THE PHOSPHORUS FAMILY
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| | ATOMIC | | MELTING
| SYMBOL | WEIGHT | DENSITY | POINT
-----------+--------+---------+---------+---------
Phosphorus | P | 31.0 | 1.8 | 43.3°
Arsenic | As | 75.0 | 5.73 | ---
Antimony | Sb | 120.2 | 6.7 | 432°
Bismuth | Bi | 208.5 | 9.8 | 270°
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~The family.~ The elements constituting this family belong in the same
group with nitrogen and therefore resemble it in a general way. They
exhibit a regular gradation of physical properties, as is shown in the
above table. The same general gradation is also found in their chemical
properties, phosphorus being an acid-forming element, while bismuth is
essentially a metal. The other two elements are intermediate in
properties.
~Compounds.~ In general the elements of the family form compounds having
similar composition, as is shown in the following table:
PH_{3} PCl_{3} PCl_{5} P_{2}O_{3} P_{2}O_{5}
AsH_{3} AsCl_{3} AsCl_{5} As_{2}O_{3} As_{2}O_{5}
SbH_{3} SbCl_{3} SbCl_{5} Sb_{2}O_{3} Sb_{2}O_{5}
.... BiCl_{3} BiCl_{5} Bi_{2}O_{3} Bi_{2}O_{5}
In the case of phosphorus, arsenic, and antimony the oxides are acid
anhydrides. Salts of at least four acids of each of these three elements
are known, the free acid in some instances being unstable. The relation
of these acids to the corresponding anhydrides may be illustrated as
follows, phosphorus being taken as an example:
P_{2}O_{3} + 3H_{2}O = 2H_{3}PO_{3} (phosphorous acid).
P_{2}O_{5} + 3H_{2}O = 2H_{3}PO_{4} (phosphoric acid).
P_{2}O_{5} + 2H_{2}O = H_{4}P_{2}O_{7} (pyrophosphoric acid).
P_{2}O_{5} + H_{2}O = 2HPO_{3} (metaphosphoric acid).
PHOSPHORUS
~History.~ The element phosphorus was discovered by the alchemist Brand,
of Hamburg, in 1669, while searching for the philosopher's stone. Owing
to its peculiar properties and the secrecy which was maintained about
its preparation, it remained a very rare and costly substance until the
demand for it in the manufacture of matches brought about its production
on a large scale.
~Occurrence.~ Owing to its great chemical activity phosphorus never occurs
free in nature. In the form of phosphates it is very abundant and widely
distributed. _Phosphorite_ and _sombrerite_ are mineral forms of calcium
phosphate, while _apatite_ consists of calcium phosphate together with
calcium fluoride or chloride. These minerals form very large deposits
and are extensively mined for use as fertilizers. Calcium phosphate is a
constituent of all fertile soil, having been supplied to the soil by the
disintegration of rocks containing it. It is the chief mineral
constituent of bones of animals, and bone ash is therefore nearly pure
calcium phosphate.
~Preparation.~ Phosphorus is now manufactured from bone ash or a pure
mineral phosphate by heating the phosphate with sand and carbon in an
electric furnace. The materials are fed in at M (Fig. 70) by the feed
screw F. The phosphorus vapor escapes at P and is condensed under
water, while the calcium silicate is tapped off as a liquid at S. The
phosphorus obtained in this way is quite impure, and is purified by
distillation.
[Illustration: Fig. 70]
~Explanation of the reaction.~ To understand the reaction which
occurs, it must be remembered that a volatile acid anhydride is
expelled from its salts when heated with an anhydride which is
not volatile. Thus, when sodium carbonate and silicon dioxide
are heated together the following reaction takes place:
Na_{2}CO_{3} + SiO_{2} = Na_{2}SiO_{3} + CO_{2}.
Silicon dioxide is a less volatile anhydride than phosphoric
anhydride (P_{2}O_{5}), and when strongly heated with a
phosphate the phosphoric anhydride is driven out, thus:
Ca_{3}(PO_{4})_{2} + 3SiO_{2} = 3CaSiO_{3} + P_{2}O_{5}.
If carbon is added before the heat is applied, the P_{2}O_{5}
is reduced to phosphorus at the same time, according to the
equation
P_{2}O_{5} + 5C = 2P + 5CO.
~Physical properties.~ The purified phosphorus is a pale yellowish,
translucent, waxy solid which melts at 43.3° and boils at 269°. It can
therefore be cast into any convenient form under warm water, and is
usually sold in the market in the form of sticks. It is quite soft and
can be easily cut with a knife, but this must always be done while the
element is covered with water, since it is extremely inflammable, and
the friction of the knife blade is almost sure to set it on fire if cut
in the air. It is not soluble in water, but is freely soluble in some
other liquids, notably in carbon disulphide. Its density is 1.8.
~Chemical properties.~ Exposed to the air phosphorus slowly combines with
oxygen, and in so doing emits a pale light, or phosphorescence, which
can be seen only in a dark place. The heat of the room may easily raise
the temperature to the kindling point of phosphorus, when it burns with
a sputtering flame, giving off dense fumes of oxide of phosphorus. It
burns with dazzling brilliancy in oxygen, and combines directly with
many other elements, especially with sulphur and the halogens. On
account of its great affinity for oxygen it is always preserved under
water.
Phosphorus is very poisonous, from 0.2 to 0.3 gram being a fatal dose.
Ground up with flour and water or similar substances, it is often used
as a poison for rats and other vermin.
~Precaution.~ The heat of the body is sufficient to raise
phosphorus above its kindling temperature, and for this reason
it should always be handled with forceps and never with the
bare fingers. Burns occasioned by it are very painful and slow
in healing.
~Red phosphorus.~ On standing, yellow phosphorus gradually undergoes a
remarkable change, being converted into a dark red powder which has a
density of 2.1. It no longer takes fire easily, neither does it dissolve
in carbon disulphide. It is not poisonous and, in fact, seems to be an
entirely different substance. The velocity of this change increases with
rise in temperature, and the red phosphorus is therefore prepared by
heating the yellow just below the boiling point (250°-300°). When
distilled and quickly condensed the red form changes back to the yellow.
This is in accordance with the general rule that when a substance
capable of existing in several allotropic forms is condensed from a gas
or crystallized from the liquid state, the more unstable variety forms
first, and this then passes into the more stable forms.
~Matches.~ The chief use of phosphorus is in the manufacture of
matches. Common matches are made by first dipping the match
sticks into some inflammable substance, such as melted
paraffin, and afterward into a paste consisting of (1)
phosphorus, (2) some oxidizing substance, such as manganese
dioxide or potassium chlorate, and (3) a binding material,
usually some kind of glue. On friction the phosphorus is
ignited, the combustion being sustained by the oxidizing agent
and communicated to the wood by the burning paraffin. In
sulphur matches the paraffin is replaced by sulphur.
In safety matches _red_ phosphorus, an oxidizing agent, and
some gritty material such as emery is placed on the side of the
box, while the match tip is provided as before with an
oxidizing agent and an easily oxidized substance, usually
antimony sulphide. The match cannot be ignited easily by
friction, save on the prepared surface.
~Compounds of phosphorus with hydrogen.~ Phosphorus forms several
compounds with hydrogen, the best known of which is phosphine (PH_{3})
analogous to ammonia (NH_{3}).
~Preparation of phosphine.~ Phosphine is usually made by heating
phosphorus with a strong solution of potassium hydroxide, the reaction
being a complicated one.
[Illustration: Fig. 71]
The experiment can be conveniently made in the apparatus shown
in Fig. 71. A strong solution of potassium hydroxide together
with several small bits of phosphorus are placed in the flask
A, and a current of coal gas is passed into the flask through
the tube B until all the air has been displaced. The gas is
then turned off and the flask is heated. Phosphine is formed in
small quantities and escapes through the delivery tube, the
exit of which is just covered by the water in the vessel C.
Each bubble of the gas as it escapes into the air takes fire,
and the product of combustion (P_{2}O_{5}) forms beautiful
small rings, which float unbroken for a considerable time in
quiet air. The pure phosphine does not take fire spontaneously.
When prepared as directed above, impurities are present which
impart this property.
~Properties.~ Phosphine is a gas of unpleasant odor and is exceedingly
poisonous. Like ammonia it forms salts with the halogen acids. Thus we
have phosphonium chloride (PH_{4}Cl) analogous to ammonium chloride
(NH_{4}Cl). The phosphonium salts are of but little importance.
~Oxides of phosphorus.~ Phosphorus forms two well-known oxides,--the
trioxide (P_{2}O_{3}) and the pentoxide (P_{2}O_{5}), sometimes called
phosphoric anhydride. When phosphorus burns in an insufficient supply of
air the product is partially the trioxide; in oxygen or an excess of air
the pentoxide is formed. The pentoxide is much the better known of the
two. It is a snow-white, voluminous powder whose most marked property is
its great attraction for water. It has no chemical action upon most
gases, so that they can be very thoroughly dried by allowing them to
pass through properly arranged vessels containing phosphorus pentoxide.
~Acids of phosphorus.~ The important acids of phosphorus are the
following:
H_{3}PO_{3} phosphorous acid.
H_{3}PO_{4} phosphoric acid.
H_{4}P_{2}O_{7} pyrophosphoric acid.
HPO_{3} metaphosphoric acid.
These may be regarded as combinations of the oxides of phosphorus with
water according to the equations given in the discussion of the
characteristics of the family.
1. _Phosphorous acid_ (H_{3}PO_{3}). Neither the acid nor its salts are
at all frequently met with in chemical operations. It can be easily
obtained, however, in the form of transparent crystals when phosphorus
trichloride is treated with water and the resulting solution is
evaporated:
PCl_{3} + 3H_{2}O = H_{3}PO_{3} + 3HCl.
Its most interesting property is its tendency to take up oxygen and pass
over into phosphoric acid.
2. _Orthophosphoric acid (phosphoric acid)_ (H_{3}PO_{4}). This acid can
be obtained by dissolving phosphorus pentoxide in boiling water, as
represented in the equation
P_{2}O_{5} + 3H_{2}O = 2H_{3}PO_{4}.
It is usually made by treating calcium phosphate with concentrated
sulphuric acid. The calcium sulphate produced in the reaction is nearly
insoluble, and can be filtered off, leaving the phosphoric acid in
solution. Very pure acid is made by oxidizing phosphorus with nitric
acid. It forms large colorless crystals which are exceedingly soluble in
water. Being a tribasic acid, it forms acid as well as normal salts.
Thus the following compounds of sodium are known:
NaH_{2}PO_{4} monosodium hydrogen phosphate.
Na_{2}HPO_{4} disodium hydrogen phosphate.
Na_{3}PO_{4} normal sodium phosphate.
These salts are sometimes called respectively primary, secondary, and
tertiary phosphates. They may be prepared by bringing together
phosphoric acid and appropriate quantities of sodium hydroxide.
Phosphoric acid also forms mixed salts, that is, salts containing two
different metals. The most familiar compound of this kind is microcosmic
salt, which has the formula Na(NH_{4})HPO_{4}.
_Orthophosphates._ The orthophosphates form an important class of salts.
The normal salts are nearly all insoluble and many of them occur in
nature. The secondary phosphates are as a rule insoluble, while most of
the primary salts are soluble.
3. _Pyrophosphoric acid_ (H_{4}P_{2}O_{7}). On heating orthophosphoric
acid to about 225° pyrophosphoric acid is formed in accordance with the
following equation:
2H_{3}PO_{4} = H_{4}P_{2}O_{7} + H_{2}O.
It is a white crystalline solid. Its salts can be prepared by heating a
secondary phosphate:
2Na_{2}HPO_{4} = Na_{4}P_{2}O_{7} + H_{2}O.
4. _Metaphosphoric acid (glacial phosphoric acid)_ (HPO_{3}). This acid
is formed when orthophosphoric acid is heated above 400°:
H_{3}PO_{4} = HPO_{3} + H_{2}O.
It is also formed when phosphorus pentoxide is treated with cold water:
P_{2}O_{5} + H_{2}O = 2HPO_{3}.
It is a white crystalline solid, and is so stable towards heat that it
can be fused and even volatilized without decomposition. On cooling from
the fused state it forms a glassy solid, and on this account is often
called glacial phosphoric acid. It possesses the property of dissolving
small quantities of metallic oxides, with the formation of compounds
which, in the case of certain metals, have characteristic colors. It is
therefore used in the detection of these metals.
While the secondary phosphates, on heating, give salts of pyrophosphoric
acid, the primary phosphates yield salts of metaphosphoric acid. The
equations representing these reactions are as follows:
2Na_{2}HPO_{4} = Na_{4}P_{3}O_{7} + H_{2}O,
NaH_{2}PO_{4} = NaPO_{3} + H_{2}O.
~Fertilizers.~ When crops are produced year after year on the same field
certain constituents of the soil essential to plant growth are removed,
and the soil becomes impoverished and unproductive. To make the land
once more fertile these constituents must be replaced. The calcium
phosphate of the mineral deposits or of bone ash serves well as a
material for restoring phosphorus to soils exhausted of that essential
element; but a more soluble substance, which the plants can more readily
assimilate, is desirable. It is better, therefore, to convert the
insoluble calcium phosphate into the soluble primary phosphate before it
is applied as fertilizer. It will be seen by reference to the formulas
for the orthophosphates (see page 244) that in a primary phosphate only
one hydrogen atom of phosphoric acid is replaced by a metal. Since the
calcium atom always replaces two hydrogen atoms, it might be thought
that there could be no primary calcium phosphate; but if the calcium
atom replaces one hydrogen atom from each of two molecules of phosphoric
acid, the salt Ca(H_{2}PO_{4})_{2} will result, and this is a primary
phosphate. It can be made by treatment of the normal phosphate with the
necessary amount of sulphuric acid, calcium sulphate being formed at the
same time, thus:
Ca_{3}(PO_{4})_{2} + 2H_{2}SO_{4} = Ca(H_{2}PO_{4})_{2} + 2CaSO_{4}.
The resulting mixture is a powder, which is sold as a fertilizer under
the name of "superphosphate of lime."
ARSENIC
~Occurrence.~ Arsenic occurs in considerable quantities in nature as the
native element, as the sulphides realgar (As_{2}S_{2}) and orpiment
(As_{2}S_{3}), as oxide (As_{2}O_{3}), and as a constituent of many
metallic sulphides, such as arsenopyrite (FeAsS).
~Preparation.~ The element is prepared by purifying the native arsenic, or
by heating the arsenopyrite in iron tubes, out of contact with air,
when the reaction expressed by the following equation occurs:
FeAsS = FeS + As.
The arsenic, being volatile, condenses in chambers connected with the
heated tubes. It is also made from the oxide by reduction with carbon:
2As_{2}O_{3} + 3C = 4As + 3CO_{2}.
~Properties.~ Arsenic is a steel-gray, metallic-looking substance of
density 5.73. Though resembling metals in appearance, it is quite
brittle, being easily powdered in a mortar. When strongly heated it
sublimes, that is, it passes into a vapor without melting, and condenses
again to a crystalline solid when the vapor is cooled. Like phosphorus
it can be obtained in several allotropic forms. It alloys readily with
some of the metals, and finds its chief use as an alloy with lead, which
is used for making shot, the alloy being harder than pure lead. When
heated on charcoal with the blowpipe it is converted into an oxide which
volatilizes, leaving the charcoal unstained by any oxide coating. It
burns readily in chlorine gas, forming arsenic trichloride,--
As + 3Cl = AsCl_{3}.
Unlike most of its compounds, the element itself is not poisonous.
~Arsine~ (AsH_{3}). When any compound containing arsenic is brought into
the presence of nascent hydrogen, arsine (AsH_{3}), corresponding to
phosphine and ammonia, is formed. The reaction when oxide of arsenic is
so treated is
As_{2}O_{3} + 12H = 2AsH_{3} + 3H_{2}O.
Arsine is a gas with a peculiar garlic-like odor, and is intensely
poisonous. A single bubble of pure gas has been known to prove fatal. It
is an unstable compound, decomposing into its elements when heated to a
moderate temperature. It is combustible, burning with a pale
bluish-white flame to form arsenic trioxide and water when air is in
excess:
2AsH_{3} + 6O = As_{2}O_{3} + 3H_{2}O.
When the supply of air is deficient water and metallic arsenic are
formed:
2AsH_{3} + 3O = 3H_{2}O + 2As.
These reactions make the detection of even minute quantities of arsenic
a very easy problem.
[Illustration: Fig. 72]
~Marsh's test for arsenic.~ The method devised by Marsh for
detecting arsenic is most frequently used, the apparatus being
shown in Fig. 72. Hydrogen is generated in the flask A by the
action of dilute sulphuric acid on zinc, is dried by passing
over calcium chloride in the tube B, and after passing
through the hard-glass tube C is ignited at the jet D. If a
substance containing arsenic is now introduced into the
generator A, the arsenic is converted into arsine by the
action of the nascent hydrogen, and passes to the jet along
with the hydrogen. If the tube C is strongly heated at some
point near the middle, the arsine is decomposed while passing
this point and the arsenic is deposited just beyond the heated
point in the form of a shining, brownish-black mirror. If the
tube is not heated, the arsine burns along with the hydrogen at
the jet. Under these conditions a small porcelain dish crowded
down into the flame is blackened by a spot of metallic arsenic,
for the arsine is decomposed by the heat of the flame, and the
arsenic, cooled below its kindling temperature by the cold
porcelain, deposits upon it as a black spot. Antimony conducts
itself in the same way as arsenic, but the antimony deposit is
more sooty in appearance. The two can also be distinguished by
the fact that sodium hypochlorite (NaClO) dissolves the arsenic
deposit, but not that formed by antimony.
~Oxides of arsenic.~ Arsenic forms two oxides, As_{2}O_{3} and
As_{2}O_{5}, corresponding to those of phosphorus. Of these arsenious
oxide, or arsenic trioxide (As_{2}O_{3}), is much better known, and is
the substance usually called white arsenic, or merely arsenic. It is
found as a mineral, but is usually obtained as a by-product in burning
pyrite in the sulphuric-acid industry. The pyrite has a small amount of
arsenopyrite in it, and when this is burned arsenious oxide is formed as
a vapor together with sulphur dioxide:
2FeAsS + 10O = Fe_{2}O_{3} + As_{2}O_{3} + 2SO_{2}.
The arsenious oxide is condensed in appropriate chambers. It is a rather
heavy substance, obtained either as a crystalline powder or as large,
vitreous lumps, resembling lumps of porcelain in appearance. It is very
poisonous, from 0.2 to 0.3 g. being a fatal dose. It is frequently given
as a poison, since it is nearly tasteless and does not act very rapidly.
This slow action is due to the fact that it is not very soluble, and
hence is absorbed slowly by the system. Arsenious oxide is also used as
a chemical reagent in glass making and in the dye industry.
~Acids of arsenic.~ Like the corresponding oxides of phosphorus, the
oxides of arsenic are acid anhydrides. In solution they combine with
bases to form salts, corresponding to the salts of the acids of
phosphorus. Thus we have salts of the following acids:
H_{3}AsO_{3} arsenious acid.
H_{3}AsO_{4} orthoarsenic acid.
H_{4}As_{2}O_{3} pyroarsenic acid.
HAsO_{3} metarsenic acid.
Several other acids of arsenic are also known. Not all of these can be
obtained as free acids, since they tend to lose water and form the
oxides. Thus, instead of obtaining arsenious acid (H_{3}AsO_{3}), the
oxide As_{2}O_{3} is obtained:
2H_{3}AsO_{3} = As_{2}O_{3} + 3H_{2}O.
Salts of all the acids are known, however, and some of them have
commercial value. Most of them are insoluble, and some of the copper
salts, which are green, are used as pigments. Paris green, which has a
complicated formula, is a well-known insecticide.
~Antidote for arsenical poisoning.~ The most efficient antidote for
arsenic poisoning is ferric hydroxide. It is prepared as needed,
according to the equation
Fe_{2}(SO_{4})_{3} + 3Mg(OH)_{2} = 2Fe(OH)_{3} + 3MgSO_{4}.
~Sulphides of arsenic.~ When hydrogen sulphide is passed into an acidified
solution containing an arsenic compound the arsenic is precipitated as a
bright yellow sulphide, thus:
2H_{3}AsO_{3} + 3H_{2}S = As_{2}S_{3} + 6H_{2}O,
2H_{3}AsO_{4} + 5H_{2}S = As_{2}S_{5} + 8H_{2}O.
In this respect arsenic resembles the metallic elements, many of which
produce sulphides under similar conditions. The sulphides of arsenic,
both those produced artificially and those found in nature, are used as
yellow pigments.
ANTIMONY
~Occurrence.~ Antimony occurs in nature chiefly as the sulphide
(Sb_{2}S_{3}), called stibnite, though it is also found as oxide and as
a constituent of many complex minerals.
~Preparation.~ Antimony is prepared from the sulphide in a very simple
manner. The sulphide is melted with scrap iron in a furnace, when the
iron combines with the sulphur to form a slag, or liquid layer of melted
iron sulphide, while the heavier liquid, antimony, settles to the bottom
and is drawn off from time to time. The reaction involved is represented
by the equation
Sb_{2}S_{3} + 3Fe = 2Sb + 3FeS.
~Physical properties.~ Antimony is a bluish-white, metallic-looking
substance whose density is 6.7. It is highly crystalline, hard, and very
brittle. It has a rather low melting point (432°) and expands very
noticeably on solidifying.
~Chemical properties.~ In chemical properties antimony resembles arsenic
in many particulars. It forms the oxides Sb_{2}O_{3} and Sb_{2}O_{5},
and in addition Sb_{2}O_{4}. It combines with the halogen elements with
great energy, burning brilliantly in chlorine to form antimony
trichloride (SbCl_{3}). When heated on charcoal with the blowpipe it is
oxidized and forms a coating of antimony oxide on the charcoal which has
a characteristic bluish-white color.
~Stibine~ (SbH_{3}). The gas stibine (SbH_{3}) is formed under conditions
which are very similar to those which produce arsine, and it closely
resembles the latter compound, though it is still less stable. It is
very poisonous.
~Acids of antimony.~ The oxides Sb_{2}O_{3} and Sb_{2}O_{5} are
weak acid anhydrides and are capable of forming two series of
acids corresponding in formulas to the acids of phosphorus and
arsenic. They are much weaker, however, and are of little
practical importance.
~Sulphides of antimony.~ Antimony resembles arsenic in that
hydrogen sulphide precipitates it as a sulphide when conducted
into an acidified solution containing an antimony compound:
2SbCl_{3} + 3H_{2}S = Sb_{2}S_{3} + 6HCl,
2SbCl_{5} + 5H_{2}S = Sb_{2}S_{5} + 10HCl.
The two sulphides of antimony are called the trisulphide and
the pentasulphide respectively. When prepared in this way they
are orange-colored substances, though the mineral stibnite is
black.
~Metallic properties of antimony.~ The physical properties of the element
are those of a metal, and the fact that its sulphide is precipitated by
hydrogen sulphide shows that it acts like a metal in a chemical way.
Many other reactions show that antimony has more of the properties of a
metal than of a non-metal. The compound Sb(OH)_{3}, corresponding to
arsenious acid, while able to act as a weak acid is also able to act as
a weak base with strong acids. For example, when treated with
concentrated hydrochloric acid antimony chloride is formed:
Sb(OH)_{3} + 3HCl = SbCl_{3} + 3H_{2}O.
A number of elements act in this same way, their hydroxides under some
conditions being weak acids and under others weak bases.
ALLOYS
Some metals when melted together thoroughly intermix, and on cooling
form a homogeneous, metallic-appearing substance called an _alloy_. Not
all metals will mix in this way, and in some cases definite chemical
compounds are formed and separate out as the mixture solidifies, thus
destroying the uniform quality of the alloy. In general the melting
point of the alloy is below the average of the melting points of its
constituents, and it is often lower than any one of them.
Antimony forms alloys with many of the metals, and its chief commercial
use is for such purposes. It imparts to its alloys high density, rather
low melting point, and the property of expanding on solidification.
Such an alloy is especially useful in type founding, where fine lines
are to be reproduced on a cast. Type metal consists of antimony, lead,
and tin. Babbitt metal, used for journal bearings in machinery, contains
the same metals in a different proportion together with a small
percentage of copper.
BISMUTH
~Occurrence.~ Bismuth is usually found in the uncombined form in nature.
It also occurs as oxide and sulphide. Most of the bismuth of commerce
comes from Saxony, and from Mexico and Colorado, but it is not an
abundant element.
~Preparation.~ It is prepared by merely heating the ore containing the
native bismuth and allowing the melted metal to run out into suitable
vessels. Other ores are converted into oxides and reduced by heating
with carbon.
~Physical properties.~ Bismuth is a heavy, crystalline, brittle metal
nearly the color of silver, but with a slightly rosy tint which
distinguishes it from other metals. It melts at a low temperature (270°)
and has a density of 9.8. It is not acted upon by the air at ordinary
temperatures.
~Chemical properties.~ When heated with the blowpipe on charcoal, bismuth
gives a coating of the oxide Bi_{2}O_{3}. This has a yellowish-brown
color which easily distinguishes it from the oxides formed by other
metals. It combines very readily with the halogen elements, powdered
bismuth burning readily in chlorine. It is not very easily acted upon by
hydrochloric acid, but nitric and sulphuric acids act upon it in the
same way that they do upon copper.
~Uses.~ Bismuth finds its chief use as a constituent of alloys,
particularly in those of low melting point. Some of these melt in hot
water. For example, Wood's metal, consisting of bismuth, lead, tin, and
cadmium, melts at 60.5°.
~Compounds of bismuth.~ Unlike the other elements of this group, bismuth
has almost no acid properties. Its chief oxide, Bi_{2}O_{3}, is basic in
its properties. It dissolves in strong acids and forms salts of bismuth:
Bi_{2}O_{3} + 6HCl = 2BiCl_{3} + 3H_{2}O,
Bi_{2}O_{3} + 6HNO_{3} = 2Bi(NO_{3})_{3} + 3H_{2}O.
The nitrate and chloride of bismuth can be obtained as well-formed
colorless crystals. When treated with water the salts are decomposed in
the manner explained in the following paragraph.
HYDROLYSIS
Many salts such as those of antimony and bismuth form solutions which
are somewhat acid in reaction, and must therefore contain hydrogen ions.
This is accounted for by the same principle suggested to explain the
fact that solutions of potassium cyanide are alkaline in reaction (p.
210). Water forms an appreciable number of hydrogen and hydroxyl ions,
and very weak bases such as bismuth hydroxide are dissociated to but a
very slight extent. When Bi^{+++} ions from bismuth chloride, which
dissociates very readily, are brought in contact with the OH^{-} ions
from water, the two come to the equilibrium expressed in the equation
Bi^{+++} + 3OH^{-} <--> Bi(OH)_{3}.
For every hydroxyl ion removed from the solution in this way a hydrogen
ion is left free, and the solution becomes acid in reaction.
Reactions of this kind and that described under potassium cyanide are
called _hydrolysis_.
DEFINITION: _Hydrolysis is the action of water upon a salt to form an
acid and a base, one of which is very slightly dissociated._
~Conditions favoring hydrolysis.~ While hydrolysis is primarily due to the
slight extent to which either the acid or the base formed is
dissociated, several other factors have an influence upon the extent to
which it will take place.
1. _Influence of mass._ Since hydrolysis is a reversible reaction, the
relative masses of the reacting substances influence the point at which
equilibrium will be reached. In the equilibrium
BiCl_{3} + 3H_{2}O <--> Bi(OH)_{3} + 3HCl
the addition of more water will result in the formation of more bismuth
hydroxide and hydrochloric acid. The addition of more hydrochloric acid
will convert some of the bismuth hydroxide into bismuth chloride.
2. _Formation of insoluble substances._ When one of the products of
hydrolysis is nearly insoluble in water the solution will become
saturated with it as soon as a very little has been formed. All in
excess of this will precipitate, and the reaction will go on until the
acid set free increases sufficiently to bring about an equilibrium. Thus
a considerable amount of bismuth and antimony hydroxides are
precipitated when water is added to the chlorides of these elements. The
greater the dilution the more hydroxide precipitates. The addition of
hydrochloric acid in considerable quantity will, however, redissolve the
precipitate.
~Partial hydrolysis.~ In many cases the hydrolysis of a salt is only
partial, resulting in the formation of basic salts instead of the free
base. Most of these basic salts are insoluble in water, which accounts
for their ready formation. Thus bismuth chloride may hydrolyze by
successive steps, as shown in the equations
BiCl_{3} + H_{2}O = Bi(OH)Cl_{2} + HCl,
BiCl_{3} + 2H_{2}O = Bi(OH)_{2}Cl + 2HCl,
BiCl_{3} + 3H_{2}O = Bi(OH)_{3} + 3HCl.
The basic salt so formed may also lose water, as shown in the equation
Bi(OH)_{2}Cl = BiOCl + H_{2}O.
The salt represented in the last equation is sometimes called bismuth
oxychloride, or bismuthyl chloride. The corresponding nitrate,
BiONO_{3}, is largely used in medicine under the name of subnitrate of
bismuth. In these two compounds the group of atoms, BiO, acts as a
univalent metallic radical and is called _bismuthyl_. Similar basic
salts are formed by the hydrolysis of antimony salts.
EXERCISES
1. Name all the elements so far studied which possess allotropic forms.
2. What compounds would you expect phosphorus to form with bromine and
iodine? Write the equations showing the action of water on these
compounds.
3. In the preparation of phosphine, why is coal gas passed into the
flask? What other gases would serve the same purpose?
4. Give the formula for the salt which phosphine forms with hydriodic
acid. Give the name of the compound.
5. Could phosphoric acid be substituted for sulphuric acid in the
preparation of the common acids?
6. Write the equations for the preparation of the three sodium salts of
orthophosphoric acid.
7. Why does a solution of disodium hydrogen phosphate react alkaline?
8. On the supposition that bone ash is pure calcium phosphate, what
weight of it would be required in the preparation of 1 kg. of
phosphorus?
9. If arsenopyrite is heated in a current of air, what products are
formed?
10. (a) Write equations for the complete combustion of hydrosulphuric
acid, methane, and arsine. (b) In what respects are the reactions
similar?
11. Write the equations for all the reactions involved in Marsh's test
for arsenic.
12. Write the names and formulas for the acids of antimony.
13. Write the equations showing the hydrolysis of antimony trichloride;
of bismuth nitrate.
14. In what respects does nitrogen resemble the members of the
phosphorus family?