~Arsenic.~--This is determined in a separate portion. Weigh up a portion

of the powder equivalent to 1 gram of the hardhead, place in a pint

flask, and boil with 10 c.c. of nitric acid. When action has ceased add

10 c.c. of the solution of the metallic portion and then hydrochloric

acid (a few drops at a time) till solution is complete. Warm gently in

dissolving, but do not boil. Dilute to about 100 c.c., render alkaline

with ammonia, and add 20 c.c. of yellow ammonium sulphide. Digest at a

gentle heat for about thirty minutes, filter, and wash. Add 50 c.c. of

magnesia mixture, shake well, allow to stand for an hour, filter, and

wash with dilute ammonia. The precipitate is dissolved and then titrated

with uranium acetate, or it is evaporated with nitric acid, ignited, and

weighed as pyrarsenate of magnesia. Calculate the result to arsenic, As.

~Sulphur.~--Weigh up a portion of the powder equivalent to 2 or 3 grams

of the hardhead. Rub up in a mortar with 5 grams of nitre and fuse in a

porcelain dish for ten minutes. Extract with water, add 20 or 30 c.c.

(as the case may be) of the solution of the "metallics." Add 10 grams of

sodic acetate, and ferric chloride until the precipitate turns brown;

dilute with water to half a litre, boil, and titrate with standard baric

chloride, as described under _Sulphur_. Report as sulphur.

A sample of hardhead examined in this way gave--

Sulphur 3.00%

Arsenic 27.10

Tin 22.2

Copper 1.64

Iron 43.2

Cobalt 2.6

------

99.74

~Examination of Tin Slags.~--In tin smelting works the term "slag" is

applied to the unfused portion of the charge. It is made up of unburnt

anthracite and small lumps of slag proper together with some buttons of

metallic tin. This is rarely, if ever, assayed. The slag proper (or, as

it is generally called, "glass") is a silicate of iron, alumina, and

lime, containing from 3 to 7 per cent. of tin. It is thus examined:--The

sample after bruising on an iron plate, is reduced to a very fine powder

by grinding in an agate mortar. In this state it is in most cases

readily decomposed by hydrochloric acid.

~Determination of Tin.~--Where the percentage of tin only is required,

take 2 grams of the powdered slag and well mix with it 20 c.c. of

hydrochloric acid, and heat to boiling. Add 1 c.c. of nitric acid, allow

to stand for fifteen minutes, dilute with water, and filter. Pass a

rapid current of sulphuretted hydrogen for some time. Allow to settle,

and filter. The precipitate, after washing with hot water, is dried, and

gently calcined until the greater part of the sulphur is burnt off. It

is then strongly ignited in the muffle (or over the blowpipe) with the

addition of a small lump of ammonic carbonate. The residue is weighed as

stannic oxide (SnO_{2}); and is calculated to metallic tin by

multiplying by 0.787. The percentage on the slag is calculated in the

usual way.

The tin is always best determined in the examination of slags by a

separate assay carried out in this way. The determination of the other

constituents is thus made:--

~Silica.~--Take 2 grams of the powdered slag and cover them, in a small

evaporating dish, with 20 c.c. of hydrochloric acid; mix well by

stirring with a glass rod; and evaporate to dryness. If (as is generally

the case) tungsten is present the solution will be blue. Take up with 20

c.c. of hydrochloric acid. Add 1 c.c. of nitric acid; and reduce by

boiling to about half the bulk. Add about 20 c.c. of water, boil, and

filter. Wash the residue with hot dilute hydrochloric acid. It consists

of silica with the tungstic acid. Wash it back into the dish; and digest

with 5 or 10 c.c. of a cold solution of ammonic carbonate. Filter; and

collect the filtrate and washings in a weighed porcelain dish. Dry the

residue, ignite strongly, and weigh as silica, SiO_{2}. In certain

exceptional cases this may contain some unaltered cassiterite, which is

easily recognised by its appearance.

~Tungsten.~--The ammonic carbonate filtrate from the silica is

evaporated to dryness, ignited strongly over the blowpipe, and weighed.

The residue is tungstic acid, WO_{3}. The tungsten may be conveniently

reported in this form, although it is probably present as a lower

oxide.

~Tin.~--The acid filtrate from the silica and tungstic acid is treated

with sulphuretted hydrogen. The sulphide of tin is filtered off. Since

the percentage of tin has been already determined, this precipitate may

be neglected; or may be treated in the same way as the previous one, so

as to check the result. Since some stannic chloride will have been lost

in the evaporation, a low result may be expected. The tin should be

reported as stannous oxide; and is calculated by multiplying the

percentage of tin by 1.136.

The filtrate from the tin is boiled rapidly down to remove sulphuretted

hydrogen; and then peroxidised with 1 or 2 c.c. of nitric acid. It is

cooled, transferred to a graduated flask, and diluted with water to 200

c.c.

~Ferrous Oxide and Alumina.~--Half the filtrate from the tin (that is,

100 c.c.) is taken, nearly neutralised with soda, and treated with

sodium acetate. The basic acetate precipitate obtained on boiling is

filtered off and washed. Reserve the filtrate. The precipitate is

dissolved off the filter with hot dilute hydrochloric acid; and the

solution thus formed is treated with a slight excess of ammonia, and

boiled. The precipitate is filtered off, washed with hot water, dried,

ignited, and weighed as mixed ferric oxide and alumina. The ignited

precipitate is then dissolved with sulphuric and hydrochloric acids; and

the iron determined in the solution by titration with the solution of

stannous chloride. The iron found is calculated to and reported as

ferrous oxide, FeO (factor = 1.286). To find the alumina, which is best

estimated by difference, multiply the iron by 1.428 to get the weight of

ferric oxide, and deduct this from the weight of alumina and ferric

oxide found. This, of course, gives the alumina. A direct determination

may be made by removing the tin from the titrated solution with

sulphuretted hydrogen, filtering, nearly neutralising with ammonia, and

boiling with a few grams of hyposulphite of soda. The precipitate,

filtered, washed, and ignited, is the alumina, which is weighed. The

direct determination gives a slightly low result.

~Oxides of Zinc and Manganese.~--These are determined in the filtrate

from the basic acetate precipitate by rendering alkaline with ammonia,

and passing a current of sulphuretted hydrogen. Generally a small, but

decided, precipitate of alumina comes down, together with sulphides of

any zinc or manganese which is present. The precipitate is allowed to

settle, dried, ignited, and weighed. The metals are separately

determined in it; and the residue is counted as alumina, and added to

that already found. The mixed precipitate amounts to from 1 to 2 per

cent. of the sample.

~Lime.~--The filtrate from the last is treated with ammonic oxalate,

boiled for a few minutes, allowed to settle, and filtered. The

precipitate is washed with hot water; dried; ignited; and weighed as

carbonate, after gentle ignition; or as lime, after strong ignition in

the muffle.

~Magnesia.~--The filtrate from the lime is treated with sodic phosphate

and ammonia. It is well mixed by stirring, and allowed to stand

overnight. The precipitate is washed with dilute ammonia, dried,

ignited, and weighed as pyrophosphate.

~Soda and Potash.~--These are determined in the remaining half of the

filtrate from the tin. The solution is rendered ammoniacal with ammonia;

and treated, first with sulphuretted hydrogen, and then with ammonium

oxalate. The precipitate is filtered off and rejected. The filtrate is

evaporated in a small porcelain dish over a Bunsen burner, or on the

sand bath; and towards the close (or earlier if the evaporation is not

proceeding well) nitric acid is added. The evaporation is carried to

dryness; and the residue heated nearly to redness. The residue, which

consists of magnesia with carbonates and chlorides of the alkalies, is

extracted with water; and filtered. The filtrate is evaporated with

hydrochloric acid in a weighed platinum dish, ignited gently, and

weighed. This gives the weight of the mixed chlorides of sodium and

potassium; which are then separated and determined as described under

_Potash_.

It must be remembered when calculating the percentage that (with the

exception of the silica, tungstic acid, and tin) the determinations have

been made on 1 gram of the sample.

The following analysis will illustrate the composition of such a slag:--

Tungstic acid 1.3%

Silica 39.4

Stannous oxide 8.1

Ferrous oxide 26.2

Alumina 14.8

Oxide of manganese traces

Lime 7.9

Magnesia 0.5

Alkalies calculated as soda 1.7

-----

99.9

TITANIUM.

Titanium only occurs as a mineral in its oxidised state, or as titanic

oxide (TiO_{2}). It is a substance which has little commercial value,

and is generally recognised as one of the rare bodies; although, in

small quantities, it is widely disseminated. It occurs in granite,

basalt, and other igneous rocks in quantities up to as much as 1 per

cent. It is also met with in clays and iron ores, and in river sands, in

which it is often associated with stream tin. The proper minerals of

titanium are rutile (TiO_{2}), titaniferous iron (titanate of iron), and

sphene (titanate and silicate of lime).

The oxide of titanium (like cassiterite and quartz) is undecomposed by

hydrochloric or nitric acid; so that it is generally found in the

residue insoluble in acids. The titanates, however, are attacked, and a

portion of the titanium dissolves; so that it must be looked for in both

the filtrate and residue. Oxide of titanium in its native form, or after

ignition, may be made soluble by fusing the finely-divided substance

with fusion mixture in a platinum dish. The resulting titanate is

dissolved out of the "melt" by cold hydrochloric acid.

The method most commonly used is fusion with bisulphate of potash. This

renders the oxide of titanium soluble in cold water. The process is as

follows:--The substance is extracted with hydrochloric and nitric acids,

and the solution reserved for further treatment; the residue is dried,

moistened with sulphuric acid, and evaporated once or twice to dryness

with hydrofluoric acid. It is then fused with bisulphate of potash, and

the "melt" extracted with cold water until all soluble matter is

removed. The solution is filtered. The residue may consist of unremoved

silica, and oxides of tantalum, niobium, and, perhaps, chromium. On the

prolonged boiling of the filtrate, the oxide of titanium (and oxide of

zirconium, if any) is precipitated.

Any titanium dissolved by the first extraction with acids is recovered

in the following way:--Sulphuretted hydrogen is passed into the acid

solution, and any precipitate that may be formed is filtered off. The

filtrate is oxidised, and the iron, aluminium, and titanium are

separated as basic acetates (see under _Iron_). The precipitate is dried

and fused with bisulphate of potash. The "melt" is extracted with cold

water, filtered if necessary, and the solution rendered first faintly

alkaline with ammonia, then very slightly acid with sulphuric acid. 30

or 40 c.c. of a saturated solution of sulphurous acid is added, and the

oxide of titanium precipitated by prolonged boiling. It is filtered off,

added to the precipitate previously got, ignited with ammonic carbonate

towards the end, and then weighed.

~Detection.~--Titanium is detected in an insoluble residue by fusing the

residue for some time in a bead of microcosmic salt. In the reducing

flame it gives a violet colour, which becomes reddish-brown if much iron

is present. In the oxidising flame it gives a colourless or whitish

bead. It is best detected in acid solutions by the deep brown or iodine

colour developed on adding hydroxyl. A solution of this can be prepared

by pouring peroxide of barium (BaO_{2}) diffused in water into dilute

hydrochloric acid (a little at a time), and keeping the acid in excess.

~Separation.~--In the usual course of an analytical separation the

hydrate of titanium will be thrown down with ferric hydrate, &c., on the

addition of ammonic chloride and ammonia. It is best separated from this

precipitate by fusion with bisulphate of potash, as already described,

but it must be remembered that the presence of much mineral acid

prevents complete precipitation when the solution is boiled. Further, if

phosphates are present, the precipitate will contain phosphoric oxide;

it may be freed from this by fusion with sodium carbonate. A very good

method of separating titanium from iron is to add tartaric acid and

ammonia to the solution, and then precipitate the iron (as sulphide)

with sulphuretted hydrogen. The filtrate contains the titanium, which is

recovered by evaporating and igniting. It may be separated from zirconia

by the action of sodium carbonate, which precipitates both; but when

concentrated, redissolves the zirconia. The separation from large

quantities of silica is best effected by evaporating with hydrofluoric

acid, which volatilises the silicon; but sulphuric acid must be present,

otherwise some titanium also will be lost, as may be seen from the

following experiments,[77] in which oxide of titanium (pure, ignited)

was evaporated to dryness with a quantity of hydrofluoric acid known by

experiment to be sufficient to volatilise 1 gram of silica.

_Without sulphuric acid_, 0.0466 gram of titanic oxide left 0.0340 gram,

showing a loss of about 25 per cent.

_With sulphuric acid_ the following results were obtained:--

Oxide taken. Left after Evaporation

and Ignition.

0.0340 gram 0.0340 gram

0.0414 " 0.0413 "

0.0520 " 0.0520 "

0.0352 " 0.0352 "

GRAVIMETRIC DETERMINATION.

The titanic hydrate thrown down by ammonia (or on boiling the solution

from the bisulphate) is collected, washed, dried, ignited strongly with

the addition of a little ammonic carbonate, and weighed. The substance

is titanic oxide (TiO_{2}), and is generally reported as such. It

contains 60.98 per cent. of titanium. It should be white, if pure

(Holland), white, yellow, or brown (Fresenius), or black (Tidy).

VOLUMETRIC METHOD.

A method has been proposed based on the reduction of titanic oxide by

zinc in hydrochloric acid solutions to the sesquioxide. The reduction is

marked by the development of a violet or green colour, the former with

chlorides and the latter when fluorides are present. The quantity of

titanium reduced is measured by titrating with permanganate of potassium

solution. The water used must be free from dissolved oxygen.

TUNGSTEN AND TUNGSTATES.

Tungsten occurs in nature only in the oxidised state, or as tungstic

acid (WO_{3}), either free, as in wolframine, or combined with oxides of

manganese and iron, as in wolfram, or with lime, as in scheelite.

Wolfram occurs associated with tin ores, the value of which is

consequently lowered. Both wolfram and scheelite are of considerable

importance as a source of tungstic acid for the manufacture of sodium

tungstate, which is used as a mordant and for some other purposes, and

as a source of metallic tungsten, which is used in steel-making.

The tungsten minerals have a high specific gravity (6 to 7.5). On

treatment with hydrochloric acid or aqua regia they are decomposed; the

yellow tungstic acid separates and remains insoluble.

Tungsten itself is insoluble in nitric acid or aqua regia; but is

converted into tungstic acid (WO_{3}) by prolonged and strong ignition

in air. Alloys containing tungsten leave tungstic acid after treatment

with nitric acid or aqua regia. Tungstic acid may be got into solution

after fusion with alkalies or alkaline carbonates. This solution gives

with hydrochloric acid a white precipitate of tungstic acid, which

becomes yellow on boiling, but the separation is not complete. Fusion

with bisulphate of potash gives a residue, which does not dissolve in

water, but is soluble in ammonic carbonate. For the assay of minerals

containing tungsten these reactions are only occasionally taken

advantage of for testing or purifying the separated tungstic acid.

~Detection.~--The minerals are easily recognised by their physical

characters, and the yellow tungstic acid separated by boiling with acids

is the best test for its presence; this, after decanting and washing,

immediately dissolves in a few drops of dilute ammonia. A solution of

tungstate acidulated with hydrochloric acid becomes intensely blue on

the addition of stannous chloride and warming. Fused in a bead of

microcosmic salt it gives a clear blue colour (reddish-brown if iron is

also present) in the reducing flame, but is colourless in the oxidising

flame.

~Solution and Separation.~--The decomposition and solution of natural

tungstates is difficult to effect owing to the separation of tungstic

acid; the method of treatment is as follows:--Boil the finely-powdered

substance with hydrochloric acid or aqua regia till it apparently ceases

to be attacked; dilute, filter, and wash with dilute hydrochloric acid.

Cover with dilute ammonia, and filter the solution, which contains

ammonic tungstate, into an evaporating dish. Treat the residue again

with acid, and again dissolve out the separated tungstic acid with

ammonia, and repeat this operation until decomposition is complete. By

this means there will be obtained--(1) a solution containing tungstate

of ammonia; (2) an insoluble residue with silicates, and oxides of tin,

niobium, tantalum, &c.; and (3) an acid solution containing the soluble

bases. The tungstate of ammonia requires simple evaporation on the

water-bath and gentle ignition in order to cause the tungstic acid to be

left in an almost pure state; possibly, it may carry a little silica.

GRAVIMETRIC DETERMINATION.

The tungstic acid is dissolved, and separated as ammonic tungstate, and,

after evaporation, is gently ignited, the heat being increased towards

the end. The residual tungstic acid is fixed, so that when the ammonia

has been driven off it may be strongly heated without loss. It is a dark

yellow or brown powder whilst hot, which becomes a light yellow on

cooling. If any reduction has taken place it will be more or less

greenish. It is weighed when cold, and is the trioxide or "tungstic

acid" (WO_{3}), which contains 79.31 per cent. of tungsten. After its

weight has been taken its purity is checked by fusing with hydric

potassic sulphate, extracting with water, and treating the residue with

ammonic carbonate. Any silica present will be left undissolved; it

should be separated and weighed, and its weight deducted from that of

the tungstic acid found.

~Determination of Tungstic Acid in Wolfram.~--Take 2 grams of the

finely-powdered sample and boil with 50 c.c. of hydrochloric acid for

half an hour, adding 5 c.c. of nitric acid towards the end. Allow to

stand overnight and boil again for 15 or 20 minutes; dilute with an

equal volume of water, and filter. Wash with dilute hydrochloric acid,

dissolve in a few c.c. of warm dilute ammonia, and dilute to 200 c.c.

with distilled water; allow to settle, and filter. Evaporate in a

weighed dish, ignite, and weigh.

The following analysis will illustrate the composition of a sample of

Cornish wolfram as brought into the market:--

Tungstic acid 50.1%

Cassiterite 10.9

Ferrous oxide 24.6

Manganous oxide 5.4

Niobic oxide, alumina, &c. 3.5

Silica 1.2

Copper oxide 2.7

Zinc oxide 0.22

Arsenic 0.51

Sulphur 0.20

-----

99.33

NIOBIC AND TANTALIC OXIDES.

These oxides are commonly met with in samples of wolfram and tinstone,

especially niobic. They are probably present in the form of columbite, a

niobate of iron and manganese; and tantalite, a tantalate of the same

metals.

On boiling with hydrochloric acid they are both liberated, and remain

for the greater part (all the niobic) in the insoluble residue with the

tungstic acid. On removing the latter with dilute ammonia they remain as

a white insoluble precipitate, very prone to run through the filter on

washing. They may be dissolved in hydrofluoric acid either at once or

after fusion with bisulphate of potash, and extraction with cold water.

To the solution in hydrofluoric acid gradually add a boiling solution of

acid potassium fluoride (HF, KF.). Potassic fluotantalate (soluble in

200 parts of water) separates out first, and afterwards potassic

fluoniobate (soluble in 12 parts of water). The separated salts (after

heating with sulphuric acid and washing out the potassium sulphate

formed) are ignited with ammonic carbonate, and weighed as tantalic

oxide (Ta_{2}O_{5}) and niobic oxide (Nb_{2}O_{5}) respectively.

They are both white powders. The oxide of niobium dissolved in a bead of

microcosmic salt gives a bluish colour in the reducing flame. The oxide

of tantalum dissolves in the bead, but gives no colour.

FOOTNOTES:

[76] This will give almost the whole of the tin; a further portion will

be got in subsequent work, and must be added to this result.

[77] Published by P. Holland, in the _Chemical News_, vol. lix. p. 27.

CHAPTER XIII.

MANGANESE, CHROMIUM, &c.

MANGANESE.

Manganese occurs mainly as black oxide (MnO_{2}) in the mineral

pyrolusite; and, in a less pure form, in psilomelane and wad. The value

of the ore depends rather on the percentage of available oxygen than on

the proportion of metal present. The results of assays are generally

reported as so much per cent. of the dioxide (MnO_{2}). In smaller

quantities it is very widely distributed. Manganese itself has a value

for steel-making; or, rather, for the making of spiegeleisen and

ferro-manganese, which are used in the Bessemer and Siemens processes.

For this purpose the percentage of the metal (Mn) is required.

Consequently the minerals of manganese may be considered in two

aspects--(1) as a source of oxygen; and (2) as a source of manganese.

These will require separate consideration.

The black oxide is mainly used in the preparation of chlorine,

liberation of which it brings about when treated with hot hydrochloric

acid, or with a mixture of common salt and sulphuric acid. The quantity

of chlorine which is obtained depends upon the proportion of dioxide

present;[78] and in assaying may either be measured by its equivalent of

iodine liberated, or by the oxidising effect on an acid solution of

ferrous sulphate. When the ore also carries substances which have a

reducing effect (such as ferrous compounds), such assays will give, not

the total dioxide (MnO_{2}), but less, by the amount required to oxidise

these impurities; and this is exactly what is required in valuing such

an ore for commercial purposes. Manganese compounds are characterised by

the readiness with which they may be converted into highly-oxidised

bodies. Solution of manganese in hydrochloric acid, rendered alkaline

with ammonia, yields a clear solution,[79] which rapidly takes up oxygen

from the air, forming a brown precipitate of the oxide (Mn_{2}O_{3}).

The addition of bromine or chlorine to such a solution determines the

precipitation of a still higher oxide (approximately MnO_{2}). On

treating a compound containing manganese with nitric acid and dioxide of

lead (PbO_{2}), the oxidation is carried still further, a

purple-coloured solution of permanganic acid (HMnO_{4} or

H_{2}O.Mn_{2}O_{7}) being formed. On fusing minerals containing (even

traces of) manganese with sodium carbonate in an open crucible, a green

"melt" is obtained which owes its colour to sodium manganate

(Na_{2}MnO_{4} or Na_{2}O.MnO_{3}). This salt is soluble in water,

forming a green solution; which, when rendered acid, rapidly changes

into the permanganate with the characteristic purple colour.

Permanganate of potash is a salt much used in assaying, with some

properties of which the student will have already become familiar.

Compounds of manganese, on boiling with strong hydrochloric acid, yield

manganous chloride[80] (MnCl_{2}).

The properties given above serve for the detection of manganese; the

higher oxides are distinguished by causing the evolution of chlorine

(with its peculiarly suffocating smell) when acted on with hydrochloric

acid; while the green "melt," with sodium carbonate, can be relied on

for the recognition of manganese itself. There is no dry assay of

manganese ores.

WET METHODS.

Strong hydrochloric acid is the best solvent for ores of manganese; but

where the proportion of dioxide (MnO_{2}) is required, the solution is

effected during the assay. The ore should be in a very fine state of

division before treatment with acids.

The separation of manganese from other metals is thus effected: Ignite,

in order to destroy any organic matter which may be present; dissolve in

hydrochloric acid, and evaporate to dryness, to separate silica. Take up

with hydrochloric acid, dilute, pass sulphuretted hydrogen, and filter.

Boil off the excess of gas, peroxidise the iron with a drop or two of

nitric acid, and separate the iron as basic acetate (as described under

_Iron_).[81] If the iron precipitate is bulky, it is dissolved in a

little hydrochloric acid, reprecipitated, and the filtrate added to the

original one. Neutralise with soda, and add bromine in excess; heat

gradually to boiling, allow to settle, and filter. The precipitate is

impure dioxide of manganese (containing alkalies and, possibly, cobalt

or nickel).

GRAVIMETRIC DETERMINATION.

Dissolve the precipitate in hydrochloric acid, and boil; add a slight

excess of carbonate of soda, warm, and filter. Wash with hot water, dry,

carefully ignite in an open Berlin crucible, and weigh. The substance is

the brown oxide (Mn_{3}O_{4}), and contains 72.05 per cent. of

manganese. If the percentage of dioxide is required it may be calculated

by multiplying the percentage of manganese by 1.582. It must be borne in

mind that the manganese should never be calculated to dioxide except

when it is known to exist in the ore only in that form.

VOLUMETRIC METHODS.

The two methods are based on the oxidising effect of manganese dioxide;

and if the metal does not already exist in this form it will require a

preliminary treatment to convert it. The following method due to Mr. J.

Pattinson[82] effects this: A quantity of the ore containing not more

than .25 grams of the metal (Mn), is dissolved in hydrochloric acid in a

pint beaker, and, if necessary, 3 or 4 c.c. of nitric acid are added to

peroxidise the iron, and ferric chloride is added if required, so that

there may be at least as much iron as manganese. Calcium carbonate is

added till the solution is slightly red; and next the redness is removed

by the cautious addition of acid; 30 c.c. of zinc chloride solution

(containing 15 grams of zinc per litre) are added, the liquid is brought

to boil and diluted to about 300 c.c. with boiling water; 60 c.c. of a

solution of bleaching powder (33 grams to the litre and filtered),

rendered slightly greenish by acid, are then run in and are followed by

3 grams of calcium carbonate suspended in 15 c.c. of boiling water.

During effervescence the beaker is covered, the precipitate is stirred,

and 2 c.c. of methylated spirit are mixed in. The precipitate is

collected on a large filter, washed with cold water, and then with hot,

till free from chlorine, which is tested for with starch and potassium

iodide. The acid ferrous sulphate solution (presently described) is then

measured into the beaker, and the precipitate, still in the paper,

added; more acid is added (if necessary), and the solution is diluted

and titrated. In place of bleaching powder solution, 90 c.c. of bromine

water (containing 22 grams per litre) may be used.

FERROUS SULPHATE ASSAY.

This method, which is the one commonly used, is based on the

determination of the amount of ferrous iron oxidised by a known weight

of the ore. It is known that 87 parts of the dioxide will oxidise 112

parts of ferrous iron;[83] therefore 1 gram will oxidise 1.287 gram of

ferrous iron, or 1 gram of ferrous iron oxidised will be equivalent to

0.7768 gram of the dioxide. The finely-divided substance containing the

dioxide is digested in a solution of a known quantity of iron in

sulphuric acid. The iron, of course, must be in excess, which excess is

determined when the ore is dissolved by titrating with standard

permanganate or bichromate of potash solution. The assay resolves itself

into one for the determination of ferrous iron, for which the standard

solutions and method of working described under _Iron_ are used.

The assay is as follows:--For rich ores, 2 grams of clean soft iron wire

are treated, in a pint flask, with 100 c.c. of dilute sulphuric acid and

warmed till dissolved. Carefully sample the ore, and in one portion

determine the "moisture at 100° C.;" grind the rest in a Wedgwood mortar

with a little pure alcohol until free from grit. This reduces the

substance to a finely-divided state and assists solution. Evaporate off

the alcohol and dry at 100° C., mix well, and keep in a weighing-bottle.

Weigh up 2 grams and add them to the solution of iron in the flask;

carefully wash it all down into the acid liquid. On rotating the flask

the ore will rapidly dissolve, but gentle heat may be used towards the

end to complete the solution. When the residue is clean and

sandy-looking, and free from black particles, the flask is cooled, and

the residual ferrous iron is determined by titration with

"permanganate." The iron thus found, deducted from the 2 grams taken,

will give the amount of iron peroxidised by the dioxide contained in the

2 grams of ore. This divided by 2 and multiplied by 77.68 will give the

percentage of dioxide in the sample, or multiplied by 49.41 will give

that of metallic manganese.

When the quantity of manganese or of the dioxide to be determined is

small, it is not necessary to use 2 grams of iron; 1 gram, or even less,

may be taken. The iron may be used in the form of a standard solution of

ferrous sulphate and portions measured off, thus saving the labour of

weighing.

~Determination of Dioxide in a Manganese Ore.~--Weigh up 1 or 2 grams of

the finely-powdered ore[84] and an equal weight of pure iron wire,

dissolve the wire in 50 or 100 c.c. of dilute sulphuric acid, and, when

solution is complete, add the ore and warm till it too is dissolved.

Cool and titrate the remaining ferrous iron with the permanganate or

bichromate of potassium solution.

For example, 0.7560 gram of pyrolusite and 1.000 gram of iron were taken

and treated as above; 13.9 c.c. of "permanganate" (standard 100 c.c. =

0.4920 gram iron) were required; this indicates that 0.0684 gram of iron

was left unoxidised by the ore. The iron oxidised, then, was 0.9316 gram

(1.000 - 0.0684); multiplying this by 0.7768, we find that 0.7237 gram

is the quantity of manganese dioxide which was present. This is

equivalent to 95.77 per cent.;

0.7560 : 0.7237 :: 100 : 95.77.

IODINE METHOD.

It has been already stated that when dioxide of manganese is boiled with

strong hydrochloric acid chlorine is given off, and that the amount of

chlorine so liberated is a measure of the dioxide present. If the

chlorine is passed into a solution of potassium iodide, an equivalent of

iodine will be set free.[85] This is apparently a very indirect way of

determining how much of the dioxide is present; but the reactions are

very sharp, and the final determination of the iodine is an easy one.

[Illustration: FIG. 60.]

The finely-powdered sample of dioxide is placed in a small flask

provided with an exit tube leading into a solution of potassic iodide

(fig. 60). On adding hydrochloric acid and boiling, the chlorine evolved

is driven into the iodide solution and there absorbed; the boiling is

continued till the steam and hydrochloric acid fumes have driven the

last portions of the chlorine out of the flask and into the solution. In

this experiment there is a strong tendency for the iodide solution to

rush back into the flask. This tendency is overcome by avoiding draughts

and regulating the heat; or by placing a lump of magnesite in the flask,

which acts by evolving carbonic acid and so producing a steady outward

pressure. When the distillation is finished the tube containing the

iodine is detached and washed out into a beaker. If the solution is

strongly acid it should be almost neutralised by the cautious addition

of dilute ammonia. If crystals of iodine have separated, potassium

iodide must be added in quantity sufficient to dissolve them. The

condenser must be kept cool whilst the chlorine is passing into it.

The solution, transferred to a beaker, is titrated with a standard

solution of sodic hyposulphite (100 c.c. = 1.27 gram iodine or 0.435

gram of dioxide of manganese). In titrating, the solution should be

cold, or not warmer than 30° C. The bulk may vary from 100 to 200 c.c.;

but it is best always to work with the same volume. The "hypo" is run in

with constant agitation until the brown colour has been reduced to a

light yellow; 5 c.c. of starch solution are then added and the titration

cautiously continued until the end is reached; the finish is indicated

by a change from blue to colourless.

The assay solution may be acidified with acetic, sulphuric, or

hydrochloric acid before titrating with "hypo;" but it must be only

faintly so. An excess of acid may be nearly neutralised with ammonia

without interference, but excess of alkali is fatal. Bicarbonate of soda

must not be used in excess; it is best to avoid it altogether. The assay

solution should be titrated at once, as it weakens on standing; and the

"hypo" solution should be standardised every two or three days, as its

strength is not constant.

_The standard solution of hyposulphite of soda_ is made by dissolving 25

grams of the salt (Na_{2}S_{2}O_{3}.5H_{2}O) in water and diluting to 1

litre. 100 c.c. are equivalent to 1.27 gram of iodine.

This solution is standardised by weighing, in a small beaker, about half

a gram of iodine, to which is added a crystal or two of potassium iodide

and a few drops of water. When dissolved, the solution is diluted to 100

c.c., and titrated in the manner described. The starch solution is made

in the manner described under the iodide copper assay. 5 c.c. are used

for each titration.

In determining the effects of variations in the condition of the assay a

solution of iodine was used, which was equivalent in strength to the

"hypo" solution. It was made by dissolving 12.7 grams of iodine with 25

grams of potassium iodide in a little water and diluting to 1 litre. 100

c.c. of this solution were found (at the time of the experiments) to be

equivalent to 102.0 c.c. of the "hypo."

~Effect of Varying Temperature.~--The bulk of the solution was 100 c.c.;

20 c.c. of iodine were taken, and 5 c.c. of starch solution were added

towards the end as indicator. These conditions are also those of the

other experiments, except where otherwise stated. Iodine being volatile,

it is to be expected that with hot solutions low results will be

obtained.

Temperature 15° 20° 40° 60° 80°

"Hypo" required 20.4 c.c. 20.4 c.c. 20.1 c.c. 19.2 c.c. 15.5 c.c.

These show that the temperature should not much exceed 20°.

~Effect of Exposure of the Iodine Solution.~--Twenty c.c. of the iodine

were diluted to 100 c.c., and exposed for varying lengths of time in

open beakers at the ordinary temperature, and then titrated.

Time exposed -- 1 day 2 days 3 days

"Hypo" required 20.4 c.c. 16.1 c.c. 13.6 c.c. 9.4 c.c.

~Effect of Varying Bulk.~--These experiments were carried out in the

usual way, bulk only varying.

Bulk 100.0 c.c. 200.0 c.c. 300.0 c.c. 500.0 c.c.

"Hypo" required 20.4 " 20.4 " 20.4 " 20.4 "

~Effect of Varying Acid.~--These experiments were under the usual

conditions, the bulk being 100 c.c. The results were--

Acetic acid -- 1.5 c.c. 30.0 c.c.

"Hypo" required 20.4 c.c. 20.7 " 20.7 "

Hydrochloric acid -- 1.5 c.c. 15.0 c.c.

"Hypo" required 20.4 c.c. 20.6 " 20.9 "

Sulphuric acid -- 0.5 c.c. 20.0 c.c.

"Hypo" required 20.4 c.c. 20.7 " 15.2 "[86]

Nitric acid -- 0.5 c.c. 10.0 c.c.

"Hypo" required 20.4 c.c. 21.5 " could not be titrated.

In the application of this titration to the assay of manganese ores,

hydrochloric and hydriodic acids are the only ones likely to be present.

~Effect of Alkalies.~--On theoretical grounds the presence of these is

known to be inadmissible. A solution rendered faintly alkaline with

ammonia required only 11.2 c.c. of "hypo;" and another, with 0.5 gram of

caustic soda, required 4.0 c.c. instead of 20.4 c.c. as in neutral

solutions.

~Effect of nearly Neutralising Hydrochloric Acid Solutions with

Ammonia.~--Provided care is taken not to add excess of ammonia, this

has a good effect, counteracting the interference of excess of acid.

Thus 20 c.c. of iodine (as before) required 20.4 c.c. of "hypo;" with 15

c.c. of hydrochloric acid 20.7 c.c. were required, but with 15 c.c. of

acid, nearly neutralised with dilute ammonia 20.4 c.c. were used.

~Effect of the Addition of Starch.~--The addition of varying quantities

of starch has no effect, provided it is added when the titration is

nearly finished, as the following experiments show:--

Starch added 1.0 c.c. 5.0 c.c. 10.0 c.c. 50.0 c.c.

"Hypo" required 20.4 " 20.4 " 20.4 " 20.5 "

But if the starch is added before the titration, the results are liable

to error.

Starch added 1.0 c.c. 50.0 c.c.

"Hypo" required 20.4 " 24.0 "

The starch should be used fresh, and is best made on the day it is used;

after four days the finishing point is not so good.

~Effect of Varying Potassium Iodide.~--An excess of iodide is always

required to keep the iodine in solution; a larger excess has little

effect.

Iodide added -- 1 gram 20 grams

"Hypo" required 20.4 c.c. 20.5 c.c. 20.6 c.c.

The 20 c.c. of iodine used, itself contained 0.5 gram of potassium

iodide.

~Effect of Foreign Salts.~--

Bicarbonate of soda

added -- 0.5 gram 1.5 gram 5.0 grams

"Hypo" required 20.4 c.c. 18.2 c.c. 17.1 c.c. 16.0 c.c.

The solution obviously must be free from bicarbonate of soda. This

should be remembered, since when titrating arsenic assays with iodine it

must be present; and students must avoid confounding the two titrations.

In some other experiments, in which 10 grams each of the salts were

taken, the following results were obtained:--

Salt added -- AmCl AmNO_{3} Am_{2}SO_{4}

"Hypo" required 20.4 c.c. 20.5 c.c. 20.3 c.c. 20.2 c.c.

Salt added NaCl NaNO_{3} Na_{2}SO_{4}

"Hypo" required 20.3 c.c. 20.4 c.c. 20.4 c.c.

~Effect of Varying Iodine.~--

Iodine added 1.0 c.c. 10.0 c.c. 20.0 c.c. 50.0 c.c. 100.0 c.c.

"Hypo" required 1.3 " 10.2 " 20.4 " 51.0 " 102.0 "

~Determination of Dioxide in a Manganese Ore.~--Weigh up 0.25 to 0.3

gram of the powdered ore; place in a flask, cover with 10 c.c. of

hydrochloric acid, and close the flask with a paraffined cork, and bulbs

(as shown in fig. 60), having previously charged the bulb with 5 grams

of potassium iodide in strong solution. Heat the flask, and boil

cautiously for about fifteen minutes. Wash the contents of the bulbs

into a large beaker, nearly (but not quite) neutralise with dilute

ammonia, and titrate with the standard "hypo."

As an example, 0.2675 gram of pyrolusite was taken, and required 60.3

c.c. of standard "hypo" (100 c.c. equal 1.185 gram iodine, or 0.4042

gram MnO_{2}), which equals 0.2437 gram of the dioxide or 91.1 per cent.

COLORIMETRIC METHOD.

When compounds of manganese free from chlorides are boiled with nitric

acid and dioxide of lead,[87] the manganese is converted into

permanganic acid, which is soluble and tints the solution violet. The

depth of colour depends on the amount of manganese present, and this

should not much exceed 10 milligrams. A quantity of substance containing

not more than this amount of manganese should be boiled for a few

minutes with 25 c.c. of a solution containing 5 c.c. of nitric acid, and

10 or 20 c.c. of dilute sulphuric acid, with 2 or 3 grams of lead

dioxide. Filter through asbestos, wash by decantation with dilute

sulphuric acid, make up with distilled water[88] to a definite bulk, and

take a measured portion for the colorimetric determination.

The standard solution of manganese is made by dissolving 0.1435 gram of

permanganate of potash (KMnO_{4}) in a little water acidulated with

nitric acid, and diluting to 1 litre. One c.c. will contain 0.05

milligram of manganese.

PRACTICAL EXERCISES.

1. What percentage of manganese (Mn) is contained in permanganate of

potash (KMnO_{4})?

2. Ten c.c. of a solution of permanganate of potash is found to oxidise

10 c.c. of an acid solution of ferrous sulphate. The manganese is

determined in the titrated solution by precipitation as dioxide and

titrating. How much of the ferrous solution will be oxidised in the

second titration?

3. What weight of potassium iodide would be just sufficient to absorb

the chlorine evolved by 0.5 gram of pure dioxide of manganese?

4. What weight of iron must be dissolved up so as to have an excess of

0.25 gram after oxidation by 1 gram of pure dioxide?

5. What weight of the brown oxide, Mn_{2}O_{4} will be left on igniting

1 gram of the pure dioxide?

CHROMIUM.

Chromium occurs in nature chiefly as chromite or chrome iron ore

(FeO_{2}Cr_{2}O_{3}, with more or less MgO and Al_{2}O_{3}), which is

the chief ore. It is a constituent of some silicates, and is frequently

met with in very small quantities in iron ores. It occurs as chromate in

crocoisite (PbCrO_{4}), and some other rare minerals.

The metal is used in steel-making. Steel containing about 0.5 per cent.

of it is rendered very hard; but its chief value is in its salts, the

chromates. These are highly-coloured compounds, generally red or yellow.

Some of the insoluble chromates are used as pigments; chromate of lead

or chrome-yellow is the most important. The soluble chromates, those of

soda and potash, are valuable chemicals, and are largely used in the

preparation of pigments, dyeing and tanning, and as oxidising agents.

Chromium forms two important classes of compounds--chromic salts,

corresponding to the oxide Cr_{2}O_{3}, and chromates, which contain the

trioxide CrO_{3}. Solutions of chromic salts are green, whilst those of

the chromates are yellow. Chromates are reduced to chromic salts by the

action of most reducing agents in the presence of an acid; and this

property is used in assaying for the volumetric determination of ferrous

iron, &c. The chromates in solution are more stable than other similar

oxidising agents, and consequently are generally used in the laboratory

as one of the standard oxidising agents for volumetric analysis. They

have the disadvantage of requiring an outside indicator. Bichromate of

potash (K_{2}Cr_{2}O_{7}) is the salt generally used for this purpose.

Chromic salts are oxidised to chromate by fusion with "fusion mixture"

and nitre, or by treating with chlorine in an alkaline solution.

Chromic salts closely resemble those of ferric iron, and in the ordinary

course of analysis chromic hydrate (green) is precipitated together with

ferric hydrate, alumina, &c., on the addition of ammonic chloride and

ammonia. The ignited oxide, Cr_{2}O_{3}, however, is not reduced on

heating to redness in a current of hydrogen.

~Detection.~--Chromium is detected by fusing the powdered substance with

"fusion mixture" and nitre. The melt is extracted with water and

filtered. The filtrate is acidified with acetic acid, and treated with a

few drops of a solution of lead acetate. A yellow precipitate indicates

chromium. Substances containing chromium impart a green colour to the

borax bead in both flames. Small quantities of chromate in neutral

solution can be found by the dark or violet-red colouration imparted

thereto on boiling with a dilute decoction of logwood.

~Solution and Separation.~--Chromates and chromic salts are generally

soluble in water or dilute acids. Chrome iron ore, however, and ignited

chromic oxide are insoluble; and the former presents considerable

difficulty on attempting to open up by the usual methods. A large number

of mixtures have been tried in order to get all the chromium in a

soluble form. Among these are the following. One part of the very

finely-powdered ore is fused with any of these mixtures.

(1) 10 parts of bisulphate of potash.

(2) 5 parts of bisulphate of potash and 5 parts of potassium fluoride.

(3) 5 parts of hydric potassic fluoride.

(4) 12 parts of bisulphate of potash; and, afterwards, with 6 parts of

carbonate of soda and 6 parts of nitre.

(5) 8 parts of borax; afterwards, with carbonate of soda till it ceases

to effervesce; then, with 3 parts of carbonate of soda and 3 of

nitre.

(6) 4 parts of borax and 6 parts of fusion mixture.

(7) 12 parts of caustic potash.

(8) 10 parts of caustic soda and 30 of magnesia.

(9) 5 parts of caustic soda and 3 of magnesia.

(10) 2 parts of carbonate of soda and 1 of lime.

(11) 6 parts of soda-lime and 2 of chlorate of potash.

(12) Sodium peroxide.

Of these, numbers 1, 2, and 3 yield the chromium in a form soluble in

dilute acids, as chromic salt. The rest in a form soluble in water, as

potassium or sodium chromate.

On boiling an insoluble chromium compound with chlorate of potash and

nitric acid, the chromium passes into solution as chromate. This method,

however, does not answer for chrome iron ore. In the fusion methods the

ore must be very finely powdered, well mixed with the fluxes, and

subjected to a prolonged fusion in a platinum vessel at a high

temperature. Undecomposed particles require re-fusion.

The aqueous extract containing the chromate is ready for volumetric

work, except in those cases where nitre has been used. For gravimetric

work the solution is acidified with hydrochloric acid, then mixed with

ammonia in slight excess, boiled, and filtered. The filtrate is

acidified with hydrochloric acid, and treated with sulphuretted

hydrogen, warmed, rendered slightly alkaline with ammonia, and the gas

again passed. The chromium is precipitated as chromic hydrate mixed with

sulphur from the reduction with sulphuretted hydrogen. It is filtered

off, washed with hot water, and ignited. It is weighed as chromic oxide.

GRAVIMETRIC DETERMINATION.

The solution containing the chromium, freed from other metals and earths

and in the form of (green) chromic salt, is heated to boiling. If any

chromate is present reduce it with sodium sulphite or sulphuretted

hydrogen. Add ammonia in slight excess, boil till the liquid is free

from a red tint, and allow to settle for a few minutes. Filter, wash

with hot water, dry, and ignite strongly in a loosely-covered crucible.

Cool, and weigh. The substance is chromic oxide, Cr_{2}O_{3}, and

contains 68.62 per cent. of chromium. It is a dark-green powder

insoluble in acids.

When, as is generally the case, the chromium exists altogether as

chromate (phosphates and arsenates being absent) it is best to proceed

as follows:--Render the solution acid with acetic acid, then add sodium

acetate to the solution and heat nearly to boiling; next treat with a

slight excess of acetate of lead, and boil. Allow to settle, and filter.

Wash the precipitate with hot water, dry in the water-oven or at a low

temperature. Transfer the precipitate to a weighed Berlin crucible, burn

the filter separately, ignite below redness, cool in the desiccator, and

weigh. The substance is lead chromate, PbCrO_{4}, and contains 16.1 per

cent. of chromium, or 23.53 per cent. of chromic oxide (Cr_{2}O_{3}).

VOLUMETRIC METHOD.

This is based on the oxidation of ferrous iron by the solution

containing the chromium as chromate. A known weight of iron (0.5, 1, or

1.5 gram, according to the quantity of chromate) is dissolved in 50 c.c.

of dilute sulphuric acid. The solution containing the chromate is added,

and the remaining ferrous iron titrated with the permanganate or

bichromate of potassium solution, as described under _Iron_. The iron

thus found is deducted from that taken, and the difference gives the

iron oxidised by the chromate. This multiplied by 0.3101 gives the

chromium, Cr, and when multiplied by 0.4529 gives the chromic oxide,

Cr_{2}O_{3}.

COLORIMETRIC METHOD.

Small quantities of chromium may be determined, after conversion into

chromate, colorimetrically. The solution, which should not contain more

than a few milligrams in 100 c.c., is acidified with acetic acid and

compared against an equal volume of water rendered acid with acetic

acid and tinted with a standard bichromate of potassium solution. This

standard bichromate is made by dissolving 2.827 grams of the salt in

water and diluting to 1 litre. One c.c. will contain 1 milligram of

chromium, Cr. The manner of working this assay is the same as that

adopted in the other colorimetric processes.

~Determination of Chromium in Steel.~[89]--Weigh up 2.4 grams, dissolve

in hydrochloric acid, and evaporate to dryness. Fuse with sodium

carbonate and nitre, extract with water, and make up to 301 c.c. Take

250 c.c. of the clear liquor, boil with hydrochloric acid, add sodium

phosphate, and then ammonia in slight excess. Heat till clear. Filter

off the precipitate, dissolve it in hydrochloric acid, and evaporate to

dryness. Take up with a little acid, filter, and precipitate with a

slight excess of ammonia. Wash, ignite, and weigh as chromium phosphate

(3Cr_{2}O_{3},2P_{2}O_{5}), which contains 42.2 per cent. of chromium.

VANADIUM.

Vanadium occurs in certain rare minerals, such as vanadinite

(3Pb_{3}(VO_{4})_{2}.PbCl_{2}), a vanadate of lead; mottramite, a

vanadate of copper and lead; and dechenite, a vanadate of lead and zinc.

It is occasionally found in iron and copper ores and in the slags from

them. In Spanish copper-precipitates it is found along with chromium,

and is probably derived from the iron used for precipitating. The

vanadates, like the chromates, are coloured compounds, generally yellow

or red. On reduction, blue solutions are got. In their general reactions

vanadates resemble phosphates.

Vanadium is detected by the red colouration produced by passing

sulphuretted hydrogen into ammoniacal solutions for some time. On adding

an acid to the filtered solution a brown precipitate of the sulphide is

produced. This gives with borax a colourless bead in the oxidising, and

a green one in the reducing, flame.

It is separated by fusing the ore with potassic nitrate, extracting with

water and precipitating with baric chloride. The precipitate is boiled

with dilute sulphuric acid, filtered, neutralised with ammonia, and

saturated with ammonic chloride. Ammonium vanadate separates out. It is

filtered off, ignited, and weighed as vanadic oxide, V_{2}O_{5},

containing 56.18 per cent. of vanadium.

MOLYBDENUM.

Molybdenum occurs in nature chiefly as molybdenite (MoS_{2}); it also

occurs in wulfenite, a molybdate of lead (PbMoO_{4}), and in molybdic

ochre (MoO_{3}).

Molybdate of ammonia is an important reagent in the determination of

phosphates, the manufacture of which compound is the chief purpose to

which molybdenum is applied.

Iron and copper ores frequently contain molybdenum, sometimes in

quantity; consequently it is met with in slags and pig-iron.

Molybdenum forms several series of salts. In those corresponding to the

lower oxides it is basic; but the trioxide (MoO_{3}) is the acid oxide

which forms a series of salts called the molybdates. All molybdenum

compounds are converted into the trioxide by boiling with nitric acid.

The trioxide is a white powder readily dissolved by ammonia. It fuses at

a red heat, and volatilises freely in contact with air. It is slightly

soluble in water.

Molybdates are easily reduced, with the production of coloured

solutions, by most reducing agents. Sulphuretted hydrogen first produces

a blue tint, and then precipitates a brown sulphide. The precipitation

as sulphide is only complete on prolonged treatment; a green colour

indicates that some molybdenum still remains in solution. The

precipitated sulphide is soluble in ammonium sulphide.

~Detection.~--Molybdenum is detected by its behaviour with sulphuretted

hydrogen. Molybdenite can only be mistaken for graphite, from which it

is easily distinguished by yielding sulphur dioxide on roasting, and by

giving, on charcoal, a yellowish white incrustation, which becomes blue

on touching it for a moment with the reducing flame. The borax-bead is

colourless in the oxidising, and dark-brown in the reducing, flame.

GRAVIMETRIC DETERMINATION.

The solution containing the molybdate is neutralised and treated with an

excess of mercurous nitrate. The precipitate is allowed to settle for

some time, filtered, and washed with a dilute solution of mercurous

nitrate. Then it is dried, transferred to a weighed Berlin crucible

containing ignited oxide of lead, mixed, ignited, and weighed. The

increase in weight gives the amount of trioxide, MoO_{3}. This contains

66.7 per cent. of molybdenum.

URANIUM.

Uranium occurs chiefly as pitchblende, which is an impure oxide

(U_{3}O_{8}). It is also found as sulphate in uranochre, johannite, &c.;

and as phosphate in the uranites, torbernite (hydrated phosphate of

uranium and copper), and autunite (hydrated phosphate of uranium and

lime). It also occurs in some rarer minerals.

The oxide is used for colouring glass; and the nitrate and acetate are

used as reagents. "Uranium yellow," used for enamel painting, is sodium

uranate. The uranates, in which the oxide of uranium acts as an acid,

are mostly insoluble and of secondary importance.

Uranium forms two families of salts, uranous and uranic; corresponding

to the oxides UO_{2} and UO_{3} respectively. The former are generally

green and the latter yellow. Uranous salts are converted into uranic by

boiling with nitric acid or other oxidising agents. Uranic salts, on the

other hand, are easily reduced by sulphuretted hydrogen, stannous

chloride or zinc. This property is made use of in determining the

quantity of uranium in pure solutions by titrating with permanganate of

potassium solution as in the case with iron.

~Detection.~--The most characteristic reaction of the uranium compounds

is their behaviour in the presence of alkaline carbonates in which they

are freely soluble; even ammonium sulphide will not precipitate uranium

from these solutions. On neutralising the carbonate with an acid a

uranate of the alkali is precipitated. Ammonia or sodic hydrate (free

from carbonates) give yellow precipitates, which are insoluble in excess

of the reagent, but are soluble in acids. Ferrocyanide of potassium

gives a reddish-brown precipitate. Uranium colours the borax-bead

yellowish-green in the oxidising, and green in the reducing, flame.

~Solution and Separation.~--The compounds of uranium are soluble in

acids. Powder the substance and evaporate with an excess of nitric acid.

Take up with hydrochloric acid, dilute, pass sulphuretted hydrogen, and

filter. Peroxidise the filtrate with a little nitric acid, add an excess

of ammonic carbonate and some ammonium sulphide, and filter. Render the

solution acid, boil; and precipitate the uranium by means of ammonia.

Filter off, and wash it with dilute ammonic chloride. Ignite, and weigh

as protosesqui-oxide, U_{3}O_{8}.

GRAVIMETRIC DETERMINATION.

The solution containing the uranium free from other metals is, if

required, first peroxidised by boiling with nitric acid. Ammonia in

slight excess is added to the nearly-boiling solution. A yellow

precipitate is formed, which is filtered off hot and washed with a

dilute solution of ammonium chloride. The precipitate is dried and

ignited; and weighed as U_{3}O_{8}, which contains 84.8 per cent. of

uranium.

VOLUMETRIC METHOD.

This is based on the precipitation of uranium as phosphate from acetic

acid solutions and the recognition of complete precipitation by testing

with potassic ferrocyanide; it is the converse of the process for the

volumetric determination of phosphate.

_The standard solution of phosphate_ is prepared by dissolving 29.835

grams of hydric sodic phosphate (Na_{2}HPO_{4}.12H_{2}O) in water and

diluting to 1 litre. 100 c.c. will be equivalent to 2 grams of uranium.

Take 1 gram of the sample (or, if poor in uranium, 2 grams) and separate

the uranium as described. Dissolve the precipitate in nitric acid and

evaporate to a small bulk, add 2 grams of sodium acetate, dilute with

water to 100 c.c., and boil. Titrate the boiling solution with the

sodium phosphate till it ceases to give a brown colouration with

potassium ferrocyanide. Calculate the percentage in the usual way.

FOOTNOTES:

[78] MnO_{2} + 4HCl = MnCl_{2} + Cl_{2} + 2H_{2}O.

[79] Provided a sufficiency of ammonic chloride is present.

[80] With some silicates, &c., a preliminary fusion with sodium

carbonate will be necessary.

[81] Instead of sodium acetate, ammonium succinate can be used.

[82] _Journ. Soc. Chem. Industry_, vol. x. p. 333.

[83] MnO_{2} + 2FeSO_{4} + 2H_{2}SO_{4} = Fe_{2}(SO_{4})_{3} + MnSO_{4}

+ 2H_{2}O.

[84] If the ore is very rich, a smaller quantity (0.75 or 1.5 gram) must

be taken; otherwise the iron will be insufficient.

[85] MnO_{2} + 4HCl = MnCl_{2} + 2H_{2}O + Cl_{2}.

Cl_{2} + 2KI = 2KCl + I_{2}.

[86] Iodine probably lost by volatilisation.

[87] Obtained as a brown powder by digesting red lead with nitric acid

and filtering.

[88] The water for dilution and the dilute sulphuric acid used for

washing should be previously tested, to see they have no reducing

action, with dilute permanganate of potassium solution.

[89] Arnold and Hardy, _Chemical News_, vol. lvii. p. 153.

CHAPTER XIV.

EARTHS, ALKALINE EARTHS, ALKALIES.

ALUMINA.

Alumina, the oxide of aluminium (Al_{2}O_{3}), is found in nature fairly

pure in the mineral corundum; transparent and coloured varieties of

which form the gems sapphire and ruby. A coarser compact variety

contaminated with oxide of iron constitutes emery. Compounded with

silica, alumina forms the base of clays and many rock-forming minerals.

China clay (or kaolin) is used as a source of alumina. Bauxite, hydrated

alumina, is also used for the same purpose--that is, for the preparation

of sulphate of alumina. The mineral cryolite is a fluoride of aluminium

and sodium.

Corundum is characterised by a high specific gravity (4.0) and extreme

hardness. By these it is distinguished from felspar and similar

minerals, which it somewhat resembles in general appearance.

Aluminium is used for a variety of small purposes: it is white, light,

and very tenacious; but owing to the difficulty of its reduction it is

expensive.

Aluminium forms one series of salts which closely resemble those of

ferric iron. It forms an interesting series of double sulphates, known

as the alums. Common potash alum is

Al_{2}(SO_{4})_{3},K_{2}SO_{4},24H_{2}O.

~Detection.~--Alumina is not precipitated from its acid solution by

sulphuretted hydrogen, but it is thrown down by ammonia (with the other

earths) as a white hydrate, soluble in soda and insoluble in ammonic

carbonate. Filtered off and ignited, it assumes, after treatment with

nitrate of cobalt before the blowpipe, a blue colour which is

characteristic. With natural compounds containing metallic oxides this

colour is masked. It is more satisfactory to make a separation in the

wet way and to test the ignited oxide.

~Separation and Solution.~--If the substance is insoluble in

hydrochloric acid it is finely powdered and fused with "fusion mixture"

with the help, in the case of corundum (which is very refractory) of a

little caustic soda or potash. The method of working is the same as

that described for opening up silicates. See under _Silica_. Corundum

cannot be powdered in Wedgwood, or even agate, mortars; since it rapidly

wears these away and becomes contaminated with their powder. It is best

to use a hard steel mortar and to extract the metallic particles from

the bruised sample with a magnet or dilute acid.

When the substance has been completely attacked and dissolved, it is

evaporated to dryness with an excess of hydrochloric acid on the

water-bath to render any silica present insoluble. The residue is

extracted with hydrochloric acid and freed from the second group of

metals by means of sulphuretted hydrogen. The filtrate from this (after

removing the sulphuretted hydrogen by boiling) is nearly neutralised,

and treated with 8 or 10 grams of hyposulphite of soda[90] in solution.

It is then boiled till the sulphurous oxide is driven off. The

precipitate is filtered off, ignited, and weighed as alumina.

It is sometimes more convenient to proceed as follows:--After boiling

off the sulphuretted hydrogen peroxidise the iron with a little nitric

acid, add a solution of ammonic chloride, and then ammonia in very

slight excess; boil, filter, wash, ignite, and weigh the oxides. These

generally consist of ferric oxide and alumina. It is a common practice

to determine the iron, calculate it to ferric oxide, and so to estimate

the alumina indirectly. This may be done either by igniting in a current

of hydrogen and estimating the iron by the weight of oxygen lost; or, by

dissolving with sulphuric and hydrochloric acids, and determining the

iron volumetrically. It should be borne in mind that these oxides will

also contain any phosphoric oxide that happened to be in the mineral.

In general analyses of samples containing alumina, it may be contained

in both the soluble and insoluble portions. In these cases it is better

to fuse the sample with "fusion mixture" before treatment with acids.

The alumina in the fused mass will exist in a state soluble in acids.

GRAVIMETRIC DETERMINATION.

Solutions containing alumina free from the other metals are diluted to a

convenient bulk and heated nearly to boiling. Add chloride of ammonium,

and then ammonia in slight excess; boil, allow to settle, filter, and

wash with hot water. Dry the precipitate, and ignite in a platinum or

porcelain crucible at the strongest heat. Cool, and weigh. The substance

is alumina, Al_{2}O_{3}, which contains 52.94 per cent. of aluminium.

It is only in special cases, such as the analysis of metals and alloys,

that it is reported as aluminium. The percentage of alumina is generally

given.

Ignited alumina is difficultly soluble in acids; it is not reduced by

hydrogen at a red heat. Ignited with ammonium chloride portions are

volatilised.

~Direct Determination of Alumina in the Presence of Iron.~--The iron and

alumina are precipitated as hydrates by ammonia. The precipitate is

dissolved in hydrochloric acid and the iron reduced to the ferrous

state. It is then added to a hot solution of potash or soda. The

solution is boiled till the precipitate settles readily, filtered, and

washed with hot water. The alumina is contained in the filtrate, which

is acidified with hydrochloric acid and the alumina precipitated

therefrom as hydrate with ammonia, as just described.

~Determination of Alumina in the Presence of Phosphates and Iron.~--For

details, see a paper by R.T. Thomson in the "Journal of the Society of

Chemical Industry," v. p. 152. The principles of the method are as

follows:--If the substance does not already contain sufficient

phosphoric oxide to saturate the alumina, some phosphate is added. The

iron is reduced to the ferrous state and phosphate of alumina

precipitated in an acetic acid solution. It is purified by

reprecipitation, ignited, and weighed as phosphate

(Al_{2}O_{3},P_{2}O_{5}), which contains 41.8 per cent. of alumina,

Al_{2}O_{3}.

EXAMINATION OF CLAYS.

~Moisture.~--Take 5 grams of the carefully-prepared sample and dry in

the water-oven till the weight is constant.

~Loss on Ignition.~--Weigh up 2 grams of the sample used for the

moisture determination, and ignite in a platinum-crucible to redness,

cool, and weigh.

~Silica and Insoluble Silicates.~--Weigh up another 2 grams of the dried

sample, and place them in a platinum dish; moisten with water, and cover

with 20 c.c. of sulphuric acid. Evaporate and heat gently to drive off

the greater portion of the free acid. Allow to cool; and repeat the

operation. Extract by boiling with dilute hydrochloric acid, filter,

wash, dry, ignite, and weigh. The quantity of insoluble silicates is

determined by dissolving out the separated silica with a strong boiling

solution of sodium carbonate. The residue (washed, dried, and ignited)

is weighed, and reported as "sand."

~Alumina and Ferrous Oxide.~--To the filtrate from the silica add "soda"

solution till nearly neutral, and then sodium acetate. Boil and filter

off the precipitate. Reserve the filtrate. Dissolve the precipitate in

hydrochloric acid, and dilute to exactly 200 c.c. Divide into two parts

of 100 c.c. each. In one determine the iron by reducing and titrating in

the way described under volumetric iron. Calculate the percentage as

ferrous oxide, unless there are reasons to the contrary, also calculate

its weight as ferric oxide. To the other portion add ammonia in slight

excess, and boil. Filter, wash with hot water, dry, ignite, and weigh as

mixed alumina and ferric oxide. The weight of the ferric oxide has

already been determined in the first portion: deduct it, and the

difference is the weight of alumina.

~Lime.~--To the reserved filtrate, concentrated by evaporation, add

ammonium oxalate and ammonia; boil, filter, ignite strongly, and weigh

as lime.

~Magnesia~ is separated from the filtrate by adding sodium phosphate. It

is weighed as magnesium pyrophosphate.

~Potash and Soda.~--These are determined in a fresh portion of the

sample by Lawrence Smith's method, as described on page 333.

THORIA.

This is an oxide of thorium, ThO_{2}. It is only found in a few rare

minerals. It is a heavy oxide, having, when strongly ignited, a specific

gravity of 9.2. In the ordinary course of analysis it will be separated

and weighed as alumina. It is separated from this and other earths by

the following method. The solution in hydrochloric acid is nearly

neutralised and then boiled with sodium hyposulphite. The thoria will be

in the precipitate. It is dissolved, and the solution heated with

ammonium oxalate in excess. The precipitate is thorium oxalate, which is

washed with hot water, dried, and ignited. It is then weighed as thoria,

ThO_{2}. Thoria which has been ignited is not readily soluble in acids.

ZIRCONIA.

The oxide of zirconium, ZrO_{2}, is found in the mineral zircon, a

silicate of zirconia, ZrSiO_{4}. When heated intensely it becomes very

luminous, and is used on this account for incandescent lights.

In the ordinary course it is thrown down by ammonia with the other

earths, from which it is thus separated:--The hydrates precipitated in

the cold, and washed with cold water, are dissolved in hydrochloric

acid, nearly neutralised with soda, and precipitated by boiling with

hyposulphite of soda. Dissolve; and from the hydrochloric acid solution

precipitate the thoria (if any) with ammonium oxalate. To the filtrate

add carbonate of ammonia, which will precipitate any titanium present.

The zirconia will be in solution, and is recovered by precipitating with

potassium sulphate, or by evaporating the solution and igniting. It is

separated from alumina by taking advantage of its insolubility in

potassic hydrate.

It is estimated in zircons in the following way:--The powdered substance

is fused with bisulphate of potash, and extracted with dilute sulphuric

acid. The residue is fused with caustic soda and extracted with water.

The portion not dissolved, consisting of zirconate of soda, is dissolved

in hydrochloric acid. The solution is diluted, filtered if necessary,

and treated with ammonia in excess. The precipitate is filtered off,

washed with hot water, dried, ignited, and weighed as zirconia, ZrO_{2}.

This is a white powder, which is insoluble in acids; even in

hydrofluoric acid it is only slightly attacked.

CERIUM.

Cerium occurs as silicate (together with the oxides of lanthanum,

didymium, iron and calcium) in the mineral cerite, which is its chief

source. It also occurs as phosphate in monazite, and as fluoride in

fluocerite. The oxalate is used in medicine. Cerium forms two classes of

salts corresponding to the oxides, cerous oxide (Ce_{2}O_{3}) and ceric

oxide (CeO_{2}). Compounds of cerium with volatile acids yield dioxide

on ignition; and this, on solution in hydrochloric acid, yields cerous

chloride and chlorine.

In the ordinary course cerium is thrown down along with alumina and the

other earths by ammonia. It is separated by dissolving the hydrates in

hydrochloric acid, and oxidizing with chlorine water. On treating with

oxalic acid, cerium, lanthanum, and didymium are precipitated as

oxalates, which on ignition are converted into oxides. These are soluble

in acids. Their solution in hydrochloric acid is nearly neutralised;

acetate of soda is then added, and an excess of sodium hypochlorite. On

boiling, the cerium is precipitated as dioxide, which is filtered off,

ignited, and weighed.

Cerium is detected by giving with borax a bead which is yellow in the

oxidising, and colourless in the reducing flame. Traces of cerium

compounds boiled with dioxide of lead and nitric acid will give a yellow

solution.

LANTHANUM AND DIDYMIUM

occur together with cerium in cerite, and are separated with that metal

as oxalates, as described under _Cerium_.

Didymium salts have a rose or violet colour, and impart (when in

sufficient quantity) the same colour to the borax bead. Solutions have a

characteristic absorption-spectrum.

The separation of lanthanum and didymium in the solution from which the

cerium has been precipitated is effected by precipitating them together

as oxalates, igniting, and dissolving in dilute nitric acid. This

solution is then evaporated to dryness and ignited, for a few minutes,

just below redness. A subnitrate of didymium is formed, and remains as

an insoluble residue on extracting with hot water. The separated salts

are treated with ammonia and ignited, and weighed as oxides (La_{2}O_{3}

and Di_{2}O_{3}).

YTTRIA.

Yttria is found in gadolinite and some other rare minerals. It is

precipitated along with the other earths by ammonia. It is distinguished

by the insolubility of its hydrate in potash, by the insolubility of its

oxalate in oxalic acid, and by not being precipitated by hyposulphite of

soda or potassium sulphate. Further, it is precipitated by potash in the

presence of tartaric acid as an insoluble tartrate. This reaction

distinguishes the members of the yttria group from most of the other

earths. The other members of the group closely resemble it, and amongst

them are erbia, terbia, ytterbia, scandia, &c.

BERYLLIA.

The oxide of beryllium, BeO (also known as glucina), occurs in nature

mainly as silicate. Beryl, the green transparent variety of which is the

emerald, is the best known of these. It is a silicate of alumina and

beryllia.[91] Some other minerals in which it occurs are phenakite,

euclase, and chrysoberyl.

In the ordinary course of analysis, beryllia will be precipitated with

alumina, &c., by ammonic hydrate. It is distinguished by the solubility

of its hydrate in ammonic carbonate, by not being precipitated by

boiling with sodium hyposulphite, and by not being precipitated by

ammonic sulphide from an ammonic carbonate solution.

The analysis of silicates containing beryllia is thus effected. The

finely powdered substance is fused with twice its weight of potassium

carbonate; and the "melt" is extracted with water, and evaporated with a

slight excess of sulphuric acid to render the silica insoluble. Treat

with water, filter, and evaporate the filtrate until a crust is formed.

Potash alum crystallises out. The liquor is poured off into a warm

strong solution of ammonium carbonate. Ferric hydrate and alumina will

be precipitated. They are filtered off, re-dissolved, and again

precipitated in ammonic carbonate solution; the combined filtrates are

boiled for some time, and acidified slightly with hydrochloric acid. The

carbon dioxide is boiled off, and the beryllia is then precipitated as

hydrate with ammonia. The hydrate is washed with hot water, dried,

ignited, and weighed as beryllia, BeO.

Beryllia has a specific gravity of 3.08. It is white, infusible, and

insoluble in water. After ignition, it is insoluble in acids, except

sulphuric, but is rendered soluble by fusion with alkalies.

Beryllia, in a solution of carbonate of ammonia, is precipitated as

carbonate on boiling in proportion as the carbonate of ammonia is

volatilised. The hydrate is dissolved by a boiling solution of ammonic

chloride, ammonia being evolved.

THE ALKALINE EARTHS.

LIME.

Lime is an oxide of calcium, CaO. It occurs abundantly in nature, but

only in a state of combination. The carbonate (CaCO_{3}), found as

limestone, chalk, and other rocks, and as the minerals calcite and

arragonite, is the most commonly occurring compound. The hydrated

sulphate, gypsum (CaSO_{4}.2H_{2}O), is common, and is used in making

"plaster of Paris." Anhydrite (CaSO_{4}) also occurs in rock masses, and

is often associated with rock salt. Phosphate of lime, in the forms of

apatite, phosphorite, coprolite, &c., is largely mined. Lime is a

component of most natural silicates. Calcium also occurs, combined with

fluorine, in the mineral fluor (CaF_{2}). In most of these the acid is

the important part of the mineral; it is only the carbonate which is

used as a source of lime.

Lime, in addition to its use in mortars and cements, is valuable as a

flux in metallurgical operations, and as a base in chemical work on a

large scale. A mixture of lime and magnesia is used in the manufacture

of basic fire-bricks.

Carbonate of lime on ignition, especially when in contact with reducing

substances, loses carbonic acid, and becomes lime. This is known as

"quicklime"; on treatment with water it becomes hot, expands, and falls

to a powder of "slaked lime" or calcium hydrate (CaH_{2}O_{2}). The

hydrate is slightly soluble in water (0.1368 gram in 100 c.c.), forming

an alkaline solution known as lime-water. Calcium hydrate is more

generally used suspended in water as "milk of lime."

As a flux it is used either as limestone or as quicklime. Silica forms

with lime a compound, calcium silicate, which is not very fusible; but

when alumina and other oxides are present, as in clays and in most rocky

substances, the addition of lime gives a very fusible slag.

~Detection.~--Calcium is detected by the reddish colour which its salts

impart to the flame. It is best to moisten with hydrochloric acid (or,

in the case of some silicates, to treat with ammonium fluoride) before

bringing the substance into the flame. When seen through a spectroscope,

it shows a large number of lines, of which a green and an orange are

most intense and characteristic. Calcium is detected in solution (after

removal of the metals by treatment with sulphuretted hydrogen and

ammonium sulphide) by boiling with ammonium oxalate and ammonia. The

lime is completely thrown down as a white precipitate. Lime is

distinguished from the other alkaline earths by forming a sulphate

insoluble in dilute alcohol, but completely soluble in a boiling

solution of ammonium sulphate.

Lime compounds are for the most part soluble in water or in dilute

hydrochloric acid. Calcium fluoride must be first converted into

sulphate by evaporation in a platinum dish with sulphuric acid.

Insoluble silicates are opened up by fusion with "fusion mixture," as

described under _Silica_.

~Separation.~--The separation of lime is effected by evaporating with

hydrochloric acid, to separate silica; and by treating with sulphuretted

hydrogen, to remove the second group of metals. If the substance

contains much iron, the solution is next oxidised by boiling with a

little nitric acid; and the iron, alumina, &c., are removed as basic

acetates. The filtrate is treated with ammonia and sulphuretted

hydrogen, and allowed to settle. The filtrate from this is heated to

boiling, treated with a solution of ammonium oxalate in excess, boiled

for five or ten minutes, allowed to settle for half an hour, and

filtered. The precipitate contains all the lime as calcium oxalate.

GRAVIMETRIC DETERMINATION.

The precipitate of calcium oxalate is washed with hot water, dried,

transferred to a weighed platinum crucible, and ignited at a

temperature not above incipient redness. This ignition converts the

oxalate into carbonate, with evolution of carbonic oxide, which burns at

the mouth of the crucible with a blue flame.[92] Generally a small

quantity of the carbonate is at the same time converted into lime. To

reconvert it into carbonate, moisten with a few drops of ammonic

carbonate solution, and dry in a water-oven. Heat gently over a Bunsen

burner, cool, and weigh. The substance is calcium carbonate (CaCO_{3}),

and contains 56 per cent. of lime (CaO). It is a white powder, and

should show no alkaline reaction with moistened litmus-paper.

Where the precipitate is small, it is better to ignite strongly over the

blowpipe, and weigh directly as lime. With larger quantities, and when

many determinations have to be made, it is easier to make the

determination volumetrically.

VOLUMETRIC METHODS.

These are carried out either by dissolving the oxalate at once in dilute

sulphuric acid, and titrating with permanganate of potassium solution;

or by calcining it to a mixture of lime and carbonate, and determining

its neutralising power with the standard solutions of acid and alkali.

~Titration with Permanganate of Potassium Solution.~--This solution is

made by dissolving 5.643 grams of the salt in water, and by diluting to

1 litre; 100 c.c. are equivalent to 0.5 gram of lime. The solution is

standardised by titrating a quantity of oxalic acid about equivalent to

the lime present in the assay; 0.5 gram of lime is equivalent to 1.125

gram of crystallised oxalic acid. The standardising may be done with

iron. The standard found for iron multiplied by 0.5 gives that for lime.

The process is as follows:--The calcium oxalate (having been

precipitated and washed, as in the gravimetric process) is washed

through the funnel into a flask with hot dilute sulphuric acid, boiled

till dissolved, diluted to 200 c.c. with water, and heated to about 80° C.

The standard solution of "permanganate" is then run in, (not too

quickly, and with constant shaking) until a permanent pink tinge is

produced. The c.c. used multiplied by the standard, and divided by the

weight of the substance taken, will give the percentage of lime.

~Estimation of Lime by Alkalimetry.~--The methods of determining the

amount of an alkali or base by means of a standard acid solution, or,

conversely, of determining an acid by means of a standard alkaline

solution, are so closely related that they are best considered under

one head. The same standard solution is applicable for many purposes,

and, consequently, it is convenient to make it of such strength that one

litre of it shall equal an equivalent in grams of any of the substances

to be determined. Such solutions are termed _normal_. For example, a

solution of hydrochloric acid (HCl = 36.5) containing 36.5 grams of real

acid per litre, would be normal and of equivalent strength to a solution

containing either 17 grams of ammonia (NH_{3} = 17) or 40 grams of sodic

hydrate (NaHO = 40) per litre. It will be seen in these cases that the

normal solution contains the molecular weight in grams per litre; and,

if solutions of these strengths be made, it will be found that they

possess equal neutralising value.

If, now, a solution containing 98 grams of sulphuric acid (H_{2}SO_{4} =

98) per litre be made, it will be found to have twice the strength of

the above solution, that is, 100 c.c. of the soda would only require 50

c.c. of the acid to neutralise it. The reason for this will be seen on

inspecting the equations:--

NaHO + HCl = NaCl + H_{2}O.

2NaHO + H_{2}SO_{4} = Na_{2}SO_{4} + 2H_{2}O.

Acids like sulphuric acid are termed bibasic, and their equivalent is

only half the molecular weight. Thus, a normal solution of sulphuric

acid would contain 49 grams (98/2) of real acid per litre. Similarly,

lime and most of the bases are bibasic, as may be seen from the

following equations; hence their equivalent will be half the molecular

weight.

2HCl + CaO = CaCl_{2} + H_{2}O.

2HCl + MgO = MgCl_{2} + H_{2}O.

_The standard normal solution of hydrochloric acid_ is made by diluting

100 c.c. of the strong acid to one litre with water. This will be

approximately normal. In order to determine its exact strength, weigh up

3 grams of recently ignited pure sodium carbonate or of the ignited

bicarbonate. Transfer to a flask and dissolve in 200 c.c. of water; when

dissolved, cool, tint faintly yellow with a few drops of a solution of

methyl orange, and run in the standard "acid " from a burette till the

yellow changes to a pink. Read off the number of c.c. used, and

calculate to how much sodium carbonate 100 c.c. of the "acid" are

equivalent. If the "acid" is strictly normal, this will be 5.3 grams. It

will probably be equivalent to more than this. Now calculate how much

strictly normal "acid" would be equivalent to the standard found. For

example: suppose the standard found is 5.5 gram of sodium carbonate,

then--

5.3 : 5.5 :: 100 : _x_

(where _x_ is the quantity of normal "acid" required).

_x_ = 103.8 c.c.

To get the "acid" of normal strength, we should then add 3.8 c.c. of

water to each 100 c.c. of the standard solution remaining. Suppose there

were left 930 c.c. of the approximate "acid," 35.3 c.c. of water must be

added and mixed. It should then be checked by another titration with

pure sodium carbonate.

_The standard solution of semi-normal "alkali."_ The best alkali for

general purposes is ammonia, but, since it is volatile (especially in

strong solutions), it is best to make it of half the usual strength, or

_semi-normal_. One litre of this will contain 8.5 grams of ammonia

(NH_{3}), and 100 c.c. of it will just neutralise 50 c.c. of the normal

"acid." Take 100 c.c. of dilute ammonia and dilute with water to one

litre. Run into a flask 50 c.c. of the standard "acid," tint with methyl

orange, and run in from a burette the solution of ammonia till

neutralised. Less than 100 c.c. will probably be used. Suppose 95 c.c.

were required, there should have been 100, hence there is a deficiency

of five. Then, for each 95 c.c. of standard "ammonia" left, add 5 c.c.

of water, and mix well. 100 c.c. will now be equivalent to 50 c.c. of

the "acid."

As an example of the application of this method, we may take the

determination of lime in limestone, marble, and similar substances.

~Determination of Lime in Limestone.~--Weigh up 1 gram of the dried

sample, and dissolve in 25 c.c. of normal acid, cool, dilute to 100

c.c., and titrate with the semi-normal solution of alkali (using

methyl-orange as an indicator). Divide the c.c. of alkali used by 2,

subtract from 25, and multiply by 0.028 to find the weight of lime. This

method is not applicable in the presence of other carbonates or oxides,

unless the weight of these substances be afterwards determined and due

correction be made.

STRONTIA.

Strontia, the oxide of strontium (SrO), occurs in nature as sulphate, in

the mineral celestine (SrSO_{4}), and as carbonate in strontianite

(SrCO_{3}). It is found in small quantities in limestones, chalk, &c.

Strontia is used in sugar-refining, and for the preparation of coloured

lights.

~Detection.~--It is detected by the crimson colour which its compounds

(when moistened with hydrochloric acid) impart to the flame. The

spectrum shows a large number of lines, of which a red, an orange, and a

blue are most characteristic.

It resembles lime in many of its compounds, but is distinguished by the

insolubility of its sulphate in a boiling solution of ammonium sulphate,

and by the insolubility of its nitrate in alcohol. From baryta, which it

also resembles, it is distinguished by not yielding an insoluble

chromate in an acetic acid solution, by the solubility of its chloride

in alcohol, and by the fact that its sulphate is converted into

carbonate on boiling with a solution formed of 3 parts of potassium

carbonate and 1 of potassium sulphate.

It is got into solution in the same manner as lime. The sulphate should

be fused with "fusion mixture," extracted with water, and thoroughly

washed. The residue will contain the strontia as carbonate, which is

readily soluble in dilute hydrochloric or nitric acid.

~Separation.~--It is separated (after removal of the silica and metals,

as described under _Lime_) by adding ammonia and ammonia carbonate, and

allowing to stand for some hours in a warm place. In the absence of

baryta or lime it is filtered off, and weighed as strontium carbonate,

which contains 70.17 per cent. of strontia. It is separated from baryta

by dissolving in a little hydrochloric acid, adding ammonia in excess,

and then acidifying with acetic acid, and precipitating the baryta with

potassium bichromate, as described under _Baryta_. The strontia is

precipitated from the filtrate by boiling for some time with a strong

solution of ammonic sulphate and a little ammonia. Fifty parts of

ammonic sulphate are required for each part of strontia or lime present.

The precipitate is filtered off, and washed first with a solution of

ammonic sulphate, and then with alcohol. It is dried, ignited and

weighed as strontium sulphate.

GRAVIMETRIC DETERMINATION.

The determination of strontia in pure solutions is best made by adding

sulphuric acid in excess and alcohol in volume equal to that of the

solution. Allow to stand overnight, filter, wash with dilute alcohol,

dry, ignite at a red heat, and weigh as sulphate (SrSO_{4}). This

contains 56.4 per cent. of strontia (SrO); or 47.7 per cent. of

strontium.

BARYTA.

Baryta, oxide of barium (BaO), commonly occurs in combination with

sulphuric oxide in the mineral barytes or heavy spar (BaSO_{4}), and in

combination with carbon dioxide in witherite (BaCO_{3}). These minerals

are not unfrequently found in large quantity (associated with galena and

other metallic sulphides) in lodes. Small isolated crystals of these are

frequently found in mining districts. Barium is a constituent of certain

mineral waters. The minerals are recognised by their high specific

gravity and their crystalline form.

Compounds of barium are often used by the assayer, more especially the

chloride and hydrate. The salts are, with the exception of the sulphate,

generally soluble in water or hydrochloric acid. In such solutions

sulphuric acid produces a white precipitate of baric sulphate, which is

practically insoluble in all acids.

The dioxide (BaO_{2}) is used for the preparation of oxygen. On strong

ignition it gives up oxygen, and is converted into baryta (BaO), which,

at a lower temperature, takes up oxygen from the air, re-forming the

dioxide.

~Detection.~--Barium is detected by the green colour its salts,

especially the chloride, give to the flame. This, viewed through the

spectroscope, shows a complicated spectrum, of which two lines in the

green are most easily recognised and characteristic. The salts of barium

give no precipitate with sulphuretted hydrogen in either acid or

alkaline solution, but with sulphuric acid they at once give a

precipitate, which is insoluble in acetate of soda. In solutions

rendered faintly acid with acetic acid, they give a yellow precipitate

with bichromate of potash. These reactions are characteristic of barium.

Baryta is got into solution in the manner described under _Lime_; but in

the case of the sulphate the substance is fused with three or four times

its weight of "fusion mixture." The "melt" is extracted with water,

washed, and the residue dissolved in dilute hydrochloric acid.

~Separation.~--The separation is thus effected:--The solution in

hydrochloric acid is evaporated to dryness, re-dissolved in hot dilute

hydrochloric acid, and sulphuric acid is added to the solution till no

further precipitate is formed. The precipitate is filtered off, and

digested with a solution of ammonium acetate or of sodium hyposulphite

at 50° or 60° C. to dissolve out any lead sulphate. The residue is

filtered off, washed, dried, and ignited. The ignited substance is mixed

with four or five times its weight of "fusion mixture," and fused in a

platinum-dish over the blowpipe for a few minutes. When cold, it is

extracted with cold water, filtered, and washed. The residue is

dissolved in dilute hydrochloric acid, and (if necessary) filtered. The

solution contains the barium as baric chloride mixed, perhaps, with

salts of strontium or lime. To separate these, ammonia is added till the

solution is alkaline, and then acetic acid in slight excess. Chromate of

baryta is then thrown down, by the addition of bichromate of potash, as

a yellow precipitate. It is allowed to settle, filtered and washed with

a solution of acetate or of nitrate of ammonia. It is dried, ignited

gently, and weighed. It is BaCrO_{4}, and contains 60.47 per cent. of

baryta.

GRAVIMETRIC DETERMINATION.

The gravimetric determination of baryta, when lime and strontia are

absent, is as follows:--The solution, if it contains much free acid, is

nearly neutralised with ammonia, and then diluted to 100 or 200 c.c. It

is heated to boiling, and dilute sulphuric acid is added till no further

precipitation takes place. The precipitate is allowed to settle for a

few minutes, decanted through a filter, and washed with hot water; and,

afterwards, dried, transferred to a porcelain crucible, and strongly

ignited in the muffle or over the blowpipe for a few minutes. It is then

cooled, and weighed as sulphate of baryta (BaSO_{4}). It contains 65.67

per cent. of baryta (BaO).

In determining the baryta in minerals which are soluble in acid, it is

precipitated direct from the hydrochloric acid solution (nearly

neutralised with ammonia) by means of sulphuric acid. The precipitated

baric sulphate is digested with a solution of ammonic acetate; and

filtered, washed, ignited, and weighed.

VOLUMETRIC DETERMINATION.

The principle and mode of working of this is the same as that given

under the Sulphur Assay; but using a standard solution of sulphuric acid

instead of one of barium chloride. The standard solution of sulphuric

acid is made to contain 32.02 grams of sulphuric acid (H_{2}SO_{4}), or

an equivalent of a soluble alkaline sulphate, per litre. 100 c.c. will

be equal to 5 grams of baryta.

Five grams of the substance are taken, and the baryta they contain

converted into carbonate (if necessary). The carbonate is dissolved in

dilute hydrochloric acid. Ten grams of sodium acetate are added, and the

solution, diluted to 500 c.c., is boiled, and titrated in the manner

described.

Lead salts must be absent in the titration, and so must strontia and

lime. Ferrous salts should be peroxidised by means of permanganate or

chlorate of potash. Other salts do not interfere.

MAGNESIA.

Magnesia, the oxide of magnesium (MgO) occurs in nature in the rare

mineral periclase (MgO); and hydrated, as brucite (MgH_{2}O_{2}). As

carbonate it occurs in large quantity as magnesite (MgCO_{3}), which is

the chief source of magnesia. Mixed with carbonate of lime, it forms

magnesian limestone and dolomite. It is present in larger or smaller

quantity in most silicates; and the minerals, serpentine, talc, steatite

and meerschaum are essentially hydrated silicates of magnesia. Soluble

magnesian salts occur in many natural waters; more especially the

sulphate and the chloride. Kieserite (MgSO_{4}.H_{2}O) occurs in

quantity at Stassfurt, and is used in the manufacture of Epsom salts.

~Detection.~--Magnesia is best detected in the wet way. Its compounds

give no colour to the flame, and the only characteristic dry reaction is

its yielding a pink mass when ignited before the blowpipe (after

treatment with a solution of cobalt nitrate). In solution, it is

recognised by giving no precipitate with ammonia or ammonic carbonate in

the presence of ammonic chloride, and by giving a white crystalline

precipitate on adding sodium phosphate or arsenate to the ammoniacal

solution.

Magnesia differs from the other alkaline earths by the solubility of its

sulphate in water.

Magnesia is dissolved by boiling with moderately strong acids; the

insoluble compounds are fused with "fusion mixture," and treated as

described under _Silicates_.

~Separation.~--It is separated by evaporating the acid solution to

dryness to render silica insoluble, and by taking up with dilute

hydrochloric acid. The solution is freed from the second group of metals

by means of sulphuretted hydrogen, and the iron, alumina, &c., are

removed with ammonic chloride, ammonia, and ammonic sulphide. The

somewhat diluted filtrate is treated, first, with ammonia, and then with

carbonate of ammonia in slight excess. It is allowed to stand for an

hour in a warm place, and then filtered. The magnesia is precipitated

from the filtrate by the addition of an excess of sodium phosphate and

ammonia. It is allowed to stand overnight, filtered, and washed with

dilute ammonia. The precipitate contains the magnesia as

ammonic-magnesic phosphate.

In cases where it is not desirable to introduce sodium salts or

phosphoric acid into the assay solution, the following method is

adopted. The solution (freed from the other alkaline earths by ammonium

carbonate) is evaporated in a small porcelain dish with nitric acid. The

residue (after removing the ammonic salts by ignition) is taken up with

a little water and a few crystals of oxalic acid, transferred to a

platinum dish, evaporated to dryness, and ignited. The residue is

extracted with small quantities of boiling water and filtered off; while

the insoluble magnesia is washed. The filtrate contains the alkalies.

The residue is ignited, and weighed as magnesia. It is MgO.

GRAVIMETRIC DETERMINATION.

The solution containing the magnesia is mixed with chloride of ammonium

and ammonia in excess. If a precipitate should form, more ammonic

chloride is required. Add sodium phosphate solution in excess, stir and

allow to stand overnight. Filter and wash the precipitate with dilute

ammonia. Dry, transfer to a platinum or porcelain crucible, and ignite

(finally at intense redness); cool, and weigh. The substance is magnesic

pyrophosphate (Mg_{2}P_{2}O_{7}), and contains 36.04 per cent. of

magnesia.

VOLUMETRIC METHOD.

The magnesia having been precipitated as ammonic-magnesic phosphate,

which is the usual separation, its weight can be determined

volumetrically by the method of titration described under _Phosphates_.

The same standard solution of uranium acetate is used. Its standard for

magnesia is got by multiplying the standard for phosphoric oxide by

0.5493. For example, if one hundred c.c. are equivalent to 0.5 gram of

phosphoric oxide, they will be equivalent to (0.5 × .5493) 0.2746 gram

of magnesia. The method of working and the conditions of the titration

are the same as for the phosphate titration. The quantity of substance

taken for assay must not contain more than 0.1 or 0.2 gram of magnesia.

After precipitating as ammonic-magnesic phosphate with sodium phosphate,

and well washing with ammonia, it is dissolved in dilute hydrochloric

acid, neutralised with ammonia, and sodic acetate and acetic acid are

added in the usual quantity. The solution is boiled and titrated.

EXAMINATION OF A LIMESTONE.

~Silica and Insoluble Silicates.~--Take one gram of the dried sample and

dissolve it in 10 c.c. of dilute hydrochloric acid; filter; wash, dry,

and ignite the residue.

~Organic Matter.~--If the residue insoluble in hydrochloric acid shows

the presence of organic matter, it must be collected on a weighed filter

and dried at 100°. On weighing, it gives the combined weights of organic

and insoluble matter. The latter is determined by igniting and weighing

again. The organic matter is calculated by difference.

~Lime.~--Where but little magnesia is present, this is determined by

titration with standard acid. Take one gram, and dissolve it in 25 c.c.

of normal hydrochloric acid. Tint with methyl-orange and titrate with

semi-normal ammonia. Divide the quantity of ammonia used by 2, deduct

this from 25, and multiply the remainder by 2.8. This gives the

percentage of lime. Where magnesia is present, the same method is

adopted, and the magnesia (which is separately determined) is afterwards

deducted. The percentage of magnesia found is multiplied by 1.4, and the

result is deducted from the apparent percentage of lime got by

titrating.

~Magnesia.~--Dissolve 2 grams of the limestone in hydrochloric acid, and

separate the lime with ammonia and ammonium oxalate. The filtrate is

treated with sodium phosphate, and the magnesia is weighed as

pyrophosphate, or titrated with uranium acetate.

~Iron.~--Dissolve 2 grams in hydrochloric acid, reduce, and titrate with

standard permanganate of potassium solution. This gives the total iron.

The ferrous iron is determined by dissolving another 2 grams in

hydrochloric acid and at once titrating with the permanganate of

potassium solution.

~Manganese.~--Dissolve 20 grams in hydrochloric acid, nearly neutralise

with soda, add sodium acetate, boil, and filter. To the filtrate add

bromine; boil, and determine the manganese in the precipitate. See page

300.

~Phosphoric Oxide.~--This is determined by dissolving the ferric acetate

precipitate from the manganese separation in hydrochloric acid, adding

ammonia in excess, and passing sulphuretted hydrogen. Filter and add to

the filtrate "magnesia mixture." The precipitate is collected, washed

with ammonia, ignited, and weighed as pyrophosphate.

THE ALKALIES.

The oxides of sodium, potassium, lithium, cæsium, and rubidium and

ammonia are grouped under this head. Of these cæsia and rubidia are

rare, and lithia comparatively so. They are easily distinguished by

their spectra. They are characterised by the solubility of almost all

their salts in water, and, consequently, are found in the solutions from

which the earths and oxides of the metals have been separated by the

usual group re-agents.

The solution from which the other substances have been separated is

evaporated to dryness, and the product ignited to remove the ammonic

salts added for the purpose of separation. The residue contains the

alkali metals generally, as chlorides or sulphates. Before determining

the quantities of the particular alkali metals present, it is best to

convert them altogether, either into chloride or sulphate, and to take

the weight of the mixed salts. It is generally more convenient to weigh

them as chlorides. They are converted into this form, if none of the

stronger acids are present, by simply evaporating with an excess of

hydrochloric acid. Nitrates are converted into chlorides by this

treatment. When sulphates or phosphates are present, the substance is

dissolved in a little water, and the sulphuric or phosphoric acid

precipitated with a slight excess of acetate of lead in the presence of

alcohol. The solution is filtered, and the excess of lead precipitated

with sulphuretted hydrogen. The filtrate from this is evaporated to

dryness with an excess of hydrochloric acid, and the residue, consisting

of the mixed chlorides, is gently ignited and weighed. In many cases

(such as the analysis of slags and of some natural silicates where the

percentage of alkalies is small) the percentage of soda and potash

(which most commonly occur) need not be separately determined. It is

sufficient to report the proportion of mixed alkalies; which is thus

ascertained:--Dissolve the ignited and weighed chlorides in 100 c.c. of

distilled water, and titrate with the standard solution of silver

nitrate (using potassic chromate as indicator) in the manner described

under _Chlorine_. The c.c. of silver nitrate used gives the weight in

milligrams of the chlorine present. Multiply this by 0.775, and deduct

the product from the weight of the mixed chlorides. This will give the

combined weight of the alkalies (Na_{2}O and K_{2}O) present. For

example, 0.0266 gram of mixed chlorides required on titrating 14.2 c.c.

of silver nitrate, which is equivalent to 0.0142 gram of chlorine. This

multiplied by 0.775 gives 0.0110 to be deducted from the weight of the

mixed chlorides.

Mixed chlorides 0.0266 gram

Deduction 0.0110 "

------

Mixed alkalies 0.0156 "

Assuming this to have been got from 1 gram of a rock, it would amount to

1.56 per cent. of "potash and soda."

The relative proportions of the potash and soda can be ascertained from

the same determination. Sodium and potassium chlorides have the

following composition:--

Sodium 39.38 Potassium 52.46

Chlorine 60.62 Chlorine 47.54

----- -----

100.00 100.00

The percentage of chlorine in the mixed chlorides is calculated. It

necessarily falls somewhere between 47.5 and 60.6 per cent., and

approaches the one or the other of these numbers as the proportion of

the sodium or potassium preponderates. Each per cent. of chlorine in

excess of 47.5 represents 7.63 per cent. of sodium chloride in the mixed

chlorides. The percentage of potash and soda in the substance can be

calculated in the usual way. Sodium chloride multiplied by 0.5302 gives

its equivalent of soda (Na_{2}O), and potassium chloride multiplied by

0.6317 gives its equivalent of potash (K_{2}O).

The weight of sodium chloride in the mixed chlorides is also calculated

thus:--Take the same example for illustration. Multiply the chlorine

found by 2.103. This gives--

(0.0142×2.103) = 0.02987.

From the product deduct the weight of the mixed chlorides found--

Product 0.02987

Mixed chlorides 0.02660

-------

Difference 0.00327

The difference multiplied by 3.6288 gives the weight of sodium chloride

in the mixture. In this case it equals 0.0118 gram. The potassium

chloride is indicated by the difference between this and the weight of

the mixed chlorides. It equals 0.0148 gram. We have now got--

Sodium chloride 0.0118 gram

Potassium chloride 0.0148 "

from 1 gram of the rock taken. Multiplying these by their factors we

have (Soda = 0.0118×0.5302; Potash 0.0148×0.6317)--

Soda = 0.625 per cent.

Potash = 0.935 "

~Concentration of the Alkalies.~--With the exception of magnesia, all

the other bases are separated from the alkalies in the ordinary course

of work without the addition of any re-agent which cannot be removed by

simple evaporation and ignition. Consequently, with substances soluble

in acids, successive treatment of the solution with sulphuretted

hydrogen, ammonia, ammonic sulphide, and ammonic carbonate, filtering,

where necessary, will yield a filtrate containing the whole of the

alkalies with ammonic salts and, perhaps, magnesia.

The filtrate is evaporated in a small porcelain dish, with the addition

of nitric acid towards the finish. It is carried to dryness and ignited.

The residue is taken up with a little water, treated with a few crystals

of oxalic acid, and again evaporated and ignited. The alkaline salts are

extracted with water, and filtered from the magnesia into a weighed

platinum dish. The solution is then evaporated with an excess of

hydrochloric acid, ignited at a low red heat, and weighed. The residue

consists of the mixed alkaline chlorides.

For substances (such as most silicates and similar bodies) not

completely decomposed by acids, Lawrence Smith's method is generally

used. This is as follows:--Take from 0.5 to 1 gram of the finely

powdered mineral, and mix, by rubbing in the mortar, with an equal

weight of ammonium chloride. Then mix with eight times as much pure

calcium carbonate, using a part of it to rinse out the mortar. Transfer

to a platinum crucible, and heat gently over a Bunsen burner until the

ammonic chloride is decomposed (five or ten minutes). Raise the heat to

redness, and continue at this temperature for about three quarters of an

hour. The crucible must be kept covered. Cool, and turn out the mass

into a 4-inch evaporating dish; wash the crucible and cover with

distilled water, and add the washings to the dish; dilute to 60 or 80

c.c., and heat to boiling. Filter and wash. Add to the filtrate about

1.5 gram of ammonium carbonate; evaporate to about 40 c.c., and add a

little more ammonic carbonate and some ammonia. Filter into a weighed

platinum dish, and evaporate to dryness. Heat gently, to drive off the

ammonic chloride, and ignite to a little below redness. Cool and weigh.

The residue consists of the mixed alkaline chlorides.

~Separation of the Alkali-Metals from each other.~--Sodium and lithium

are separated from the other alkali-metals by taking advantage of the

solubility of their chlorides in the presence of platinic chloride; and

from one another by the formation of an almost insoluble lithic

phosphate on boiling with a solution of sodium phosphate in a slightly

alkaline solution. Cæsium, rubidium, and potassium yield precipitates

with platinic chloride, which are somewhat soluble, and must be

precipitated from concentrated solutions. Cæsium and rubidium are

separated from potassium by fractional precipitation with platinum

chloride. Their platino-chlorides, being less soluble than that of

potassium, are precipitated first. One hundred parts of boiling water

dissolve 5.18 of the potassium platino-chloride, 0.634 of the rubidium

salt, and 0.377 of the corresponding cæsium compound. The separation of

lithium, cæsium, and rubidium is seldom called for, owing to their

rarity. The details of the separation of potassium from sodium are

described under _Potassium_. Ammonia compounds are sharply marked off

from the rest by their volatility, and it is always assumed that they

have been removed by ignition; if left in the solution, they would count

as potassium compounds. They will be considered under _Ammonia_.

SODIUM.

Sodium is the commonest of the alkali metals. It is found in nature

chiefly combined with chlorine as "common salt" (NaCl). This mineral is

the source from which the various compounds of sodium in use are

prepared. Sodium occurs abundantly as nitrate (NaNO_{3}) in Chili

saltpetre, and as silicate in various minerals, such as albite (or

soda-felspar).

It occurs as fluoride in cryolite (Na_{3}AlF_{6}), and as carbonate in

natron, &c. Sulphates are also found. Sodium is very widely diffused,

few substances being free from it.

The detection of sodium is easy and certain, owing to the strong yellow

colour its salts impart to the flame; this, when viewed by the

spectroscope, shows a single yellow line.[93] The extreme delicacy of

this test limits its value, because of the wide diffusion of sodium

salts. It is more satisfactory to separate the chloride, which may be

recognised by its taste, flame coloration, fusibility, and negative

action with reagents. The chloride dissolved in a few drops of water

gives with potassium metantimoniate, a white precipitate of the

corresponding sodium salt.

Sodium salts are dissolved out from most compounds on treatment with

water or dilute acids. Insoluble silicates are decomposed and the alkali

rendered soluble by Lawrence Smith's method, which has just been

described. The separation of the sodium from the mixed chlorides is

effected in the following way:--The chlorides are dissolved in a little

water and the potassium separated as platino-chloride. The soluble

sodium platino-chloride, with the excess of platinum, is boiled, mixed

with sulphuric acid, evaporated to dryness, and ignited. On extracting

with water, filtering, evaporating, and igniting, sodium sulphate is

left, and is weighed as such.

It is more usual, and quite as satisfactory, to calculate the weight of

the sodium chloride by difference from that of the mixed chlorides, by

subtracting that of the potassium chloride, which is separately

determined. For example, 1 gram of a rock gave--Mixed chlorides, 0.0266

gram, and 0.0486 gram of potassic platino-chloride. This last is

equivalent to 0.0149 gram of potassium chloride.

Mixed chlorides found 0.0266

Deduct potassium chloride 0.0149

------

Leaves sodium chloride 0.0117

The weight of sodium chloride found, multiplied by 0.5302, gives the

weight of the soda (Na_{2}O).

GRAVIMETRIC DETERMINATION.

The solution, which must contain no other metal than sodium, is

evaporated in a weighed platinum crucible or dish. Towards the finish an

excess, not too great, of sulphuric acid is added, and the evaporation

is continued under a loosely fitting cover. The residue is ignited over

the blowpipe, a fragment of ammonic carbonate being added towards the

end, when fumes of sulphuric acid cease to be evolved. This ensures the

removal of the excess of acid. The crucible is cooled in the desiccator,

and weighed. The substance is sulphate of soda (Na_{2}SO_{4}), and

contains 43.66 per cent. of soda (Na_{2}O), or 32.38 per cent. of sodium

(Na).

VOLUMETRIC METHODS.

There are various methods used for the different compounds of sodium.

There is no one method of general application. Thus with "common salt"

the chlorine is determined volumetrically; and the sodium, after

deducting for the other impurities, is estimated by difference.

With sodic carbonate and caustic soda, a given weight of the sample is

titrated with standard acid, and the equivalent of soda estimated from

the alkalinity of the solution.

With sodium sulphate, a modification of the same method is used. To a

solution of 3.55 grams of the salt contained in a half-litre flask, 250

c.c. of a solution of baryta water is added. The volume is made up to

500 c.c. with water. The solution is mixed and filtered. Half of the

filtrate is measured off, treated with a current of carbonic acid, and

then boiled. It is transferred to a half-litre flask, diluted to the

mark, shaken up, and filtered. 250 c.c. of the filtrate, representing a

quarter of the sample taken, is then titrated with standard acid. The

standard acid is made by diluting 250 c.c. of the normal acid to 1

litre. The c.c. of acid used multiplied by 2 gives the percentage. A

correction must be made to counteract the effect of impurities in the

baryta as well as errors inherent in the process. This is small, and its

amount is determined by an experiment with 3.55 grams of pure sodium

sulphate.

EXAMINATION OF COMMON SALT.

~Moisture.~--Powder and weigh up 10 grams of the sample into a platinum

dish. Dry in a water oven for an hour, and afterwards heat to bare

redness over a Bunsen burner. Cool, and weigh. The loss gives the water.

~Chlorine.~--Weigh up two separate lots of 1 gram each; dissolve in 100

c.c. of water, and determine the chlorine by titrating with the standard

silver nitrate solution, using chromate of potash as indicator. See

_Chlorine_.

~Insoluble Matter.~--Dissolve 10 grams of the salt in water with the

help of a little hydrochloric acid. Filter off the sediment, wash,

ignite, and weigh. This residue is chiefly sand. Dilute the nitrate to

500 c.c.

~Lime.~--Take 250 c.c. of the filtrate, render ammoniacal and add

ammonium oxalate; wash, dry, and ignite the precipitate. Weigh as lime

(CaO).

~Magnesia.~--To the filtrate from the lime add phosphate of soda. Allow

to stand overnight, filter, wash with dilute ammonia, dry, ignite, and

weigh as pyrophosphate.

~Sulphuric Oxide.~--To the remaining 250 c.c. of the filtrate from the

"insoluble," add an excess of barium chloride. Collect, wash, dry,

ignite, and weigh the barium sulphate.

~Sodium.~--It is estimated by difference.

The following may be taken as an example:--

Moisture 0.35

Insoluble matter 0.40

Lime 0.40

Magnesia 0.05

Sulphuric oxide 0.60

Chlorine 59.60

Sodium 38.60

------

100.00

POTASSIUM.

Potassium occurs in nature as chloride, in the mineral sylvine (KCl),

and more abundantly combined with magnesium chloride, in earnallite

(KCl.MgCl_{2}.6H_{2}O). It occurs as nitrate in nitre (KNO_{3}), and as

silicate in many minerals, such as orthoclase (or potash-felspar) and

muscovite (or potash-mica).

Potassium compounds are detected by the characteristic violet colour

they impart to the flame. The presence of sodium salts masks this tint,

but the interference can be neutralised by viewing the flame through a

piece of blue glass. Viewed through the spectroscope, it shows a

characteristic line in the red and another in the violet. These,

however, are not so easy to recognise or obtain as the sodium one.

Concentrated solutions of potassium salts give a yellow crystalline

precipitate with platinum chloride, and a white crystalline one with the

acid tartrate of soda. For these tests the solution is best neutral.

These tests are only applicable in the absence of compounds other than

those of potassium and sodium.

GRAVIMETRIC DETERMINATION.

This process serves for its separation from sodium. Take 1 gram of the

sample and dissolve it in an evaporating dish with 50 c.c. of water.

Acidify with hydrochloric acid in quantity sufficient (if the metals are

present as chlorides) to make it acid, or, if other acids are present,

in at least such quantity as will provide the equivalent of chlorine.

Add 3 grams of platinum, in solution as platinum chloride, and evaporate

on a water-bath to a stiff paste, but not to dryness. Moisten with a few

drops of platinic chloride solution without breaking up the paste by

stirring. Cover with 20 c.c. of strong alcohol, and wash the crystals as

much as possible by rotating the dish. Allow to settle for a few

moments, and decant through a filter. Wash in the same way two or three

times until the colour of the filtrate shows that the excess of the

platinum chloride used is removed. Wash the precipitate on to the filter

with a jet of alcohol from the wash-bottle; clean the filter-paper,

using as little alcohol as possible. Dry in the water-oven for an hour.

Brush the precipitate into a weighed dish, and weigh it. It is potassium

platino-chloride (K_{2}PtCl_{6}), and contains 16.03 per cent. of

potassium, or 30.56 per cent. of potassium chloride (KCl), which is

equivalent to 19.3 per cent. of potash (K_{2}O).

If the filter-paper is not free from precipitate, burn it and weigh

separately. The excess of weight over that of the ash will be due to

platinum and potassic chloride (Pt and 2KCl). This multiplied by 1.413

will give the weight of the potassic platino-chloride from which it was

formed. It must be added to the weight of the main precipitate.

The mixed alkaline chlorides obtained in the usual course of analysis

are treated in this manner; the quantity of platinum added must be about

three times as much as the mixed chlorides weigh.

VOLUMETRIC METHODS.

These are the same as with soda.

~Examination of Commercial Carbonate of Potash.~--The impurities to be

determined are moisture, silica, and insoluble matter, chlorine,

sulphuric oxide, and oxide of iron. These determinations are made in the

ways described under the examination of common salt.

The ~potassium~ is determined by converting it into chloride and

precipitating with platinum chloride, &c., as just described.

~Available Alkali.~--Weigh up 23.5 grams of the sample, dissolve in

water, and make up to 500 c.c. Take 50 c.c., tint with methyl orange,

and titrate with the normal solution of acid. The c.c. of acid used

multiplied by 2 gives the percentage of available alkali calculated as

potash (K_{2}O).

~Soda.~--This is calculated indirectly in the following way:--Deduct

from the potassium found the quantity required for combination with the

chlorine and sulphuric oxide present, and calculate the remainder to

potash (K_{2}O). The apparent surplus excess of available alkali is the

measure of the soda present.

~Carbon Dioxide.~--The c.c. of acid used in the available alkali

determination, multiplied by 2.2 and divided by 2.35, gives the

percentage of carbon dioxide.

LITHIUM.

Lithia, the oxide of lithium (Li_{2}O), occurs in quantities of 3 or 4

per cent. in various silicates, such as lepidolite (or lithia-mica),

spodumene, and petalite. It also occurs as phosphate in triphyline. It

is a constituent of the water of certain mineral springs. A spring at

Wheal Clifford contained as much as 0.372 gram of lithium chloride per

litre. In small quantities, lithia is very widely diffused.

The ~Detection~ of lithia is rendered easy by the spectroscope; its

spectrum shows a red line lying about midway between the yellow sodium

line and the red one of potassium. It also shows a faint yellow line.

The colour of the flame (a crimson) is characteristic.

The reactions of the lithium compounds lie between those of the alkalies

and of the alkaline earths. Solutions are not precipitated by tartaric

acid nor by platinic chloride. The oxide is slowly soluble in water.

The carbonate is not freely soluble. Lithia is completely precipitated

by sodic phosphate, especially in hot alkaline solutions.

In its determination the mixed alkaline chlorides obtained in the

separation of the alkalies are dissolved in water, a solution of soda is

added in slight excess, and the lithia precipitated with _sodic_

phosphate. Before filtering, it is evaporated to dryness and extracted

with hot water rendered slightly ammoniacal. The residue is transferred

to a filter, dried, ignited, and weighed. The precipitate is lithium

phosphate (3Li_{2}O, P_{2}O_{5}), and contains 38.8 per cent. of lithia.

The separation of lithia from magnesia is not given by the usual

authorities. Wohler recommends evaporating the solution to dryness with

carbonate of soda. On extracting the residue with water, the lithia

dissolves out and is determined in the filtrate. One hundred parts of

water dissolve, at the ordinary temperature, 0.769 parts of lithium

carbonate (Li_{2}CO_{3}); the basic magnesia compound is almost

insoluble in the absence of carbon dioxide and ammonium salts.

CAESIUM.

The oxide of caesium, caesia (Cs_{2}O), is found associated with lithia

in lepidolite, &c., and, together with rubidium, in many mineral waters.

The mineral pollux is essentially a silicate of alumina and caesia; it

contains 34.0 per cent. of the latter oxide.

Caesium is best detected by the spectroscope, its spectrum being

characterised by two lines in the blue and one in the red; the latter is

about midway between the lithium and sodium lines.

If not detected by the spectroscope, or specially looked for, caesia

would, in the ordinary course of work, be separated with the potash and

weighed as potassium platino-chloride.

Caesia is separated from all the other alkalies by adding to the acid

solution of the mixed chlorides a strongly acid cold solution of

antimonious chloride. The acid used must be hydrochloric. The caesium is

precipitated as a white crystalline precipitate (CsCl.SbCl_{3}), which

is filtered off, and washed, when cold, with strong hydrochloric acid;

since it is decomposed by water or on warming. The precipitate is washed

into a beaker, and treated with sulphuretted hydrogen; after filtering

off the sulphide of antimony, the solution leaves, on evaporation, the

caesium as chloride.

RUBIDIUM.

Rubidium occurs widely diffused in nature, but in very small quantities.

It is generally associated with caesium.

It is detected by the spectroscope, which shows two violet lines and two

dark red ones. Like caesium, it is precipitated with platinic chloride,

and in the ordinary course of work would be weighed as potassium. It is

separated from potassium by fractional precipitation with platinic

chloride. Rubidium platino-chloride is much less soluble than the

potassium salt.

AMMONIUM.

It is usual to look upon the salts of ammonia as containing a compound

radical (NH_{4} = Am), which resembles in many respects the metals of

the alkalies. Ammonium occurs in nature as chloride in sal ammoniac

(AmCl), as sulphate in mascagnine (Am_{2}SO_{4}), as phosphate in

struvite (AmMgPO_{4}.12H_{2}O). Minerals containing ammonium are rare,

and are chiefly found either in volcanic districts or associated with

guano. Ammonia and ammonium sulphide occur in the waters of certain

Tuscan lagoons, which are largely worked for the boracic acid they

contain. The crude boracic acid from this source contains from 5 to 10

per cent. of ammonium salts. It is from these that the purer forms of

ammonium compounds of commerce known as "from volcanic ammonia" are

derived. But the bulk of the ammonia of commerce is prepared from the

ammoniacal liquors obtained as bye-products in the working of certain

forms of blast furnaces and coke ovens, and more especially in

gas-making.

Ammonia hardly comes within the objects of assaying; but it is largely

used in the laboratory, and the assayer is not unfrequently called on to

determine it. Ammonium salts are mostly soluble in water. In strong

solutions they give a yellow precipitate of ammonium platino-chloride on

the addition of chloride of platinum; and with the acid tartrate of soda

yield a white precipitate of hydric ammonic tartrate. These reactions

are similar to those produced with potassium compounds.

Heated with a base, such as lime or sodic hydrate, ammonium salts are

decomposed, yielding ammonia gas (NH_{3}), which is readily soluble in

water. The solution of this substance is known as ammonic hydrate or

"ammonia."

They are volatilised on ignition; either with, or without, decomposition

according to the acid present. This fact is of importance in analytical

work; since it allows of the use of alkaline solutions and reagents

which leave nothing behind on heating. It must be remembered, however,

that, although ammonic chloride is volatile, it cannot be volatilised in

the presence of substances which form volatile chlorides without loss of

the latter. For example: ferric oxide and alumina are thus lost,

volatilising as chlorides; and there are some other compounds (notably

ammonic magnesic arsenate) which on heating to redness suffer reduction.

The presence of ammonic chloride in such cases must be avoided.

~Detection.~--Compounds of ammonium are detected by their evolving

ammonia when mixed or heated with any of the stronger bases. The ammonia

is recognised by its odour, by its alkaline reaction with litmus paper,

and by yielding white fumes, when brought in contact with fuming acid.

In consequence of the use of ammonium salts and ammonia as reagents, it

is necessary to make a special test for and determination of

ammonium.[94] In the ordinary course of work it will be "lost on

ignition." The determination presents little difficulty, and is based on

the method used for its detection.

[Illustration: FIG. 61.]

~Solution and Separation.~--Although ammonium salts are soluble in

water, there is no necessity for dissolving them. The compound

containing the ammonia is boiled with an alkaline solution; and the

liberated ammonia condensed and collected. The substance is weighed out

into a flask of about 200 c.c. capacity. The flask is closed with a

rubber cork perforated to carry a 20 c.c. pipette and a bulb exit tube.

The latter is connected with a receiver, which is a small flask

containing dilute hydrochloric acid (fig. 61). The flask containing the

substance is corked, and the greater part of the soda solution is run in

from the pipette. The solution is then boiled. The ammonia volatilises,

and is carried over into the hydrochloric acid, with which it combines

to form ammonic chloride. The distillation is carried on gently until

the bulk of the liquid is driven over. The ammonia in the receiver will

be mixed only with the excess of hydrochloric acid. This separation is

used in all determinations.

GRAVIMETRIC DETERMINATION.

The contents of the flask are transferred to a weighed platinum dish,

and evaporated on the water-bath. It is dried until the weight is

constant. The chloride of ammonium remains as a white mass which, after

cooling in a desiccator, is weighed. It contains 33.72 per cent. of

ammonium (NH_{4}), or 31.85 per cent. of ammonia (NH_{3}). On heating

over the Bunsen burner it is completely volatilised, leaving no residue.

VOLUMETRIC DETERMINATION.

Weigh up 1.7 gram of the substance and place it in the flask. Measure

off 50 c.c. of the normal solution of acid, place them in the receiver,

and dilute with an equal volume of water. Run in through the pipette (by

opening the clip) 20 c.c. of a strong solution of soda, boil until the

ammonia has passed over, and then aspirate a current of air through the

apparatus. Disconnect the receiver, and tint its contents with methyl

orange. Titrate the residual acid with a semi-normal solution of alkali.

Divide the c.c. of the "alkali" solution used by 2, and deduct from the

50 c.c. The difference will give the number of c.c. of the normal acid

solution neutralised by the ammonia distilled over. Each c.c. of "acid"

so neutralised, represents 1 per cent. of ammonia in the sample. If the

results are to be reported as ammonium, 1.8 gram of the sample is taken

instead of 1.7 gram.

COLORIMETRIC DETERMINATION.

This is effected by means of "Nessler's" reagent, which strikes a brown

colour with traces of ammonia, even with a few hundredths of a milligram

in 100 c.c. of liquid. With larger quantities of ammonia the reagent

gives a precipitate. This reagent is a strongly alkaline solution of

potassic mercuric iodide; and is thus made:--

_Nessler's solution_: Dissolve 17 grams of mercuric chloride in 300 c.c.

of water; and add the solution to one of 35 grams of potassium iodide in

100 c.c. of water until a permanent precipitate is produced. Both

solutions must be cold. Then make up to a litre by adding a 20 per cent.

solution of potash. Add more of the mercuric chloride (a little at a

time) until a permanent precipitate is again formed. Allow to settle,

decant, and use the clear liquor. Four or five c.c. are used for each

100 c.c. of liquid to be tested.