 0.01
The solids were made up as under:--
Per 100,000 of the Water.
Potash 0.38
Soda 2.01
Magnesia 1.44
Lime 10.55
Ferric oxide 0.01
Silica 0.30
Sulphuric oxide 3.69
Nitrogen pentoxide 0.11
Carbon dioxide 8.38
Chlorine 1.70
Volatile and organic matter 0.66
-----
29.23
Less oxygen equivalent to chlorine found 0.39
-----
28.84
For the preparation of distilled water, the apparatus shown in fig. 64
is convenient for laboratory use. It consists of a copper retort heated
by a ring gas-burner, and connected with a worm-condenser.
[Illustration: FIG. 64]
PRACTICAL EXERCISE.
A mineral, on analysis, gave the following results:--Water, 44.94 per
cent.; sulphuric oxide, 28.72 per cent.; ferrous iron, 13.92 per cent.;
ferric iron, 0.35 per cent.; copper, 6.1 per cent. The mineral was
soluble in water, and showed nothing else on testing. How would you
report the analysis? Calculate the formula for the salt.
THE HALOGENS.
There is a group of closely allied elements to which the name halogen
(salt-producer) has been given. It comprises chlorine, bromine, iodine,
and fluorine. These elements combine directly with metals, forming as
many series of salts (chlorides, bromides, iodides, and fluorides),
corresponding to the respective oxides, but differing in their formulæ
by having two atoms of the halogen in the place of one atom of oxygen.
For example, ferrous oxide is FeO and ferrous chloride is FeCl_{2}, and,
again, ferric oxide is Fe_{2}O_{3}, whilst ferric chloride is
Fe_{2}Cl_{6}. These salts differ from the carbonates, nitrates, &c., in
containing no oxygen. Consequently, it is incorrect to speak of such
compounds as chloride of potash, fluoride of lime, &c., since potash and
lime are oxides. It is important to bear this in mind in reporting
analyses in which determinations have been made, say, of chlorine,
magnesia, and potash, or of fluorine, silica, and alumina. It is
necessary in all such cases to deduct from the total an amount of oxygen
equivalent to the halogen found, except, of course, where the base has
been determined and recorded as metal. Compounds containing oxides and
fluorides, &c., do not lend themselves to the method of determining the
halogen by difference. For example, topaz, which, according to Dana, has
the formula Al_{2}SiO_{4}F_{2}, would yield in the ordinary course of
analysis--
Alumina 55.4%
Silica 32.6
Fluorine 20.6
-----
108.6
The oxygen equivalent to 20.6 per cent. fluorine may be found by
multiplying the percentage of fluorine by 0.421; it is 8.7 per cent.,
and must be deducted. The analysis would then be reported thus:--
Alumina 55.4%
Silica 32.6
Fluorine 20.6
-----
108.6
Less oxygen equivalent to fluorine 8.7
-----
99.9
Take as an illustration the following actual analysis by F.W. Clarke and
J.S. Diller:--
Alumina 57.38%
Silica 31.92
Fluorine 16.99
Potash 0.15
Soda 1.33
Water 0.20
------
107.97
Deduct oxygen equivalent 7.16
------
100.81
In calculating the factor for the "oxygen equivalent," divide the weight
of one atom of oxygen (16) by the weight of two atoms of the halogen;
for example, with chlorine it would be 16/71 or 0.2253; with bromine,
16/160 or 0.1000; with iodine, 16/254 or 0.063; and with fluorine, 16/38
or 0.421.
CHLORINE AND CHLORIDES.
Chlorine occurs in nature chiefly combined with sodium, as halite or
rock salt (NaCl). With potassium it forms sylvine (KCl), and, together
with magnesium, carnallite (KCl.MgCl_{2}.6H_{2}O). Of the metalliferous
minerals containing chlorine, kerargyrite, or horn silver (AgCl), and
atacamite, an oxychloride of copper (CuCl_{2}.3Cu(HO)_{2}.) are the most
important. Apatite (phosphate of lime) and pyromorphite (phosphate of
lead) contain a considerable amount of it. Chlorine is a gas of a
greenish colour, possessing a characteristic odour, and moderately
soluble in water. It does not occur native, and is generally prepared by
the action of an oxidising agent on hydrochloric acid. It combines
directly with metals at the ordinary temperature (even with platinum and
gold), forming chlorides, which (except in the case of silver) are
soluble.
It is important in metallurgy, because of the extensive use of it in
extracting gold by "chloridising" processes. It is also used in refining
gold.
~Detection.~--Compounds containing the oxides of chlorine are not found
in nature, because of the readiness with which they lose oxygen. By
reduction they yield a chloride; the form in which chlorine is met with
in minerals. In testing, the compound supposed to contain a chloride is
boiled with water, or, in some cases, dilute nitric acid. To the clear
solution containing nitric acid a few drops of nitrate of silver
solution are added. If, on shaking, a white curdy precipitate, soluble
in ammonia, separates out, it is sufficiently satisfactory evidence of
the presence of chlorides.
~Solution and Separation.~--The chlorides are generally soluble in
water, and are got into solution by extracting with warm dilute nitric
acid. Or, if insoluble, the substance is fused with carbonate of soda,
extracted with water, and the filtrate acidified with nitric acid. For
the determination, it is not necessary to obtain the solution of the
chloride free from other acids or metals. If tin, antimony, mercury, or
platinum is present, it is best to separate by means of sulphuretted
hydrogen. The chloride is determined in the solution after removal of
the excess of the gas. Where traces of chlorides are being looked for, a
blank experiment is made to determine the quantity introduced with the
reagents. One hundred c.c. of ordinary water contains from 1 to 3
milligrams of chlorine. On the addition of nitrate of silver to the
nitric acid solution, chloride of silver separates out. This is free
from other substances, except, perhaps, bromide and iodide.
GRAVIMETRIC DETERMINATION.
Freely mix the solution containing the chloride with dilute nitric acid,
filter (if necessary), and treat with nitrate of silver. Heat nearly to
boiling, and, when the precipitate has settled, filter, and wash with
hot distilled water. Dry, and transfer to a weighed Berlin crucible.
Burn the filter-paper separately, and convert any reduced silver into
chloride by alternate treatment with drops of nitric and of hydrochloric
acid. Add the main portion to this, and heat cautiously till the edges
of the mass show signs of fusing (about 260°). Cool in the desiccator
and weigh. The substance is chloride of silver (AgCl), and contains
24.73 per cent. of chlorine.
The precipitated chloride is filtered and washed as soon as possible
after settling, since on exposure to light it becomes purple, and loses
a small amount of chlorine.
VOLUMETRIC METHOD.
There are several volumetric methods; but that based on the
precipitation of silver chloride in neutral solution, by means of a
standard solution of silver nitrate (using potassium chromate as
indicator), is preferred. Silver chromate is a red-coloured salt; and,
when silver nitrate is added to a solution containing both chloride and
chromate, the development of the red colour marks off sharply the point
at which the chloride is used up. Silver chromate is decomposed and
consequently decolorised by solution of any chloride. The solution for
this method must be neutral, since free acid prevents the formation of
the red silver chromate. If not already neutral, it is neutralised by
titrating cautiously with a solution of soda. In a neutral solution,
other substances (such as phosphates and arsenates) also yield a
precipitate with a solution of nitrate of silver; and will count as
chloride if they are not removed.
_The Standard Solution of Nitrate of Silver_ is made by dissolving 23.94
grams of the salt (AgNO_{3}) in distilled water, and diluting to 1
litre; 100 c.c. are equal to 0.5 gram of chlorine.
The _indicator_ is made by adding silver nitrate to a strong neutral
solution of yellow chromate of potash (K_{2}CrO_{4}), till a permanent
red precipitate is formed. The solution is allowed to settle, and the
clear liquid decanted into a stoppered bottle labelled "chromate
indicator for chlorine."
Standardise the silver nitrate by weighing up 0.5 gram of pure sodium
chloride (or potassium chloride). Transfer to a flask and dissolve in
distilled water; dilute to 100 c.c. Fill an ordinary burette with the
standard silver solution, and (after adjusting) run into the flask a
quantity sufficient to throw down the greater part of the chlorine. Add
a few drops of the chromate indicator and continue the addition of the
silver nitrate until the yellow colour of the solution becomes
permanently tinted red, after shaking. This shows that the chlorine is
all precipitated, and that the chromate is beginning to come down. The
further addition of a couple of drops of the silver solution will cause
a marked difference in the tint. Read off the quantity run in, and
calculate the standard. One gram of sodium chloride contains 0.6062 gram
of chlorine; and 1 gram of potassium chloride contains 0.4754 gram.
For the determination of small quantities of chloride (a few
milligrams), the same method is used; but the standard solution is
diluted so that each c.c. is equal to 1 milligram of chlorine; and the
chromate indicator is added before titrating. The standard solution is
made by measuring off 200 c.c. of the solution described above, and
diluting with distilled water to 1 litre.
BROMINE AND BROMIDES.
Bromine closely resembles chlorine in the nature of its compounds. It
does not occur free in nature, but is occasionally found in combination
with silver as bromargyrite (AgBr) and, together with chloride, in
embolite. It mainly occurs as alkaline bromides in certain natural
waters. Nearly all the bromine of commerce is derived from the mother
liquors of salt-works--_i.e._, the liquors from which the common salt
has been crystallised out. Bromine combines directly with the metals,
forming a series of salts--the bromides. In ordinary work they are
separated with, and (except when specially tested for) counted as,
chlorides. They are detected by adding chlorine water to the suspected
solution and shaking up with carbon bisulphide. Bromine colours the
latter brown.
IODINE AND IODIDES.
Iodine does not occur in nature in the free state; and iodides are rare,
iodargyrite or iodide of silver (AgI) being the only one which ranks as
a mineral species. Iodates are found associated with Chili saltpetre,
which is an important source of the element.
Iodine and Iodides are largely used in the laboratory, and have already
been frequently referred to. It is used as an oxidising agent in a
similar manner as permanganate and bichromate of potash, especially in
the determinations of copper, arsenic, antimony, and manganese.
Iodine is not readily soluble in water; but dissolves easily in a
concentrated solution of potassium iodide. Its solutions are strongly
coloured; a drop of a dilute solution colours a large volume of water
decidedly yellow; on the addition of starch paste, this becomes blue.
The delicacy of this reaction is taken advantage of in titrations to
determine when free iodine is present. The blue colour may be
alternately developed and removed by the addition of iodine (or an
oxidising agent) and hyposulphite of soda (or some other reducing
agent). In decolorising, the solution changes from blue or black to
colourless or pale yellow according to circumstances. Sometimes the
solution, instead of remaining colourless, gradually develops a blue
which recurs in spite of the further addition of the reducing agent. In
these cases the conditions of the assay have been departed from, or (and
this is more often the case) there is some substance present capable of
liberating iodine.
Iodine forms a series of salts--the iodides--resembling in many respects
the chlorides. These can be obtained by direct combination of the metals
with iodine.
~Detection.~--Free iodine is best recognised by the violet vapours
evolved from the solution on heating, and by the blue or black colour
which it strikes on the addition of starch paste. Iodides are detected
by boiling with strong solutions of ferric sulphate or chloride. Iodine
is liberated, distilled over, and collected. Chlorine also liberates
iodine from iodides; and this reaction is frequently made use of in
assaying. A process based on this is described under _Manganese_. All
substances which liberate chlorine on boiling with hydrochloric acid
(dioxides, bichromates, permanganates, &c.) are determined in a similar
way.
~Solution and Separation.~--Most iodides are soluble in water or dilute
acids. The separation is effected by distilling the substance with
solution of ferric sulphate, and collecting the vapour in a dilute
solution of sulphurous acid or arsenite of soda. On the completion of
the distillation, the iodine will be in the distillate as iodide; and
the gravimetric determination is made on this.
GRAVIMETRIC DETERMINATION.
To the solution containing the iodine, as iodide, and which is free from
chlorides (and bromides), add a little dilute nitric acid and nitrate of
silver till no further precipitate is produced. Filter off, wash with
hot water, and dry. Clean the filter-paper as much as possible, and burn
it. Collect the ash in a weighed porcelain crucible, add the main
portion, and heat to incipient fusion; cool, and weigh. The substance is
silver iodide, and contains 54.03 per cent. of iodine.
VOLUMETRIC METHOD.
This is for the titration of free iodine, and is practically that which
is described under _Manganese_. The substance to be determined is
distilled with ferric sulphate, and the iodine is collected in a
solution of potassium iodide, in which it readily dissolves. If flaky
crystals separate out in the receiver, more potassium iodide crystals
are added. When the distillation is finished, the receiver is
disconnected, and its contents washed out into a beaker and titrated
with "hypo." The standard solution of "hypo" is made by dissolving 19.58
grams of hyposulphite of soda (Na_{2}S_{2}O_{3}.5H_{2}O) in water and
diluting to 1 litre; 100 c.c. are equal to 1 gram of iodine. To
standardise the solution, weigh up 0.25 gram of pure iodine in a small
beaker. Add 2 or 3 crystals of potassium iodide; cover with water; and,
when dissolved, dilute to 50 or 100 c.c. Titrate, and calculate the
standard.
FLUORINE AND FLUORIDES.
Fluorine is frequently met with as calcium fluoride or fluor-spar
(CaF_{2}). It occurs less abundantly as cryolite (Na_{3}AlF_{6}), a
fluoride of aluminium and sodium, which is used in glass-making. Certain
other rarer fluorides are occasionally met with. Fluorine is also found
in apatite, and in some silicates, such as topaz, tourmaline, micas, &c.
Hydrofluoric acid is used for etching glass and opening up silicates. It
attacks silica, forming fluoride of silicon (SiF_{4}), which is
volatile. Silica is by this means eliminated from other oxides, which,
in the presence of sulphuric acid, are fixed. The commercial acid is
seldom pure, and generally weak; and the acid itself is dangerously
obnoxious. The use of ammonium fluoride (or sodium fluoride) and a
mineral acid is more convenient. Determinations of this kind are made in
platinum dishes enclosed in lead or copper vessels in a well-ventilated
place. Fluor-spar is useful as a flux in dry assaying; it renders slags,
which would otherwise be pasty, quite fluid. Fluorides generally are
fusible, and impart fusibility to substances with which they form weak
compounds. Their fluxing action does not depend on the removal of
silicon as fluoride.
~Detection.~--Fluorides in small quantity are easily overlooked unless
specially sought for. In larger amounts they are recognised by the
property hydrofluoric acid has of etching glass. A watch-glass is
warmed, and a layer of wax is melted over the convex side. When cold,
some lines are engraved on the waxed surface with any sharp-pointed
instrument. The substance to be tested is powdered; and moistened, in a
platinum dish, with sulphuric acid. The watch-glass is filled with cold
water and supported over the dish. The dish is then carefully warmed,
but not sufficiently to melt the wax. After a minute or two, the glass
is taken off, and the wax removed. If the substance contained fluorine,
the characters will be found permanently etched on the glass. An equally
good, but more rapid, test is to mix the powdered substance with some
silica, and to heat the mixture in a test tube with sulphuric acid.
Silicon fluoride is evolved, and, if a moistened glass rod is held in
the tube, it becomes coated with a white deposit of silica, formed by
the decomposition of the silicon fluoride by the water. This is also
used as a test for silica; but in this case the substance is mixed with
a fluoride, and the experiment must obviously be carried out in a
platinum vessel.
~Separation and Determination.~--The determination of fluorine is
difficult. In the case of fluorides free from silicates (such as
fluor-spar), it is determined indirectly by decomposing a weighed
portion with sulphuric acid, evaporating, igniting, and weighing the
residual sulphate. The increase in weight multiplied by 0.655 gives the
weight of fluorine.
In the presence of silica this method does not answer, because of the
volatilisation of silicon fluoride. In these cases Wöhler adopted the
following plan, which resembles that for the indirect determination of
carbon dioxide. Mix the weighed substance in a small flask with powdered
silica and sulphuric acid. The mouth of the flask is closed with a cork
carrying a tube which is filled, the first half with calcium chloride
and the second half with pumice coated with dried copper sulphate. The
apparatus is weighed quickly, and then warmed till decomposition is
complete. A current of dry air is aspirated for a minute or two; and the
apparatus again weighed. The loss in weight gives that of the silicon
fluoride (SiF_{4}), which, multiplied by 0.7307, gives the weight of
fluorine.
Fresenius uses the same reaction; but collects and weighs the silicon
fluoride. The finely powdered and dried substance is mixed with ten or
fifteen times its weight of ignited and powdered silica. The mixture is
introduced into a small dry flask connected on one side with a series of
drying-tubes, and on the other with an empty tube (to condense any
sulphuric acid). To this last is joined a drying-tube containing
chloride of calcium and anhydrous copper sulphate. This is directly
connected with a series of three weighed tubes in which the fluoride of
silicon is collected. The last of these is joined to another
drying-tube. The first weighed tube contains pumice and cotton wool,
moistened with water; the second tube contains soda-lime as well as (in
the upper half of the second limb) fused calcium chloride between plugs
of wool; the third tube is filled half with soda-lime and half with
fused calcium chloride. The distilling-flask containing the substance
mixed with silica is charged with 40 or 50 c.c. of sulphuric acid, and
placed on the hot plate. Alongside it is placed a similar dry flask
containing a thermometer, and the temperature in this is kept at 150° or
160° C. A current of air is sent through the tubes during the operation,
which takes from one to three hours for from 0.1 to 1 gram of the
substance. A correction is made by deducting 0.001 gram for every hour
the dried air has been passed through. The increase in weight of the
three tubes gives the weight of the silicon fluoride.
Penfield uses a similar arrangement, but passes his silicon fluoride
into an alcoholic solution of potassium chloride. Silica and potassium
silico-fluoride are precipitated, and hydrochloric acid is set
free.[100] The acid thus liberated is titrated, with a standard solution
of alkali, in the alcoholic solution, and from the amount of free acid
found the fluorine is calculated. The weight of hydrochloric acid (HCl)
found, multiplied by 1.562, gives the weight of the fluorine. With this
method of working, fewer ~U~-tubes are required. The exit tube from the
flask is bent so as to form a small ~V~, which is kept cool in water;
this is directly connected with the ~U~-tube containing the alcoholic
solution of potassium chloride. The flask with the assay is heated for
about two hours, and a current of dry air is aspirated throughout the
determination. Fluoride of silicon is a gas not easily condensed to a
liquid: but is immediately decomposed by water or moist air.
FOOTNOTES:
[95] This will require two or three hours to thoroughly complete. It is
best to powder the oxide first produced, and recalcine.
[96] No magnetic oxide was formed.
[97] For example:--
CaO + 2HCl = CaCl_{2} + H_{2}O.
PbO + H_{2}SO_{4} = PbSO_{4} + H_{2}O.
MgO + 2HNO_{3} = Mg(NO_{3})_{2} + H_{2}O.
Al_{2}O_{3} + 6HCl = Al_{2}Cl_{6} + 3H_{2}O.
Fe_{2}O_{3} + 3H_{2}SO_{4} = Fe_{2}(SO_{4})_{3} + 3H_{2}O.
[98] Fe_{2}O_{3}: 2FeO:: 0.2877: 0.2589.
[99] 100 c.c. contain 1 gram of sulphuric acid.
[100] 3SiF_{4} + 4KCl + 2H_{2}O = 2K_{2}SiF_{6} + SiO_{2} + 4HCl.
CHAPTER XVI.
SULPHUR AND SULPHATES.
Sulphur occurs native in volcanic districts, and is mined in Sicily,
Italy, and California in considerable quantities. Combined with metals
(sulphides), it is common in all mineral districts. Iron pyrites
(FeS_{2}) is the most abundant source of this element. Sulphates, such
as gypsum, are fairly common, but have no value so far as the sulphur in
them is concerned. In coal it exists as an impurity, occurring partly as
a constituent of organic compounds.
Sulphur, whether free or combined with metals, forms, on burning,
sulphurous oxide (SO_{2}), which by the action of oxidising agents and
water is converted into sulphuric acid. It forms two oxides, sulphurous
(SO_{2}) and sulphuric (SO_{3}), which combine with bases to form
sulphites and sulphates. Sulphites are of little importance to the
assayer, and are converted into sulphates by the action of nitric acid
and other oxidising agents.
The native sulphides, when acted on with hydrochloric acid, give off
sulphuretted hydrogen; with nitric acid or aqua regia, sulphates are
formed, and more or less sulphur separated.
Sulphur is detected in sulphides by the irritating odour of sulphurous
oxide given off on roasting, by the evolution of sulphuretted hydrogen
when treated with hydrochloric acid, or by a white precipitate of barium
sulphate formed when barium chloride is added to the aqua regia
solution.
~Dry Assay.~--There is no method of general application. Free or native
sulphur may be volatilised, condensed, and weighed, but pyrites only
gives up a portion of its sulphur when heated in a closed vessel, while
most sulphides, and all sulphates, give up none at all.
In the determination of sulphur in brimstone, 10 grams of the substance
are taken, placed in a small porcelain dish, heated over a Bunsen burner
in a well-ventilated place, and ignited. When the sulphur has been
completely burnt off, the residue (which consists chiefly of sand) is
collected and weighed. In a separate portion the moisture and arsenic
are determined; the amounts of these are deducted from the loss in the
first experiment. The difference, multiplied by 10, gives the percentage
of sulphur.
WET METHODS.
~Solution.~--All sulphates, excepting those of lead, barium, strontium,
and lime, are soluble in water or dilute acid. All sulphides, except
cinnabar, are converted into sulphates by the action of nitric acid at a
gentle heat; or, better, by the action of a mixture of three volumes of
nitric acid and one volume of hydrochloric acid. This last attacks
cinnabar as well. With most substances it is difficult to convert the
whole of the sulphur into sulphuric acid. The sulphur separates out at
first as a dark spongy mass, which (on continued treatment) changes to
light-coloured flakes. When the solution becomes concentrated and the
temperature rises sufficiently, the sulphur fuses into one or more
honey-coloured globules which, owing to the small surface they oppose to
the acid, are very slowly oxidised. It is not desirable to assist the
formation of these globules; therefore, the temperature is kept as low
as possible, and strong nitric acid is used. When such globules form, it
is best to allow the solution to cool, when the globules will solidify.
They can then be filtered off and picked out from the insoluble residue,
dried, weighed, ignited, and again weighed, the loss being counted as
sulphur. With iron pyrites this difficulty seldom occurs.
Metallic sulphides when fused with an excess of nitre are completely
oxidised. If the ore is rich in sulphur, some inert body (such as sodium
chloride, or, better, sodium carbonate) is added to dilute the action.
With pure sulphur, the action is so energetic as to cause an explosion,
so that care should be taken. With burnt ores (incompletely calcined
pyrites), there is sufficient oxide of iron present to prevent too rapid
action.
These fusions with nitre are best conducted in a platinum dish covered
with a piece of platinum foil. The ore is ground with the nitre to
ensure complete mixing. The heat need not be excessive, so that a single
Bunsen burner placed beneath the dish will suffice; if the bottom of the
dish is seen to be red-hot, it is sufficient. On cooling and extracting
with water, the sulphur will pass into solution as potassium sulphate,
which is then filtered off from the insoluble oxides of iron, copper,
&c. The filtrate, after having been treated with a large excess of
hydrochloric acid, evaporated to dryness, and re-dissolved in water, is
ready for the determination.
Lead sulphate may be dissolved by boiling with ammonium acetate. The
insoluble sulphates of barium, strontium, and lime, are decomposed by
fusing with 4 or 5 times their weight of "fusion mixture." The alkaline
sulphates are then dissolved out with water, and filtered off from the
insoluble residue. The filtrate is rendered acid with hydrochloric acid.
~Separation.~--The determination of the sulphuric acid in these
solutions by precipitation with barium chloride also serves as a
separation; but in hot acid solutions containing copper, and more
especially iron salts, the baric sulphate has a strong tendency to carry
down amounts of those bodies, varying, no doubt, with the conditions of
the precipitation. Boiling hydrochloric acid fails to completely extract
them. Moreover, the use of hot concentrated hydrochloric acid causes a
loss by dissolving barium sulphate. Nitric acid and nitrates must be
decomposed by prolonged boiling and evaporation with hydrochloric acid.
The iron may be removed by adding a slight excess of ammonia to the
faintly acid solution, filtering off, and washing the precipitated
ferric hydrate with hot water. By slightly acidulating the filtrate with
hydrochloric acid, it will be rendered ready for the determination.
GRAVIMETRIC METHOD.
This assay is one of those which strikingly shows the necessity of
getting the assay solution under proper conditions, in order to obtain
satisfactory results. The method has been repeatedly investigated, and
the conclusion arrived at, "that it can be correct only by accident."
Yet there are many chemists who get good results, and place considerable
faith in its accuracy. This can only be due to differences in the manner
of working. It is generally understood that nitric acid or nitrates must
be absent; and our experience fully confirms this. Precipitations in
nitrate solutions are worthless, as the following experiments show. In
each experiment the bulk of the solution was 150 c.c. The solutions
contained 10 grams of nitre, were freely acid with hydrochloric acid,
and were precipitated (while boiling) with slight excess of baric
chloride.
Sulphuric acid taken 0.020 gram 0.050 gram 0.100 gram
" found 0.019 " 0.047 " 0.098 "
" taken 0.500 " 1.004 " 1.000 "
" found 0.526 " 1.126 " 1.126 "
All the precipitates were boiled with hydrochloric acid, and thoroughly
washed before weighing. The results of some other experiments on this
subject are given under "sulphur" in the "examination of commercial
copper," page 207.
The solution having been obtained free from nitrates and chlorates (and
containing but little free hydrochloric acid), is largely diluted,
heated to boiling, and precipitated with a moderate excess of a solution
of chloride of barium (8 parts of the crystallized barium chloride are
sufficient for 1 of sulphur). It is allowed to settle for half-an-hour,
and then decanted through a filter. The precipitate is shaken up with
boiling water, rendered slightly acid, filtered, washed, dried, ignited,
and weighed. The ignited precipitate, when pure, is white, and is not
decomposed at a red heat; it is barium sulphate (BaSO_{4}), and contains
13.73 per cent. of sulphur, or 34.33 per cent. of sulphuric oxide
(SO_{3}).
~Determination of Sulphur in Pyrites.~--Weigh up half a gram of the
dried and powdered sample, and treat with 10 c.c. of a mixture of 3
volumes of nitric acid and 1 volume of hydrochloric acid, occasionally
heating. Evaporate to dryness, treat with 5 c.c. of hydrochloric acid,
and again evaporate; take up with 1 c.c. of hydrochloric acid and 100
c.c. of hot water, filter through a small filter, and wash. The residue
may contain sulphates of lead, barium, or lime; it must be separately
examined, if the total sulphur is wanted. The filtrate is heated, and
rendered slightly alkaline with ammonia. Filter off the precipitated
ferric hydrate through a quick filter, and wash with hot water. If
necessary, evaporate the bulk to about 200 c.c., render faintly acid
with hydrochloric acid, and add 20 c.c. of solution of barium chloride;
allow to stand for half-an-hour, and decant through a filter. Wash with
hot water, dry, ignite, and weigh. Pure pyrites contains 53.33 per cent.
of sulphur.
VOLUMETRIC METHOD.
This is based upon the easy conversion of all sulphur compounds into
sulphates by fusion with nitre or by oxidation with nitric acid; and on
the determination of the sulphate formed by titration in an acetic acid
solution with baric chloride.[101] The finishing point is determined by
filtering off portions of the assay solution, and testing with sulphuric
acid. A slight excess of baric chloride will cause a precipitate.
The process may be divided into--(1) the preparation of the solution,
and (2) the titration.
~Preparation of the Solution.~--Weigh up from 1 to 5 grams of the dried
and powdered substance, and mix intimately with 4 grams of powdered
nitre; clean out the mortar with another gram of nitre, and add this as
a cover. Heat in a platinum crucible for fifteen minutes at a low
temperature; cool, and extract with water in an evaporating-dish about
9 inches across, and holding 700 or 800 c.c. Add 10 grams of sodium
acetate and 10 c.c. of acetic acid, and dilute to half a litre. Boil.
The solution is ready for titrating. Substances which lose sulphur on
heating (such as pyrites) are thus treated:--Weigh up 1 gram, and
evaporate nearly to dryness with 10 c.c. each of nitric and hydrochloric
acids. Take up with 10 c.c. of hydrochloric acid, and again boil down to
a small bulk; dilute and transfer to a 9-inch evaporating-dish; add 10
grams of sodium acetate and 5 c.c. of acetic acid, dilute to half a
litre, and boil. The solution is ready for titrating. Sulphates may be
dissolved up in the dish itself with the help of a c.c. or so of
hydrochloric acid; sodium acetate and acetic acid are then added; and,
after dilution and boiling, the solutions are at once titrated.
The solution before titration must contain no free mineral acid, but 5
or 10 c.c. of acetic acid should be present. It must contain 10 grams of
sodium acetate, or sufficient to convert any free mineral acid into its
corresponding sodic salt; or, if chlorides, nitrates or sulphates of the
metals are present, sufficient to decompose them. If a precipitation
occurs, as is the case with ferric salts, &c., the solution is titrated
with the precipitate in it.
~The Titration.~--_The standard solution of barium chloride_ is made by
dissolving 76.25 grams of the crystallized salt (BaCl_{2}.2H_{2}O) in
distilled water, and diluting to 1 litre. 100 c.c. will equal 1 gram of
sulphur. As indicator, use dilute sulphuric acid. The strength of the
solution may be checked by the titration of 5 grams of ferrous sulphate
(oxidized with permanganate of potassium or a few drops of nitric acid),
which should require 57.5 c.c. of the barium chloride solution; or any
pure sulphate of known composition can be used; anhydrous salts should
be preferred.
[Illustration: FIG. 65.]
Fill an ordinary 100 c.c. burette with the solution of barium chloride.
The evaporating dish containing the assay solution is placed on a round
burner (as shown in fig. 65), and the solution is kept steadily boiling.
An ordinary Bunsen-burner flame will cause bumping, and should not be
used. Run in the standard solution in quantity known to be
insufficient; then withdraw a portion of about 2 c.c., with a pipette,
and filter through a fine filter-paper into a test tube. Run in another
0.5 c.c. of the standard solution, and withdraw and filter into a test
tube another portion of 2 c.c.; and continue this operation until
half-a-dozen or more portions have been drawn off. The test tubes should
be arranged in order in a stand resting on a piece of paper, so that
each test tube representing 0.5 c.c. of the standard baric chloride may
have its value recorded beneath it (fig. 66). Add to each test tube 3
drops of dilute sulphuric acid; that which shows the first appearance of
a precipitate marks the point at which the titration is complete.
Suppose, for example, that the test tube marked 48.5 c.c. shows no
precipitate, while that at 49.0 c.c. shows one, it is evident that the
finishing point lies between these readings. With a little practice, one
can judge from the appearance of the precipitate in the 49 c.c. tube,
whether 1/4 c.c. should be deducted or not.
[Illustration: FIG. 66.]
It is better to add dilute sulphuric acid, and to watch for the
appearance of a precipitate in the test tube, than to add baric chloride
and to look for its non-appearance; besides, baric chloride is much less
likely to be present in a test tube as impurity than sulphates are. In
this way the chance of error from what are termed "accidental causes" is
diminished.
The following experiments show the effect of variation in the conditions
of titration:--
Make _a standard solution of sulphuric acid_ by diluting 43.65 grams of
sulphuric acid (sp. g. 1.6165) to 1 litre: 100 c.c. will contain 1 gram
of sulphur. An equivalent solution may be made by dissolving 100.62
grams of sodium sulphate crystals (Na_{2}SO_{4}.10H_{2}O), or 86.88
grams of ferrous sulphate (FeSO_{4}.7H_{2}O), in water (oxidising the
latter), and diluting to 1 litre.
The order in which these experiments are given is that in which they
were made in an investigation into the conditions under which the
titration could most accurately be effected.
~Effect of Hydrochloric and Nitric Acids.~--The titrations were
performed in the manner already described, but sodic acetate and acetic
acid were absent. Twenty c.c. of the standard solution of sulphuric acid
were used.
Hydrochloric acid present 0.0 c.c. 1.0 c.c. 2.0 c.c. 5.0 c.c.
"Baric chloride" required 20.0 " 20.0 " 19.7 " 12.5 "
Nitric acid present 0.0 c.c. 1.0 c.c. 2.0 c.c. 5.0 c.c.
"Baric chloride" required 20.0 " 19.5 " 18.0 " 10.0 "
These show clearly the interference of free mineral acids, although very
dilute hydrochloric acid (1 c.c. in 500 of water) has no effect.
~Effect of Acetic and Citric Acids.~--A similar series of experiments
with these acids gave the following results:--
Acetic acid present 0.0 c.c. 5.0 c.c. 50.0 c.c. 100.0 c.c.
"Baric chloride" required 20.0 " 20.0 " 20.0 " 20.0 "
Citric acid present 0 gram 1 gram 5 grams
"Baric chloride" required 20.0 c.c. 20.0 c.c. 20.0 c.c.
These acids do not interfere.
~Effect of Sodic Acetate and Acetic Acid.~--In each of these experiments
5 c.c. of acetic acid was present.
Sodium acetate added 0 gram 1 gram 10 grams 50 grams
"Baric chloride" required 20.0 c.c. 20.0 c.c. 20.0 c.c. 20.0 c.c.
As sodic acetate and acetic acid did not interfere, it became desirable
to make some experiments on the _finishing point_. The first object
sought for was the smallest amount of the standard baric chloride in 500
c.c. of water, required to give an indication when tested in the manner
already described.
Baric Chloride
Conditions of Assay Solution. required.
Water only 0.05 c.c.
With 10 grams of sodium acetate and 5 c.c. of
acetic acid 0.05 "
The same with 5 grams of nitre 0.10 "
Like the last, but with 5 grams of salt instead of
nitre 0.10 "
These show that as small an amount of baric chloride solution as is
equal to only 0.000002 gram of sulphur in the 2 c.c. of solution tested
yields a decided precipitate on the addition of 3 drops of sulphuric
acid.
To determine whether the same finishing point is obtained on testing the
filtered portions in the test tubes with baric chloride as is obtained
on testing with sulphuric acid, a titration was made with 20 c.c. of
standard solution of sulphuric acid, together with the usual quantities
of sodic acetate and acetic acid; and two lots of 2 c.c. each were
filtered into two sets of test tubes after each addition of the standard
baric chloride. To one series 3 drops of baric chloride solution were
added, and to the other 3 drops of sulphuric acid. The results were--
With Dilute With Baric
"Baric Chloride" added. Sulphuric Acid. Chloride Solution.
19.5 c.c. Clear Cloudy
19.75 " Clear Cloudy
20.0 " Finished Finished
20.25 " Cloudy Clear
20.5 " Cloudy Clear
The two methods of testing give the same result. But this balance is
disturbed in the presence of much nitre, the indications with baric
chloride being disturbed by an opalescence for some c.c. beyond the
finishing point. In solutions containing free hydrochloric or nitric
acid, a precipitate is obtained with either baric chloride or sulphuric
acid.
~Effect of Varying Sulphur.~--In these and the subsequent experiments
the titrations were performed in the presence of 10 grams of sodic
acetate and 10 c.c. of acetic acid in the manner already described.
Standard sulphuric
acid used 5.0 c.c. 10.0 c.c. 20.0 c.c. 50.0 c.c. 100.0 c.c.
"Baric chloride"
required 5.0 " 10.0 " 20.0 " 50.0 " 100.0 "
~Effect of Varying Temperature.~--With 5 c.c. of standard sulphuric acid
titrated at 15° C., 5 c.c. of baric chloride were required; but with
larger quantities the results were altogether unsatisfactory when
titrated cold.
~Effect of Varying Bulk.~--
Bulk 100.0 c.c. 200.0 c.c. 500.0 c.c. 1000.0 c.c.
"Baric chloride"
required 20.0 " 20.0 " 20.0 " 20.5 "
Considerable variation in bulk has no effect, but 500 c.c. is the most
convenient volume to work with. It is well to occasionally replace the
water boiled off during titration.
~Effect of Foreign Salts.~--In all these experiments 20 c.c. of
"sulphuric acid" were used, and the titration was performed in the
ordinary way.
_Sodic chloride_ added 0 gram 5 grams 10 grams
"Baric chloride" required 20.0 c.c. 20.0 c.c. 19.7 c.c.
_Ammonic chloride_ added 0 gram 5 grams 10 grams
"Baric chloride" required 20.0 c.c. 20.0 c.c. 19.5 c.c.
_Calcic chloride_ added 0 gram 1 gram 2 grams 5 grams
"Baric chloride" required 20.0 c.c. 20.0 c.c. 19.2 c.c. 19.0 c.c.
_Zinc chloride_ added 0 gram 1 gram 3 grams 5 grams
"Baric chloride" required 20.0 c.c. 20.0 c.c. 20.0 c.c. 20.0 c.c.
_Ferrous chloride_ added 0 gram 1 gram 3 grams 5 grams
"Baric chloride" required 20.0 c.c. 19.7 c.c. 19.5 c.c. 19.0 c.c.
_Ferric chloride_ added 0 gram 1 gram 3 grams 5 grams
"Baric chloride" required 20.0 c.c. 20.0 c.c. 20.0 c.c. 20.0 c.c.
_Copper chloride_ added 0 gram 1 gram 3 grams 5 grams
"Baric chloride" required 20.0 c.c. 20.0 c.c. 20.0 c.c. 20.0 c.c.
_Potassic Nitrate_ added 0 gram 1 gram 5 grams 10 grams
"Baric chloride" required 20.0 c.c. 20.0 c.c. 20.0 c.c. 19.0 c.c.
_Potassic Nitrite_ added 0 gram 1 gram 5 grams
"Baric chloride" required 20.0 c.c. 20.0 c.c. 20.0 c.c.
_Sodic phosphate_ added 0 gram 1 gram
"Baric chloride" required 20.0 c.c. 22.5 c.c.
_Sodic arsenate_ added 0 gram 1 gram
"Baric chloride" required 20.0 c.c. 20.5 c.c.
In the absence of ferric salts, phosphates and arsenates count as
sulphur.
In two series of experiments for determining the effect of varying
amounts of sulphur in the form of ferrous sulphate, we obtained the
following results:--In the first series the assay solution was prepared
in the manner we have described for _Pyrites_; and in the second series,
by fusion with nitre.
Sulphur added 0.050 gram 0.100 gram 0.200 gram
"Baric chloride" required (1) 5.0 c.c. 10.0 c.c. 20.0 c.c.
" " (2) 4.7 " 10.0 " 20.0 "
Sulphur added 0.500 gram 1.000 gram
"Baric chloride" required (1) 50.0 c.c. 100.0 c.c.
" " (2) 50.0 " 100.0 "
More than 5 grams of nitre must not be used in an assay; and, since the
requisite amount of nitre considerably exceeds that sufficient to
oxidise the sulphur, not more than 0.5 gram of unoxidised sulphur should
be present in the portion of the sample weighed up for determination.
When the amount of sulphur present is not known within reasonable
limits, the test portions may be tried with a drop of baric chloride
solution instead of sulphuric acid, so that the diminishing quantity of
precipitate may give warning of an approach to the finishing point.
~Determination of Sulphur in Blende.~--Weigh up 1 gram of dried and
powdered blende, and mix and fuse with 5 grams of nitre in the manner
described. Place the dish and its contents in the titrating-dish,
extract with water, add 10 grams of sodium acetate and 10 c.c. of acetic
acid, remove and wash the platinum-dish, and dilute to 500 c.c.; boil
and titrate. In the example, duplicate determinations required (a)
32.0 c.c., (b) 32.25 c.c., giving an average of 32.1 per cent. of
sulphur.
~Determination of Sulphur in Chalcopyrite~ (Yellow Copper Ore).--Take 1
gram of the finely-powdered sample, and 5 grams of nitre. Sprinkle a
little of the nitre in a small Wedgwood mortar, place the ore on it, and
cover with 2 or 3 grams more of the nitre. Rub up together, and transfer
to a small porcelain dish; clean out the mortar with the rest of the
nitre, and add to the contents of the dish. Cover with a piece of
platinum foil, and heat gently with a Bunsen burner till the nitre melts
and the stuff shows signs of deflagrating; remove the heat, and allow
the action to go on by itself for a minute or so, then heat over the
Bunsen burner for 10 minutes. Cool; transfer the whole to the
titrating-dish; boil with 500 c.c. of water; remove the small dish and
foil; add sodic acetate and acetic acid, and titrate.
For example, 1 gram required 34.5 c.c. of "barium chloride" (standard =
1.005 gram S), which is equivalent to 34.7 per cent. sulphur. The
theoretical percentage is 34.8.
~Determination of Sulphur in Chalcocite~ (Grey Copper Ore).--Proceed as
in the last experiment but, since the action with nitre is more
moderate, no special precautions need be taken on heating. A platinum
dish may be used.
An example which was heated for 30 minutes required 20.5 c.c. of the
barium chloride solution. This is equivalent to 20.6 per cent. of
sulphur. The theoretical yield is 20.2 per cent.
~Determination of Sulphur in Pyrites.~--Take 1 gram of the
finely-powdered sample, cover with 10 c.c. of nitric acid, and, when
action has ceased, evaporate to a small bulk. Add 3 or 4 c.c. of
hydrochloric acid, and again evaporate to a paste. Take up with 1 or 2
c.c. of dilute hydrochloric acid, dilute with water, transfer to a
titrating-dish, add 10 grams of sodic acetate and 5 c.c. of acetic acid,
and dilute with water to 500 c.c. Boil and titrate.
An example with 1 gram of a pure crystallized pyrites required 52.7 c.c.
of the barium chloride solution, which is equivalent to 53.0 per cent.
of sulphur. Theory requires 53.3 per cent. of sulphur.
~Determination of Sulphur in Mispickel.~--Take 1 gram of the powdered
ore and evaporate with 10 c.c. of nitric acid, and take up with 3 or 4
c.c. of hydrochloric acid. If any globules of sulphur remain, again
evaporate with nitric acid. Dilute, and transfer to the titrating-dish.
Add 10 grams of sodic acetate, dilute with water, boil, and titrate. The
mispickel carries (according to theory) exactly sufficient iron to
precipitate the arsenic as ferric arsenate in an acetic acid solution,
so no more iron need be added. The ferric arsenate will separate out as
a yellowish-white flocculent precipitate.
An example required, in duplicate experiment, 18.5 c.c. and 18.7 c.c. of
barium chloride, equivalent to 18.7 per cent. of sulphur. The formula,
FeS_{2}.FeAs_{2}, requires 19.6 percent., but the sulphur generally
varies considerably from this amount.
~Determination of Sulphur in Burnt Ores.~--Take 5 grams of the dried and
powdered ore, and rub up with 4 grams of nitre; transfer to the
platinum-dish; clean out the mortar with another gram of nitre, and add
this as a cover. Heat, and extract with water as before; add the sodium
acetate and acetic acid; and titrate. Burnt ores carry from 2.5 to 5 per
cent. of sulphur. A series of four determinations gave:--
"Baric Chloride" Required. Percentage of Sulphur. Gravimetric Results.
12.6 c.c. 2.52 % 2.45 %
29.9 " 5.98 " 5.84 "
18.1 " 3.62 " 3.53 "
22.0 " 4.40 " 4.43 "
For ores carrying less than 1 per cent. of sulphur, take 10 grams for
the assay.
~Determination of Sulphuric Oxide (SO_{3}) in Sulphates.~--When the
sulphur exists in the sample received by the assayer in an oxidised
state as sulphate, it is usual to report it in terms of sulphuric oxide
(SO_{3}). In this case, the metal must also be reported as oxide. For
example, an analysis of copper sulphate would be thus reported:--
Oxide of copper (CuO) 31.8 %
Sulphuric oxide (SO_{3}) 32.1 "
Water 36.1 "
-----
100.0
The percentage of sulphur multiplied by 2.5 gives the percentage of
sulphuric oxide. Thus a sample of copper sulphate containing 12.85 per
cent. of sulphur will contain 12.85 × 2.5 or 32.12 per cent. of
sulphuric oxide.
In minerals and metallurgical products, it is common to find the sulphur
in both conditions--_i.e._, as sulphate and sulphide. Generally in these
the percentage of sulphur only is wanted; but this will depend entirely
on commercial requirements, and not on the fancy of the assayer.
Soluble sulphates are determined separately by extracting with small
quantities of cold water, so as to avoid the separation of basic
sulphates, or, if the sulphides present are not at the same time
attacked, by dilute hydrochloric acid. Lead sulphate may be extracted by
boiling with ammonic acetate; whilst barium, strontium, and, perhaps,
calcium sulphate, will be mainly found in the residue insoluble in
acids.
Weigh up from 2 to 5 grams of the material according to the amount of
sulphur judged to be present, and dissolve them in the titrating-dish
with 1 c.c. of hydrochloric acid and 50 c.c. of water. Add 10 grams of
sodic acetate, and 10 c.c. of acetic acid; dilute, boil, and titrate. In
the case of ferric salts, half the quantity of acetic acid will be
better, as then the ferric iron will be precipitated, and a colourless
solution will be left, in which the end reaction is more readily
distinguished.
Determined in this way, 5 gram samples of the following salts gave the
results indicated below:--
"Barium Chloride"
Salt. Required. Sulphuric Oxide.
Copper sulphate 64.25 c.c. 32.12 %
Magnesium sulphate 65.25 " 32.62 "
Zinc sulphate 56.25 " 28.12 "
Ferrous sulphate 58.25 " 29.12 "
Sodium sulphate 51.25 " 25.60 "
~Determination of Sulphuric Oxide in Barytes~ (Heavy spar).--Fuse 2
grams of the powdered mineral with 5 grams of "fusion mixture" for five
minutes; and, when cold, extract with water. Filter, acidulate the
filtrate with an excess of 10 c.c. of acetic acid, dilute, boil, and
titrate. For example, a transparent crystallised sample required 27.0
c.c. of barium chloride, which is equivalent to 13.6 per cent. of
sulphur, or 34.0 per cent. of sulphuric oxide. Theory requires 34.3 per
cent. of the latter. Since both carbonate of soda and potash are liable
to contain sulphates, a blank determination should be made on 5 grams of
the "fusion mixture," and the amount found be deducted from that got in
the assay.
PRACTICAL EXERCISES.
1. The price of sulphur in an ore being 4-1/2d. per unit in the
northern markets, what would be the price of a ton of ore containing 49
per cent. of sulphur? What would be the effect on the price of an error
of 0.25 per cent. in the assay?
2. Pyrites carries 50 per cent. of sulphur, and on calcining yields 70
per cent. of its weight of burnt ore. Supposing the burnt ore carries
3.5 per cent. of sulphur, what proportion of the sulphur will have been
removed in the calcining?
3. How would blende compare with pyrites as a source of sulphur for
sulphuric acid making?
4. How would you determine the percentage of sulphuric oxide in a sample
of gypsum? What is sulphuric oxide, and what relation does it bear to
sulphur?
5. A mineral contains 20.7 per cent. of water, 32.4 per cent. of lime,
and 18.6 per cent. of sulphur. What is its probable composition? What
experiment would you try to determine the accuracy of your conclusion?
SELENIUM
occurs in nature combined with copper, mercury, and lead, in certain
rare minerals. In small quantities it is found in many ores. It is
detected in solution by the red precipitate produced on boiling the acid
solution with sodium sulphite. This reaction is used for its
determination.
~Solution.~--The solution is effected by boiling with nitric acid or
aqua regia, or by fusing with nitre. To separate the selenium, the
solution is evaporated with an excess of hydrochloric acid and a little
sodium or potassium chloride. This destroys any nitric acid that may be
present, and reduces selenic acid (H_{2}SeO_{4}) to selenious
(H_{2}SeO_{3}). The solution is diluted with water, and treated with a
solution of sulphite of soda. It is warmed, and at last boiled. The
selenium separates as a red precipitate, which (on boiling) becomes
denser and black. It is collected on a weighed filter, washed with hot
water, dried at 100° C., and weighed as pure selenium.
Selenium can be precipitated with sulphuretted hydrogen as a sulphide,
which is readily soluble in ammonium sulphide. This sulphide may be
oxidised with hydrochloric acid and chlorate of potash; and the selenium
separated in the manner described.
TELLURIUM.
Tellurium occurs in nature, native, and in combination with gold,
silver, bismuth and lead. It is sometimes met with in assaying gold
ores. It may be detected by the purple colour it imparts to strong
sulphuric acid when dissolved in the cold, and by the black precipitate
of metallic tellurium which its solutions yield on treatment with a
reducing agent. Telluric acid is reduced to tellurous (with evolution of
chlorine) on boiling with hydrochloric acid.
~Solution~ is effected by boiling with aqua regia, or by fusing with
nitre and sodium carbonate.
~Separation.~--Tellurium closely resembles selenium in its reactions. It
is separated and determined in the same way. Like it, it forms a
sulphide soluble in ammonium sulphide. It is distinguished from selenium
by the insolubility, in a solution of cyanide of potassium, of the metal
precipitated by sodium sulphite; whereas selenium dissolves, forming a
soluble potassic seleno-cyanide.[102]
For the determination, solution is effected by fusing with nitre and
sodium carbonate, dissolving out the tellurate of potash with water, and
boiling with hydrochloric acid. Tellurous compounds are formed, with
evolution of chlorine; and the solution, on treating with a reducing
agent (such as sulphurous acid or stannous chloride), yields metallic
tellurium; which is washed, dried at 100° C., and weighed.
FOOTNOTES:
[101] BaCl_{2} + Na_{2}SO_{4} = BaSO_{4} + 2NaCl.
[102] Se + KCy = KCySe.
CHAPTER XVII.
ARSENIC, PHOSPHORUS, NITROGEN.
ARSENIC.
The chief source of the arsenic of commerce is arsenical pyrites, or
mispickel, which contains about 45 per cent. of arsenic (As). Arsenic
also occurs as a constituent of several comparatively rare minerals;
and, as an impurity, it is very widely distributed. White arsenic is an
oxide of arsenic, and is obtained by roasting arsenical ores, and
refining the material (crude arsenic), which condenses in the flues.
Arsenic itself is volatile, and many of its compounds have the same
property. It forms two well-defined series of salts, corresponding to
the oxides: arsenious oxide (As_{2}O_{3}), and arsenic oxide
(As_{2}O_{5}). These combine with bases to form arsenites and arsenates
respectively. Boiling with nitric acid converts the lower into the
higher oxide; and powerful reducing-agents, such as cuprous chloride,
have the opposite effect.
Arsenic may be detected by dissolving the substance in hydrochloric
acid, or in aqua regia (avoiding an excess of nitric acid), and adding a
little of this solution to the contents of a small flask in which
hydrogen is being made by the action of zinc and hydrochloric acid. The
ignited jet of hydrogen assumes a blue colour if arsenic is present, and
a cold porcelain dish held in the flame (fig. 67) becomes coated with a
dark deposit of metallic arsenic. Antimony produces a similar effect,
but is distinguished by the insolubility of its deposit in a cold
solution of bleaching-powder.
[Illustration FIG. 67.]
Arsenites are distinguished by the volatility of the chloride; by
decolorising a solution of permanganate of potassium, and by immediately
giving a yellow precipitate with sulphuretted hydrogen. Arsenates are
distinguished (after converting into soda salts by boiling with
carbonate of soda and neutralising) by giving with nitrate of silver a
red precipitate, and with "magnesia mixture" a white crystalline one.
~Dry Assay.~--There is no dry assay which is trustworthy. The following
method is sometimes used to find the proportion of arsenious oxide in
"crude arsenic":--Weigh up 5 grams of the dried sample, and place them
in a clean dry test-tube about 6 inches long. Tie a small filter-paper
over the mouth of the tube, so as to prevent air-currents. Heat the tube
cautiously so as to sublime off the white arsenic into the upper part of
the tube. Cut off the bottom of the test-tube by wetting whilst hot.
Scrape out the arsenic and weigh it. The weight gives an approximate
idea of the quantity, and the colour of the quality, of the white
arsenic obtainable from the sample. Some workers (sellers) weigh the
residue, and determine the white arsenic by difference. In determining
the percentage of moisture in these samples, the substance is dried on a
water-bath or in a water-oven.
WET METHODS.
~Solution.~--Where, as in crude arsenic, the substance is arsenious
oxide (As_{2}O_{3}) mixed with impurities, the arsenic is best got into
solution by warming with caustic soda, and neutralising the excess with
hydrochloric acid; it will be present as sodium arsenite. Metals and
alloys are acted on by means of nitric acid; or the arsenic may be at
the same time dissolved and separated by distilling with a strongly-acid
solution of ferric chloride, in the way described under _Volumetric
Methods_.
With minerals, mattes, &c., solution is thus effected:--The
finely-powdered substance is mixed (in a large platinum or porcelain
crucible) with from six to ten times its weight of a mixture of equal
parts of carbonate of soda and nitre. The mass is then heated gradually
to fusion, and kept for a few minutes in that state. When cold, it is
extracted with warm water, and filtered from the insoluble residue. The
solution, acidified with nitric acid and boiled, contains the arsenic as
sodium arsenate. With mispickel, and those substances which easily give
off arsenic on heating, the substance is first treated with nitric acid,
evaporated to dryness, and then the residue is treated in the way just
described.
When the arsenic is present as arsenite or arsenide, distillation with
an acid solution of ferric chloride will give the whole of the arsenic
in the distillate free from any metal except, perhaps, tin as stannic
chloride. With arsenates, dissolve the substance in acid and then add an
excess of soda. Pass sulphuretted hydrogen into the solution; warm, and
filter. Acidulate the filtrate, and pass sulphuretted hydrogen. Decant
off the liquid through a filter, and digest the precipitate with ammonic
carbonate; filter, and re-precipitate with hydrochloric acid and
sulphuretted hydrogen. Allow to stand in a warm place, and filter off
the yellow sulphide of arsenic. Wash it into a beaker, clean the
filter-paper (if necessary) with a drop or two of dilute ammonia;
evaporate with 10 c.c. of dilute nitric acid to a small bulk; dilute;
and filter off the globules of sulphur. The filtrate contains the
arsenic as arsenic acid.
GRAVIMETRIC METHOD.
Having got the arsenic into solution as arsenic acid, and in a volume
not much exceeding 50 c.c., add about 20 c.c. of dilute ammonia and 20
c.c. of "magnesia mixture." Stir with a glass rod, and allow to settle
overnight. Filter, and wash with dilute ammonia, avoiding the use of
large quantities of wash water. Dry, transfer the precipitate to a
Berlin crucible, and clean the filter-paper thoroughly. Burn this paper
carefully and completely; and add the ash to the contents of the
crucible, together with 4 or 5 drops of nitric acid. Evaporate with a
Bunsen burner, and slowly ignite, finishing off with the blow-pipe or
muffle. Cool, and weigh. The ignited precipitate is pyrarsenate of
magnesia (Mg_{2}As_{2}O_{7}), and contains 48.4 per cent. of arsenic
(As).
Instead of igniting the precipitate with nitric acid, it may be
collected on a weighed filter-paper, dried at 100° C., and weighed as
ammonic-magnesic arsenate (2AmMgAsO_{4}.H_{2}O), which contains 39.5 per
cent. of arsenic. The results in this case are likely to be a little
higher. The drying is very tedious, and is likely to leave behind more
water than is allowed for in the formula. In a series of determinations
in which the arsenic was weighed in both forms, the results were:--
Ammonic-magnesic Arsenic Magnesium Pyrarsenate Arsenic
Arsenate in grams. in grams. in grams. in grams.
0.0080 0.0032 0.0065 0.0031
0.0400 0.0158 0.0330 0.0160
0.0799 0.0316 0.0633 0.0306
0.1600 0.0632 0.1287 0.0623
0.4000 0.1580 0.3205 0.1551
0.7990 0.3156 0.6435 0.3114
VOLUMETRIC METHODS.
There are two methods: one for determining the arsenic in the lower, and
the other in the higher state of oxidation. In the first-mentioned
method this is done by titrating with a standard solution of iodine; and
in the latter with a solution of uranium acetate. Where the arsenic
already exists as arsenious oxide, or where it is most conveniently
separated by distillation as arsenious chloride, the iodine method
should be used; but when the arsenic is separated as ammonic-magnesic
arsenate or as sulphide, the uranium acetate titration should be
adopted.
IODINE PROCESS.
This is based on the fact that sodium arsenite in a solution containing
an excess of bicarbonate of soda is indirectly oxidised by iodine to
sodium arsenate,[103] and that an excess of iodine may be recognised by
the blue colour it strikes with starch. The process is divided into two
parts--(1) the preparation of the solution, and (2) the titration.
~Preparation of the Solution.~--For substances like crude arsenic, in
which the arsenic is present as arsenious oxide, the method is as
follows:--Take a portion which shall contain from 0.25 to 0.5 gram of
the oxide, place in a beaker, and cover with 10 c.c. of sodic hydrate
solution; warm till dissolved, put a small piece of litmus paper in the
solution, and render acid with dilute hydrochloric acid. Add 2 grams of
bicarbonate of soda in solution, filter (if necessary), and dilute to
100 c.c. The solution is now ready for titrating.
[Illustration: FIG. 68.]
Where the arsenic has to be separated as arsenious chloride, the process
is as follows:[104]--Weigh up 1 gram of the finely-powdered ore (metals
should be hammered out into a thin foil or be used as filings), and
place in a 16-ounce flask provided with a well-fitting cork, and
connected with a ~U~-tube, as shown in the drawing (fig. 68). The
~U~-tube should contain 2 or 3 c.c. of water, and is cooled by being
placed in a jar or large beaker of cold water. The water used for
cooling should be renewed for each assay.
Pour on the assay in the flask 50 c.c. of a "ferric chloride mixture,"
made by dissolving 600 grams of calcium chloride and 300 grams of ferric
chloride in 600 c.c. of hydrochloric acid, and making up to 1 litre with
water.
Firmly cork up the apparatus, and boil over a small Bunsen-burner flame
for fifteen or twenty minutes, but avoid evaporating to dryness.
Disconnect the flask, and pour away its contents at once to prevent
breakage of the flask by their solidification. The arsenic will be
condensed in the ~U~-tube, together with the greater part of the
hydrochloric acid; transfer the distillate to a beaker washing out the
tube two or three times with water; add a small piece of litmus paper;
neutralise with ammonia; render faintly _acid_ with dilute hydrochloric
acid; add 2 grams of bicarbonate of soda in solution; and dilute to 250
c.c. The solution is now ready for titrating.
The arsenic comes over in the early part of the distillation, as will be
seen from the following experiment, made on 1 gram of copper
precipitate; in which experiment the distillate was collected in
separate portions at equal intervals, and the arsenic in each portion
determined:--
Time Iodine Equivalent to Arsenic
Distilling. Required. in the Distillate.
5 minutes 12.0 c.c. 0.0450 gram
5 " 0.17 " 0.0005 "
5 " 0.0 "
5 " 0.0 "
To dryness 0.0 "
The volume of each distillate was about 5 c.c.
In this operation the metals are converted into chlorides by the action
of ferric chloride, which gives up a part of its chlorine, and becomes
reduced to the ferrous salt. The calcium chloride does not enter into
the chemical reaction, but raises the temperature at which the solution
boils, and is essential for the completion of the distillation.[105] Two
experiments with material containing 3.48 per cent. of arsenic gave--(1)
with ferric chloride alone, 2.74 per cent.; and (2) with the addition of
calcium chloride, 3.48 per cent.
It is always necessary to make a blank determination with 1 gram of
electrotype copper, to find out the amount of arsenic in the ferric
chloride mixture.[106] Unfortunately, a correction is always required.
This amounts to about 0.15 per cent. of arsenic on each assay, even when
the mixture has been purified; and this constitutes the weakness of the
method, since, in some cases, the correction is as much as, or even
greater than, the percentage to be determined.
The acid distillate containing the arsenious chloride may be left for an
hour or so without much fear of oxidation; but it is safer to neutralise
and then to add the bicarbonate of soda, as the following experiments
show. Several portions of a solution, each having a bulk of 100 c.c.,
were exposed for varying lengths of time, and the arsenic in each
determined.
_____________________________________________________________________
| | | |
| | Acid Solutions. | Neutralised Solutions. |
| Time Exposed. | "Iodine" Arsenic Found. |"Iodine" Arsenic Found. |
| | Required. | Required |
|_______________|___________________________|_________________________|
| | | |
| -- | 18.2 c.c. = 0.0136 gram | 18.1 c.c. = 0.0136 gram |
| 1 hour | 18.2 " = 0.0136 " | 18.2 " = 0.0136 " |
| 2 hours | 17.7 " = 0.0133 " | 18.0 " = 0.0135 " |
| 4 " | 17.5 " = 0.0131 " | 18.4 " = 0.0138 " |
| 5 " | 17.0 " = 0.0127 " | 18.3 " = 0.0137 " |
|_______________|___________________________|_________________________|
~The Titration.~--Make a _standard solution of iodine_ by weighing up in
a beaker 16.933 grams of iodine and 30 grams of potassium iodide in
crystals; add a few c.c. of water, and, when dissolved, dilute to 1
litre: 100 c.c. will equal 0.500 gram of arsenic.
A solution of starch similar to that used in the iodide-copper assay
will be required. Use 2 c.c. for each assay. Variations in the quantity
of starch used do not interfere; but the solution must be freshly
prepared, as after seven or eight days it becomes useless.
To standardise the iodine solution, weigh up 0.3 gram of white arsenic;
dissolve in caustic soda; neutralise; after acidulating, add 2 grams of
bicarbonate of soda and 2 c.c. of the starch solution, and dilute to 200
c.c. with cold water. Fill a burette having a glass stop-cock with the
iodine solution, and run it into the solution of arsenic, rapidly at
first, and then more cautiously, till a final drop produces a blue
colour throughout the solution. Calculate the standard in the usual way.
White arsenic contains 75.76 per cent. of arsenic.
The following experiments show the effect of variation in the conditions
of the titration:--
Make a solution of arsenic by dissolving 6.60 grams of white arsenic in
100 c.c. of sodic hydrate solution; render slightly acid with
hydrochloric acid; add 10 grains of bicarbonate of soda, and dilute to 1
litre: 100 c.c. will contain 0.50 gram of arsenic.
~Effect of Varying Temperature.~--The reaction goes on very quickly in
the cold, and, since there is no occasion for heating, all titrations
should therefore be carried out cold.
~Effect of Varying Bulk.~--In these experiments, 20 c.c. of arsenic
solution were taken, 2 grams of bicarbonate of soda and 2 c.c. of starch
solution added, and water supplied to the required bulk. The results
were:--
Bulk 50.0 c.c. 100.0 c.c. 250.0 c.c. 500.0 c.c.
"Iodine" required 20.0 " 20.0 " 20.0 " 20.0 "
Considerable variation in bulk does not interfere.
~Effect of Varying Bicarbonate of Soda.~--This salt must be present in
each titration in considerable excess, to prevent the interference of
free acid. The bicarbonate must be dissolved without heating, as neutral
carbonates should be avoided.
Bicarbonate added 1 gram 2 grams 5 grams 10 grams
"Iodine" required 20.1 c.c. 20.0 c.c. 20.1 c.c. 20.0 c.c.
These results show that large variation in the quantity of bicarbonate
has no effect.
~Effect of Free Acid.~--In these experiments, the arsenic taken, the
starch, and the bulk were as before, but no bicarbonate was added. In
one case the solution was rendered acid with 5 c.c. of acetic acid, and
in the other with 5 c.c. of hydrochloric acid; in both cases the
interference was strongly marked, and no satisfactory finishing point
could be obtained. This was much more marked with the hydrochloric acid.
~Effect of Foreign Salts.~--The process for getting the arsenic into
solution will exclude all metals except tin, but the solution will be
charged with sodium or ammonium salts in the process of neutralising, so
that it is only necessary to see if these cause any interference. The
alkaline hydrates, including ammonia, are plainly inadmissible, since no
free iodine can exist in their presence. Monocarbonates similarly
interfere, but to a much less extent; hence the necessity for rendering
the assay distinctly acid before adding the bicarbonate of soda.
With 20 c.c. of arsenic solution; and with bulk, soda, and starch as
before, the results obtained were:--
"Iodine" required.
With 20 grams of ammonic chloride 20.0 c.c.
" 20 grams of sodium chloride 20.0 "
" 20 grams of sodium acetate 20.0 "
" 0.050 gram of tin, as stannic chloride 19.6 "
Without any addition 20.0 "
The interference of the stannic salt is probably mechanical, the
precipitate carrying down some arsenious acid.
~Effect of Varying Arsenic.~--With bulk, starch, and soda as before, but
with varying arsenic, the results were:--
Arsenic added 1.0 c.c. 10.0 c.c. 20.0 c.c. 50.0 c.c. 100.0 c.c.
"Iodine" required 1.1 " 9.9 " 20.0 " 50.0 " 100.0 "
~Determination of Arsenic in Metallic Copper.~--Put 1 gram of the copper
filings, freed from particles of the file with a magnet, into a
16-oz.-flask; and distil with the ferric chloride mixture, as above
described. Neutralise the distillate; acidify; add bicarbonate of soda
and starch; dilute; and titrate with the standard solution of
iodine.[107] Make a blank determination with 1 gram of electrotype
copper, proceeding exactly as with the assay; and deduct the amount of
arsenic found in this experiment from that previously obtained.
Working in this way on a copper containing 0.38 per cent. of arsenic and
0.80 per cent. of antimony, 0.38 per cent. of arsenic was found.
~Determination of White Arsenic in Crude Arsenic.~--Weigh out 1 gram of
the dried and powdered substance (or 0.5 gram if rich), and digest with
10 c.c. of a 10 per cent. solution of soda; dilute to about 50 c.c., and
filter. Render faintly acid with hydrochloric acid, and filter (if
necessary); add 2 or 3 grams of bicarbonate of soda in solution, then 5
c.c. of starch, and titrate the cold solution with the standard solution
of iodine.
The following is an example:--
1 gram of crude arsenic required 53.7 c.c. "Iodine;"
100 c.c. "Iodine" = 0.6000 gram white arsenic;
100 : 53.7 :: 0.6 : 0.3222, or 32.2 per cent.
With the test-tube method of dry assaying, this same sample gave results
varying from 33 to 35 per cent. of white arsenic, which (judging from
its appearance) was impure.
URANIC ACETATE PROCESS.
This may be looked upon as an alternative to the gravimetric method. It
is applicable in all cases where the arsenic exists in solution as
arsenic acid or as arsenate of soda. The process may be considered in
two parts: (1) the preparation of the solution, and (2) the titration.
~Preparation of the Solution.~--If the arsenic has been separated as
sulphide, it is sufficient to attack it with 10 or 15 c.c. of nitric
acid, and to heat gently till dissolved, avoiding too high a temperature
at first. Afterwards continue the heat till the separated sulphur runs
into globules, and the bulk of the acid has been reduced to 3 or 4 c.c.
Dilute with 20 or 30 c.c. of water; put in a piece of litmus paper; and
add dilute ammonia until just alkaline. Then add 5 c.c. of the sodium
acetate and acetic acid solution (which should make the solution
distinctly acid); dilute to 150 c.c., and heat to boiling. The solution
is ready for titrating.
When the arsenic exists in a nitric acid solution mixed with much
copper, it is separated in the way described under _Examination of
Commercial Copper_ (Arsenic and Phosphorus), pages 208, 209.
If the arsenic has been separated as ammonium-magnesium arsenate, and
phosphates are known to be absent; dissolve the precipitate (after
filtering, but without washing) in dilute hydrochloric acid. Add dilute
ammonia till a slight precipitate is formed, and then 5 c.c. of the
sodium acetate and acetic acid solution; dilute to 150 c.c., and heat to
boiling. Titrate.
If phosphates are present (which will always be the case if they were
present in the original substance, and no separation with sulphuretted
hydrogen has been made), the phosphorus will count in the subsequent
titration as arsenic (one part of phosphorus counting as 2.4 parts of
arsenic). It will be necessary to dissolve the mixed arsenate and
phosphate of magnesia in hydrochloric acid. Add about four or five times
as much iron (as ferric chloride) as the combined phosphorus and arsenic
present will unite with, and separate by the "basic acetate" process as
described under PHOSPHORUS in the _Examination of Commercial Copper_,
page 209. Obviously, when phosphates are present, it is easier to
separate the arsenic as sulphide than to precipitate it with the
"magnesia mixture."
~The Titration.~--The _standard solution of uranium acetate_ is made by
dissolving 34.1 grams of the salt (with the help of 25 c.c. of acetic
acid) in water; and diluting to 1 litre. The water and acid are added a
little at a time, and warmed till solution is effected; then cooled, and
diluted to the required volume: 100 c.c. will equal 0.50 gram of
arsenic.
The _sodic acetate and acetic acid solution_ is made by dissolving 100
grams of sodic acetate in 500 c.c. of acetic acid, and diluting with
water to 1 litre. Five c.c. are used for each assay.
The solution of potassic ferrocyanide used as _indicator_ is made by
dissolving 10 grams of the salt in 100 c.c. of water.
To standardise the solution of uranium acetate, weigh up a quantity of
white arsenic (As_{2}O_{3}) which shall be about equivalent to the
arsenic contained in the assay (0.1 or 0.2 gram); transfer to a flask,
and dissolve in 10 c.c. of nitric acid with the aid of heat. Evaporate
to a small bulk (taking care to avoid the presence of hydrochloric
acid); dilute with water; add a small piece of litmus paper; render
faintly alkaline with ammonia; then add 5 c.c. of the sodic acetate
mixture; dilute to 150 c.c.; and heat to boiling.
Fill an ordinary burette with the uranium acetate solution, and run into
the assay a quantity known to be insufficient. Again heat for a minute
or two. Arrange a series of drops of the solution of ferrocyanide of
potassium on a porcelain slab, and, with the help of a glass rod, bring
a drop of the assay solution in contact with one of these. If no colour
is produced, run in the uranium acetate, 1 c.c. at a time, testing after
each addition, till a brown colour is developed. It is best to overdo
the assay, and to count back. It is not necessary to filter off a
portion of the assay before testing with the "ferrocyanide," since the
precipitate (uranic arsenate) has no effect.
The following experiments show the effect of variation in the conditions
of titration. Make a solution of arsenic acid by dissolving 4.95 grams
of arsenious acid (As_{2}O_{3}) in a covered beaker with 35 c.c. of
nitric acid; evaporate down to 7 or 8 c.c.; and dilute with water to 1
litre: 100 c.c. will contain 0.375 gram of arsenic. Use 20 c.c. for each
experiment.
~Effect of Varying Temperature.~--It is generally recommended to titrate
the boiling solution, since it is possible that the precipitation is
only complete on boiling. Low results are obtained in a cold solution,
the apparent excess of uranium acetate striking a colour at once; on
boiling, however, it ceases to do so; consequently, the solution should
always be boiled directly before testing.
In four experiments made in the way described, but with 20 c.c. of a
solution of arsenic acid stronger than that given (100 c.c. = 0.5 gram
As), the results at varying temperatures were:--
Temperature 15° C. 30° C. 70° C. 100° C.
"Uranium" required 18.0 c.c. 18.5 c.c. 18.5 c.c. 18.7 c.c.
~Effect of Varying Bulk.~--These experiments were like those last
mentioned, but were titrated boiling, and the volume was varied:--
Bulk 50.0 c.c. 100.0 c.c. 200.0 c.c. 300.0 c.c.
"Uranium" required 14.0 " 14.0 " 14.5 " 15.0 "
Considerable variations in bulk are to be avoided.
~Effect of Varying Sodium Acetate.~--These experiments were carried out
like those last noticed, but the bulk was 150 c.c., and varying amounts
of sodic acetate were added in excess of the quantity used in the
experiments previously described:--
Sodic acetate added 0 gram 1 gram 10 grams 20 grams
"Uranium" required 14.5 c.c. 14.5 c.c. 16.0 c.c. 18.0 c.c.
It is evidently important that the quantity of this salt present in each
titration be measured out, so as to avoid variation.
~Effect of Varying the Sodium Acetate and Acetic Acid Solution.~--Acetic
Acid also affects the results, but in the opposite direction, by
preventing the precipitation of uranium arsenate. With varying volumes
of the solution now under notice, the results were:--
Solution added 0.0 c.c. 5.0 c.c. 10.0 c.c. 15.0 c.c.
"Uranium" required 14.5 " 14.5 " 14.5 " 14.0 "
Solution added 20.0 " 30.0 " 40.0 " 50.0 "
"Uranium" required 13.2 " 10.0 " 6.0 " 2.0 "
These show that the quantity ordered (5 c.c.) must be adhered to.
~Effect of Foreign Salts.~--In these experiments, 10 grams of the salt
(the effect of which it was desired to determine) were added to a
solution in other respects resembling those previously used:--
Salt added {Ammonic Ammonic Ammonic Magnesium
{sulphate nitrate chloride sulphate
"Uranium" required 15.5 c.c. 15.5 c.c. 15.3 c.c. 15.3 c.c.
Without any addition, 15.0 c.c. were required; and in another
experiment, in which 30 grams of ammonic salts were present, 15.6 c.c.
of uranium solution were required. Such variations in the amount of
ammonic salts as occur in ordinary working are unimportant.
Phosphates, of course, interfere. In fact, the uranium acetate solution
can be standardised by titrating with a known weight of phosphate, and
calculating its equivalent of arsenic. Thus, in an experiment with 0.6
gram of hydric sodic phosphate (Na_{2}HPO_{4}.12H_{2}O), equivalent to
0.05195 gram of phosphorus, or 0.1256 gram of arsenic, 23.25 c.c. of a
solution of uranium acetate were required. The same solution
standardised with white arsenic gave a standard of which 100 c.c. =
0.5333 gram arsenic. On this standard the 0.6 gram of sodic phosphate
should have required 23.5 c.c.
Experiments in which 0.1 gram of bismuth and 0.1 gram of antimony were
present with 0.1 gram of arsenic, showed no interference on the
titration. Ferric or aluminic salts would remove their equivalent of
arsenic, and, consequently, must be removed before titrating.
~Effect of Varying Arsenic.~--Varying amounts of metallic arsenic were
weighed up and dissolved in nitric acid, &c., and titrated:--
Arsenic taken 0.010 gram 0.050 gram 0.100 gram 0.200 gram
Arsenic found 0.010 " 0.050 " 0.100 " 0.197 "
These experiments show that the method yields good results within these
limits.
~Determination of Arsenic in Mispickel.~--Weigh up 1 gram of the dried
and powdered ore, and evaporate to near dryness with 20 c.c. of dilute
nitric acid. Make up to 100 c.c. with water, and pass sulphuretted
hydrogen to reduce the ferric iron to the ferrous state, then add 20
c.c. of dilute ammonia, and again pass sulphuretted hydrogen. Warm,
filter, and evaporate the filtrate to drive off the excess of ammonia;
then add 10 c.c. of nitric acid, and boil down till the sulphide of
arsenic at first precipitated is dissolved; neutralise; add 5 c.c. of
sodium acetate and acetic acid solution; transfer to a pint flask, boil,
and titrate.
For example, an impure sample of ore required, in duplicate assay of
half a gram each, when treated in the above-mentioned way, 39.6 and 39.5
c.c. of the uranium acetate solution (100 c.c. = 0.537 gram of arsenic),
equivalent to 0.2114 gram of arsenic, or 42.3 per cent.
An alternative method is as follows. Powder the ore very finely and
weigh up .5 gram. Place in a 2-3/4 inch berlin dish and add strong
nitric acid, one drop at a time until the action ceases; with care there
need be no very violent reaction. Dry over a water bath. Cover with 2
grams of nitre and over this spread 5 grams of a mixture of equal parts
of nitre and carbonate of soda. Fuse in a muffle or over a large gentle
blow-pipe flame for 4 or 5 minutes. This will spoil the dish. Allow to
cool and boil out in a larger dish with 100 c.c. of water. Filter and
wash into an 8 oz. flask. Acidify the liquor with nitric and boil down
to about 100 c.c. The acid should not be in too large excess, but an
excess is needed to destroy nitrites. Neutralise with soda or ammonia.
Add 5 c.c. of the mixture of sodium acetate and acetic acid. Titrate
with uranium acetate.
~Determination of Arsenic (As) in Crude Arsenic.~--The method given
under the iodine titration simply determines that portion of the arsenic
which is present in the substance as arsenious oxide or white arsenic.
The following method will give the total arsenic in the sample. It would
be incorrect to report this as so much per cent. of arsenious oxide,
although it may be reported as so much per cent. of arsenic equivalent
to so much per cent. of white arsenic, thus:--
Arsenic 30.0 per cent.
Equivalent to white arsenic 39.6 "
The equivalent of white arsenic is calculated by multiplying the
percentage of arsenic by 1.32. The method of determining the percentage
of arsenic is as follows:---Boil 1 gram of the sample with 10 c.c. of
nitric acid. When the bulk of the solution has been reduced to one-half,
and red fumes are no longer evolved, dilute with a little water, and
filter into a flask. Neutralise the filtrate, add 5 c.c. of sodic
acetate solution, boil and filter. The precipitate (ferric arsenate) is
transferred to a small beaker, treated with 5 c.c. of dilute ammonia,
and sulphuretted hydrogen passed through it. The iron sulphide is
filtered off, and the filtrate evaporated with an excess of nitric acid.
When the solution is clear, it is neutralised, and 1 or 2 c.c. of sodic
acetate solution having been added, is then mixed with the first
filtrate. The solution is boiled and titrated.
A sample treated in this way required 49.2 c.c. of the uranium acetate
solution (100 c.c. = 0.537 gram of arsenic), equivalent to 26.4 per
cent.
~Determination of Arsenic in Brimstone.~--Take 10 grams of the
substance, and powder in a mortar; rub up with 10 c.c. of dilute ammonia
and a little water; rinse into a pint flask; pass a current of
sulphuretted hydrogen; and warm on a hot plate for a few minutes.
Filter, acidulate the filtrate with sulphuric acid; filter off the
precipitate; attack it with 10 c.c. of nitric acid; and proceed as in
the other determinations.
PRACTICAL EXERCISES.
1. Mispickel contains 45.0 per cent. of arsenic, to how much white
arsenic will this be equivalent?
2. How would you make a standard solution of iodine so that 100 c.c.
shall be equivalent to 1 gram of white arsenic?
3. What weight of arsenic is contained in 1 gram of pyrarsenate of
magnesia, and what weight of ammonic-magnesic arsenate would it be
equivalent to?
4. The residue, after heating 10 grams of crude arsenic, weighed 0.62
gram. What information does this give as to the composition of the
substance? If another 10 grams of the substance, heated on a water-bath,
lost 0.43 gram, what conclusions would you draw, and how would you
report your results?
5. If a sample of copper contained 0.5 per cent. of arsenic, and 1 gram
of it were taken for an assay, how much standard uranium acetate
solution would be required in the titration?
PHOSPHORUS AND PHOSPHATES.
Phosphorus rarely occurs among minerals except in its highest oxidized
state, phosphoric oxide (P_{2}O_{5}), in which it occurs abundantly as
"rock phosphate," a variety of apatite which is mainly phosphate of
lime. Phosphates of most of the metallic oxides are found. Phosphoric
oxide in small quantities is widely diffused, and is a constituent of
most rocks. Its presence in varying amounts in iron ores is a matter of
importance, since it affects the quality of the iron obtainable from
them.
Phosphorus occurs in alloys in the unoxidized state. It is directly
combined with the metal, forming a phosphide. In this manner it occurs
in meteoric iron. The alloy phosphor-bronze is made up of copper, tin,
zinc, and phosphorus.
Phosphates are mined in large quantities for the use of manure
manufacturers, and for making phosphorus.
Phosphorus and arsenic closely resemble each other in their chemical
properties, more especially those which the assayer makes use of for
their determination. Phosphorus forms several series of salts; but the
phosphates are the only ones which need be considered. Pyrophosphate of
magnesia, which is the form in which phosphoric oxide is generally
weighed, differs from the ordinary phosphate in the proportion of base
to acid. Metaphosphates differ in the same way. If these are present, it
must be remembered they act differently with some reagents from the
ordinary phosphates, which are called orthophosphates. They are,
however, all convertible into orthophosphates by some means which will
remove their base, such as fusion with alkaline carbonates, boiling with
strong acids, &c.[108]
Phosphides are converted into phosphates by the action of nitric acid or
other oxidizing agents. Dilute acids, when they act on the substance,
evolve phosphuretted hydrogen (PH_{3}). The student should be on his
guard against losing phosphorus in this manner.
There is no dry assay for phosphorus. All assays for it are made either
gravimetrically or volumetrically.
The separation of phosphoric oxide is made as follows:--The ore or metal
is dissolved in acid and evaporated, to render the silica insoluble. It
is taken up with hydrochloric acid, diluted with water, and treated with
sulphuretted hydrogen. The filtrate is boiled, to get rid of the excess
of gas, and treated with nitric acid, to peroxidize the iron present. If
the iron is not present in more than sufficient quantity to form ferric
phosphate with all the phosphorus present, some ferric chloride is
added. The iron is then separated as basic acetate. The precipitate will
contain the phosphorus, together with any arsenic acid not reduced by
the sulphuretted hydrogen. The precipitate should have a decided brown
colour. The precipitate is washed, transferred to a flask, and treated
first with ammonia, and then with a current of sulphuretted hydrogen.
The filtrate from this (acidulated with hydrochloric acid, and, if
necessary, filtered) contains the phosphorus as phosphoric acid. This
method is not applicable in the presence of alumina, chromium, titanium,
or tin, if the solution is effected with nitric acid. The precipitate
obtained by the action of nitric acid on tin retains any phosphoric or
arsenic oxide that may be present.
A method of separation more generally applicable and more convenient to
work is based on the precipitation of a yellow phospho-molybdate of
ammonia,[109] by the action of an excess of ammonic molybdate upon a
solution of a phosphate in nitric acid. Dissolve the substance by
treatment with acid, and evaporate to dryness. Take up with 10 c.c. of
nitric acid, and add 20 grams of ammonic nitrate, together with a little
water. Next put in the solution of ammonium molybdate solution in the
proportion of about 50 c.c. for each 0.1 gram of phosphoric oxide judged
to be present. Warm to about 80° C., and allow to stand for an hour.
Filter, and wash with a 10 per cent. solution of ammonic nitrate. It is
not necessary that the whole of the precipitate be placed on the filter;
but the beaker must be completely cleaned. Dissolve the precipitate off
the filter with dilute ammonia, and run the solution into the original
beaker. Run in from a burette, slowly and with stirring, "magnesia
mixture," using about 15 c.c. for each 0.1 gram of phosphoric oxide.
Allow to stand for one hour. The white crystalline precipitate contains
the phosphorus as ammonium-magnesium phosphate.
Phosphate of lead is decomposed by sulphuric acid; the lead is converted
into the insoluble lead sulphate, and the phosphoric acid is dissolved.
Phosphate of copper and phosphate of iron may be treated with
sulphuretted hydrogen; the former in an acid, and the latter in an
alkaline, solution. Phosphate of alumina is generally weighed without
separation of the alumina, since this requires a fusion. In all cases
the aim is to get the phosphoric oxide either free, or combined with
some metal whose phosphate is soluble in ammonia.
Joulie's method of separation is as follows:--One to ten grams of the
sample are treated with hydrochloric acid, and evaporated to dryness
with the addition (if any pyrites is present) of a little nitric acid.
The residue is taken up with hydrochloric acid, cooled, transferred to a
graduated flask, and diluted to the mark. It is then shaken up, filtered
through a dry filter, and a measured portion (containing about 0.05 gram
of phosphoric acid) transferred to a small beaker. Ten c.c. of a
citric-acid solution of magnesia[110] is added, and then an excess of
ammonia. If an immediate precipitate is formed, a fresh portion must be
measured out and treated with 20 c.c. of the citrate of magnesia
solution and with ammonia as before. The beaker is put aside for from
two to twelve hours. The precipitate is then filtered off and washed
with weak ammonia; it contains the phosphorus as ammonium-magnesium
phosphate.
GRAVIMETRIC DETERMINATION.
If the phosphate is not already in the form of ammonic-magnesic
phosphate, it is converted into this by the addition to its solution of
an excess of ammonia and "magnesia mixture." In order to get the
precipitate pure, the "magnesia mixture" is run in gradually (by drops)
from a burette, with constant stirring. A white crystalline precipitate
at once falls, if much phosphorus is present; but, if there is only a
small quantity, it may be an hour or two before it shows itself. The
solution is best allowed to rest for twelve or fifteen hours (overnight)
before filtering. The presence of tartaric acid should be avoided; and
the appearance of the precipitate should be crystalline. The solution is
decanted through a filter, and the precipitate washed with dilute
ammonia, using as little as may be necessary. The precipitate is dried,
transferred to a weighed Berlin or platinum crucible; the filter-paper
is carefully burnt, and its ash added to the precipitate, which is then
ignited, at first gently over a Bunsen burner, and then more strongly
over the blowpipe or in the muffle. The residue is a white mass of
magnesium pyrophosphate containing 27.92 per cent. of phosphorus, or
63.96 per cent. of phosphoric oxide.
VOLUMETRIC METHOD.
Instead of separating and weighing this compound, the phosphoric oxide
in it can be determined by titration. In many cases the ore may be
dissolved and immediately titrated without previous separation. It is
better, however, to carry the separation so far as to get phosphoric
acid, an alkaline phosphate, or the magnesia precipitate. It may then be
prepared for titration in the following way:--The precipitate in the
last case (without much washing) is dissolved in a little hydrochloric
acid, and the solution in any case rendered fairly acid. Dilute ammonia
is added till it is just alkaline, and then 5 c.c. of the sodic acetate
and acetic acid mixture (as described under the Arsenic Assay). This
should yield a clear distinctly-acid solution. It is diluted to 100 or
150 c.c., heated to boiling, and titrated with the uranium acetate
solution, using that of potassic ferrocyanide as indicator.
The _standard solution_ required is made by dissolving 35 grams of
uranium acetate in water with the aid of 25 c.c. of acetic acid, and
diluting to 1 litre.
An _equivalent solution of phosphoric oxide_ is made by dissolving 25.21
grams of crystallised hydric disodic phosphate (HNa_{2}PO_{4}.12H_{2}O)
in water, and making up to 1 litre. 100 c.c. will contain 0.5 gram of
phosphoric oxide (P_{2}O_{5}), or 0.2183 gram of phosphorus. In making
this solution, transparent crystals only must be used. The uranium
acetate solution is only approximately equivalent to this, so that its
exact standard must be determined.