An Elementary Study of Chemistry
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William McPherson >> An Elementary Study of Chemistry
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EXERCISES
1. Account for the fact that copper has been used for so long a time.
2. Write equations for the action of concentrated sulphuric and nitric
acids upon the metals of this family.
3. How would you account for the fact that normal copper sulphate is
slightly acid to litmus?
4. Contrast the action of heat on cupric nitrate and mercuric nitrate.
5. State reasons why mercury is adapted for use in thermometers and
barometers.
6. How could you distinguish between mercurous chloride and mercuric
chloride?
7. Write equations for the preparation of mercuric and mercurous
iodides.
8. How would you account for the fact that solutions of the different
salts of a metal usually have the same color?
9. Crude silver usually contains iron and lead. What would become of
these metals in refining by parting with sulphuric acid?
10. In the amalgamation process for extracting silver, how does ferrous
chloride convert silver chloride into silver? Write equation. Why is the
silver sulphide first changed into silver chloride?
11. What impurities would you expect to find in the copper sulphate
prepared from the refining of silver?
12. How could you prepare pure silver chloride from a silver coin?
13. Mercuric nitrate and silver nitrate are both white solids soluble in
water. How could you distinguish between them?
14. Account for the fact that sulphur waters turn a silver coin black;
also for the fact that a silver spoon is blackened by foods (eggs, for
example) containing sulphur.
15. When a solution of silver nitrate is added to a solution of
potassium chlorate no precipitate forms. How do you account for the fact
that a precipitate of silver chloride is not formed?
CHAPTER XXIX
TIN AND LEAD
====================================================================
| | | | |
| SYMBOL | ATOMIC | DENSITY | MELTING | COMMON OXIDES
| | WEIGHT | | POINT |
_____|________|________|_________|_________|________________________
| | | | |
Tin | Sn | 119.0 | 7.35 | 235 deg. | SnO SnO_{2}
Lead | Pb | 206.9 | 11.38 | 327 deg. | PbO Pb_{3}O_{4} PbO_{2}
====================================================================
~The family.~ Tin and lead, together with silicon and germanium, form a
family in Group IV of the periodic table. Silicon has been discussed
along with the non-metals, while germanium, on account of its rarity,
needs only to be mentioned.
The other family of Group IV includes carbon, already described, and a
number of rare elements.
TIN
~Occurrence.~ Tin is found in nature chiefly as the oxide (SnO_{2}),
called cassiterite or tinstone. The most famous mines are those of
Cornwall in England, and of the Malay Peninsula and East India Islands;
in small amounts tinstone is found in many other localities.
~Metallurgy.~ The metallurgy of tin is very simple. The ore, separated as
far as possible from earthy materials, is mixed with carbon and heated
in a furnace, the reduction taking place readily. The equation is
SnO_{2} + C = Sn + CO_{2}.
The metal is often purified by carefully heating it until it is partly
melted; the pure tin melts first and can be drained away from the
impurities.
~Properties.~ Pure tin, called block tin, is a soft white metal with a
silver-like appearance and luster; it melts readily (235 deg.) and is
somewhat lighter than copper, having a density of 7.3. It is quite
malleable and can be rolled out into very thin sheets, forming tin foil;
most tin foil, however, contains a good deal of lead.
Under ordinary conditions it is quite unchanged by air or moisture, but
at a high temperature it burns in air, forming the oxide SnO_{2}. Dilute
acids have no effect upon it, but concentrated acids attack it readily.
Concentrated hydrochloric acid changes it into the chloride
Sn + 2HCl = SnCl_{2} + 2H.
With sulphuric acid tin sulphate and sulphur dioxide are formed:
Sn + 2H_{2}SO_{4} = SnSO_{4} + SO_{2} + 2H_{2}O
Concentrated nitric acid oxidizes it, forming a white insoluble compound
of the formula H_{2}SnO_{3}, called metastannic acid:
3Sn + 4HNO_{3} + H_{2}O = 3H_{2}SnO_{3} + 4NO.
~Uses of tin.~ A great deal of tin is made into tin plate by dipping thin
steel sheets into the melted metal. Owing to the way in which tin
resists the action of air and dilute acids, tin plate is used in many
ways, such as in roofing, and in the manufacture of tin cans, cooking
vessels, and similar articles.
Many useful alloys contain tin, some of which have been mentioned in
connection with copper. When tin is alloyed with other metals of low
melting point, soft, easily melted alloys are formed which are used for
friction bearings in machinery; tin, antimony, lead, and bismuth are the
chief constituents of these alloys. Pewter and soft solder are alloys of
tin and lead.
~Compounds of tin.~ Tin forms two series of compounds: the stannous, in
which the tin is divalent, illustrated in the compounds SnO, SnS,
SnCl_{2}; the stannic, in which it is tetravalent as shown in the
compounds SnO_{2}, SnS_{2}. There is also an acid, H_{2}SnO_{3}, called
stannic acid, which forms a series of salts called stannates. While this
acid has the same composition as metastannic acid, the two are quite
different in their chemical properties. This difference is probably due
to the different arrangement of the atoms in the molecules of the two
substances. Only a few compounds of tin need be mentioned.
~Stannic oxide~ (SnO_{2}). Stannic oxide is of interest, since it is the
chief compound of tin found in nature. It is sometimes found in
good-sized crystals, but as prepared in the laboratory is a white
powder. When fused with potassium hydroxide it forms potassium stannate,
acting very much like silicon dioxide:
SnO_{2} + 2KOH = K_{2}SnO_{3} + H_{2}O.
~Chlorides of tin.~ Stannous chloride is prepared by dissolving tin in
concentrated hydrochloric acid and evaporating the solution to
crystallization. The crystals which are obtained have the composition
SnCl_{2}.2H_{2}O, and are known as tin crystals. By treating a solution
of stannous chloride with aqua regia, stannic chloride is formed:
SnCl_{2} + 2Cl = SnCl_{4}.
The salt which crystallizes from such a solution has the composition
SnCl_{4}.5H_{2}O, and is known commercially as oxymuriate of tin. If
metallic tin is heated in a current of dry chlorine, the anhydrous
chloride (SnCl_{4}) is obtained as a heavy colorless liquid which fumes
strongly on exposure to air.
The ease with which stannous chloride takes up chlorine to form stannic
chloride makes it a good reducing agent in many reactions, changing the
higher chlorides of metals to lower ones. Thus mercuric chloride is
changed into mercurous chloride:
SnCl_{2} + 2HgCl_{2} = SnCl_{4} + 2HgCl.
If the stannous chloride is in excess, the reaction may go further,
producing metallic mercury:
SnCl_{2} + 2HgCl = SnCl_{4} + 2Hg.
Ferric chloride is in like manner reduced to ferrous chloride:
SnCl_{3} + 2FeCl_{3} = SnCl_{4} + 2FeCl_{2}.
The chlorides of tin, as well as the alkali stannates, are much used as
mordants in dyeing processes. The hydroxides of tin and free stannic
acid, which are easily liberated from these compounds, possess in very
marked degree the power of fixing dyes upon fibers, as explained under
aluminium.
LEAD
~Occurrence.~ Lead is found in nature chiefly as the sulphide (PbS),
called galena; to a much smaller extent it occurs as carbonate,
sulphate, chromate, and in a few other forms. Practically all the lead
of commerce is made from galena, two general methods of metallurgy being
in use.
~Metallurgy.~ 1. The sulphide is melted with scrap iron, when iron
sulphide and metallic lead are formed; the liquid lead, being the
heavier, sinks to the bottom of the vessel and can be drawn off:
PbS + Fe = Pb + FeS.
2. The sulphide is roasted in the air until a part of it has been
changed into oxide and sulphate. The air is then shut off and the
heating continued, the reactions indicated in the following equations
taking place:
2PbO + PbS = 3Pb + SO_{2},
PbSO_{4} + PbS = 2Pb + 2SO_{2}.
The lead so prepared usually contains small amounts of silver, arsenic,
antimony, copper, and other metals. The silver is removed by Parkes's
method, as described under silver, and the other metals in various ways.
The lead of commerce is one of the purest commercial metals, containing
as a rule only a few tenths per cent of impurities.
~Properties.~ Lead is a heavy metal (den. = 11.33) which has a brilliant
silvery luster on a freshly cut surface, but which soon tarnishes to a
dull blue-gray color. It is soft, easily fused (melting at 327 deg.), and
quite malleable, but has little toughness or strength.
It is not acted upon to any great extent by the oxygen of the air under
ordinary conditions, but is changed into oxide at a high temperature.
With the exception of hydrochloric and sulphuric acids, most acids, even
very weak ones, act upon it, forming soluble lead salts. Hot,
concentrated hydrochloric and sulphuric acids also attack it to a slight
extent.
~Uses.~ Lead is employed in the manufacture of lead pipes and in large
storage batteries. In the form of sheet lead it is used in lining the
chambers of sulphuric acid works and in the preparation of paint
pigments. Some alloys of lead, such as solder and pewter (lead and tin),
shot (lead and arsenic), and soft bearing metals, are widely used. Type
metal consists of lead, antimony, and sometimes tin. Compounds of lead
form several important pigments.
~Compounds of lead.~ In nearly all its compounds lead has a valence of 2,
but a few corresponding to stannic compounds have a valence of 4.
~Lead oxides.~ Lead forms a number of oxides, the most important of which
are litharge, red lead or minium, and lead peroxide.
1. _Litharge_ (PbO). This oxide forms when lead is oxidized at a rather
low temperature, and is obtained as a by-product in silver refining. It
is a pale yellow powder, and has a number of commercial uses. It is
easily soluble in nitric acid:
PbO + 2HNO_{3} = Pb(NO_{3})_{2} + H_{2}O.
2. _Red lead, or minium_ (Pb_{3}O_{4}). Minium is prepared by heating
lead (or litharge) to a high temperature in the air. It is a heavy
powder of a beautiful red color, and is much used as a pigment.
3. _Lead peroxide_ (PbO_{2}). This is left as a residue when minium is
heated with nitric acid:
Pb_{3}O_{4} + 4HNO_{3} = 2Pb(NO_{3})_{2} + PbO_{2} + 2H_{2}O.
It is a brown powder which easily gives up a part of its oxygen and,
like manganese dioxide and barium dioxide, is a good oxidizing agent.
~Soluble salts of lead.~ The soluble salts of lead can be made by dissolving
(Pb(C_{2}H_{3}O_{2})_{2}.3H_{2}O), litharge in acids. Lead acetate
called sugar of lead, and lead nitrate (Pb(NO_{3})_{2}) are the most
familiar examples. They are while crystalline solids and are poisonous
in character.
~Insoluble salts of lead; lead carbonate.~ While the normal carbonate of
lead (PbCO_{3}) is found to some extent, in nature and can be prepared
in the laboratory, basic carbonates of varying composition are much more
easy to obtain. One of the simplest of these has the composition
2PbCO_{3}.Pb(OH)_{2}. A mixture of such carbonates is called white lead.
This is prepared on a large scale as a paint pigment and as a body for
paints which are to be colored with other substances.
~White lead.~ White lead is an amorphous white substance which,
when mixed with oil, has great covering power, that is, it
spreads out in an even waxy film, free from streaks and lumps,
and covers the entire surface upon which it is spread. Its
disadvantage as a pigment lies in the fact that it gradually
blackens when exposed to sulphur compounds, which are often
present in the air, forming black lead sulphide (PbS).
~Technical preparation of white lead.~ Different methods are used
in the preparation of white lead, but the old one known as the
Dutch process is still the principal one employed. In this
process, earthenware pots about ten inches high and of the
shape shown in Fig. 89 are used. In the bottom A is placed a
3% solution of acetic acid (vinegar answers the purpose very
well). The space above this is filled with thin, perforated,
circular pieces of lead, supported by the flange B of the
pot. These pots are placed close together on a bed of tan bark
on the floor of a room known as the corroding room. They are
covered over with boards, upon which tan bark is placed, and
another row of pots is placed on this. In this way the room is
filled. The white lead is formed by the fumes of the acetic
acid, together with the carbon dioxide set free in the
fermentation of the tan bark acting on the lead. About three
months are required to complete the process.
[Illustration 1: Fig. 89]
~Lead sulphide~ (PbS). In nature this compound occurs in highly
crystalline condition, the crystals having much the same luster as pure
lead. It is readily prepared in the laboratory as a black precipitate,
by the action of hydrosulphuric acid upon soluble lead salts:
Pb(NO_{3})_{2} + H_{2}S = PbS + 2HNO_{3}.
It is insoluble both in water and in dilute acids.
~Other insoluble salts.~ Lead chromate (PbCrO_{4}) is a yellow substance
produced by the action of a soluble lead salt upon a soluble chromate,
thus:
K_{2}CrO_{4} + Pb(NO_{3})_{2} = PbCrO_{4} + 2 KNO_{3}.
It is used as a yellow pigment. Lead sulphate (PbSO_{4}) is a white
substance sometimes found in nature and easily prepared by
precipitation. Lead chloride (PbCl_{2}) is likewise a white substance
nearly insoluble in cold water, but readily soluble in boiling water.
~Thorium and cerium.~ These elements are found in a few rare
minerals, especially in the monazite sand of the Carolinas and
Brazil. The oxides of these elements are used in the
preparation of the Welsbach mantles for gas lights, because of
the intense light given out when a mixture of the oxides is
heated. These mantles contain the oxides of cerium and thorium
in the ratio of about 1% of the former to 99% of the latter.
Compounds of thorium, like those of radium, are found to
possess radio-activity, but in a less degree.
EXERCISES
1. How could you detect lead if present in tin foil?
2. Stannous chloride reduces gold chloride (AuCl_{3}) to gold. Give
equation.
3. What are the products of hydrolysis when stannic chloride is used as
a mordant?
4. How could you detect arsenic, antimony, or copper in lead?
5. Why is lead so extensively used for making water pipes?
6. What sulphates other than lead are insoluble?
7. Could lead nitrate be used in place of barium chloride in testing for
sulphates?
8. How much lead peroxide could be obtained from 1 kg. of minium?
9. The purity of white lead is usually determined by observing the
volume of carbon dioxide given off when it is treated with an acid. What
acid should be used? On the supposition that it has the formula
2PbCO_{3}.Pb(OH)_{2}, how nearly pure was a sample if 1 g. gave 30 cc.
of carbon dioxide at 20 deg. and 750 mm.?
10. Silicon belongs in the same family with tin and lead. In what
respects are these elements similar?
11. What weight of tin could be obtained by the reduction of 1 ton of
cassiterite?
12. What reaction would you expect to take place when lead peroxide is
treated with hydrochloric acid?
13. White lead is often adulterated with barytes. Suggest a method for
detecting it, if present, in a given example of white lead.
CHAPTER XXX
MANGANESE AND CHROMIUM
====================================================================
| | | | |
| SYMBOL | ATOMIC | DENSITY | MELTING | FORMULAS OF ACIDS
| | WEIGHT | | POINT |
__________|________|________|_________|_________|___________________
| | | | |
Manganese | Mn | 55.0 | 8.01 | 1900 deg. | H_{2}MnO_{4} and
| | | | | HMnO_{4}
Chromium | Cr | 52.1 | 7.3 | 3000 deg. | H_{2}CrO_{4} and
| | | | | H_{2}Cr_{2}O_{7}
====================================================================
~General.~ Manganese and chromium, while belonging to different families,
have so many features in common in their chemical conduct that they may
be studied together with advantage. They differ from most of the
elements so far studied in that they can act either as acid-forming or
base-forming elements. As base-forming elements each of the metals forms
two series of salts. In the one series, designated by the suffix "ous,"
the metal is divalent; in the other series, designated by the suffix
"ic," the metal is trivalent. Only the manganous and the chromic salts,
however, are of importance. The acids in which these elements play the
part of a non-metal are unstable, but their salts are usually stable,
and some of them are important compounds.
MANGANESE
~Occurrence.~ Manganese is found in nature chiefly as the dioxide MnO_{2},
called pyrolusite. In smaller amounts it occurs as the oxides
Mn_{2}O_{3} and Mn_{3}O_{4}, and as the carbonate MnCO_{3}. Some iron
ores also contain manganese.
~Preparation and properties.~ The element is difficult to prepare in pure
condition and has no commercial applications. It can be prepared,
however, by reducing the oxide with aluminium powder or by the use of
the electric furnace, with carbon as the reducing agent. The metal
somewhat resembles iron in appearance, but is harder, less fusible, and
more readily acted upon by air and moisture. Acids readily dissolve it,
forming manganous salts.
~Oxides of manganese.~ The following oxides of manganese are known: MnO,
Mn_{2}O_{3}, Mn_{3}O_{4}, MnO_{2}, and Mn_{2}O_{7}. Only one of these,
the dioxide, needs special mention.
~Manganese dioxide~ (_pyrolusite_) (MnO_{2}). This substance is the most
abundant manganese compound found in nature, and is the ore from which
all other compounds of manganese are made. It is a hard, brittle, black
substance which is valuable as an oxidizing agent. It will be recalled
that it is used in the preparation of chlorine and oxygen, in
decolorizing glass which contains iron, and in the manufacture of
ferromanganese.
~Compounds containing manganese as a base-forming element.~ As has been
stated previously, manganese forms two series of salts. The most
important of these salts, all of which belong to the manganous series,
are the following:
Manganous chloride MnCl_{2}.4H_{2}O.
Manganous sulphide MnS.
Manganous sulphate MnSO_{4}.4H_{2}O.
Manganous carbonate MnCO_{3}.
Manganous hydroxide Mn(OH)_{2}.
The chloride and sulphate may be prepared by heating the dioxide with
hydrochloric and sulphuric acids respectively:
MnO_{2} + 4HCl = MnCl_{2} + 2H_{2}O + 2Cl,
MnO_{2} + H_{2}SO_{4} = MnSO_{4} + H_{2}O + O.
The sulphide, carbonate, and hydroxide, being insoluble, may be prepared
from a solution of the chloride or sulphate by precipitation with the
appropriate reagents. Most of the manganous salts are rose colored. They
not only have formulas similar to the ferrous salts, but resemble them
in many of their chemical properties.
~Compounds containing manganese as an acid-forming element.~ Manganese
forms two unstable acids, namely, manganic acid and permanganic acid.
While these acids are of little interest, some of their salts,
especially the permanganates, are important compounds.
~Manganic acid and manganates.~ When manganese dioxide is fused with an
alkali and an oxidizing agent a green compound is formed. The equation,
when caustic potash is used, is as follows:
MnO_{2} + 2KOH + O = K_{2}MnO_{4} + H_{2}O.
The green compound (K_{2}MnO_{4}) is called potassium manganate, and is
a salt of the unstable manganic acid (H_{2}MnO_{4}). The manganates are
all very unstable.
~Permanganic acid and the permanganates.~ When carbon dioxide is passed
through a solution of a manganate a part of the manganese is changed
into manganese dioxide, while the remainder forms a salt of the unstable
acid HMnO_{4}, called permanganic acid. The equation is
3K_{2}MnO_{4} + 2CO_{2} = MnO_{2} + 2KMnO_{4} + 2K_{2}CO_{3}.
Potassium permanganate (KMnO_{4}) crystallizes in purple-black needles
and is very soluble in water, forming an intensely purple solution. All
other permanganates, as well as permanganic acid itself, give solutions
of the same color.
~Oxidizing properties of the permanganates.~ The permanganates are
remarkable for their strong oxidizing properties. When used as an
oxidizing agent the permanganate is itself reduced, the exact character
of the products formed from it depending upon whether the oxidation
takes place (1) in an alkaline or neutral solution, or (2) in an acid
solution.
1. _Oxidation in alkaline or neutral solution._ When the solution is
either alkaline or neutral the potassium and the manganese of the
permanganate are both converted into hydroxides, as shown in the
equation
2KMnO_{4} + 5H_{2}O = 2Mn(OH)_{4} + 2KOH + 3O.
2. _Oxidation in acid solution._ When free acid such as sulphuric is
present, the potassium and the manganese are both changed into salts of
the acid:
2KMnO_{4} + 3H_{2}SO_{4} = K_{2}SO_{4} + 2MnSO_{4} + 3H_{2}O + 5O.
Under ordinary conditions, however, neither one of these reactions takes
place except in the presence of a third substance which is capable of
oxidation. The oxygen is not given off in the free state, as the
equations show, but is used up in effecting oxidation.
Potassium permanganate is particularly valuable as an oxidizing agent
not only because it acts readily either in acid or in alkaline solution,
but also because the reaction takes place so easily that often it is not
even necessary to heat the solution to secure action. The substance
finds many uses in the laboratory, especially in analytical work. It is
also used as an antiseptic as well as a disinfectant.
CHROMIUM
~Occurrence.~ The ore from which all chromium compounds are made is
chromite, or chrome iron ore (FeCr_{2}O_{4}). This is found most
abundantly in New Caledonia and Turkey. The element also occurs in small
quantities in many other minerals, especially in crocoisite (PbCrO_{4}),
in which mineral it was first discovered.
~Preparation.~ Chromium, like manganese, is very hard to reduce from its
ores, owing to its great affinity for oxygen. It can, however, be made
by the same methods which have proved successful with manganese.
Considerable quantities of an alloy of chromium with iron, called
ferrochromium, are now produced for the steel industry.
~Properties.~ Chromium is a very hard metal of about the same density as
iron. It is one of the most infusible of the metals, requiring a
temperature little short of 3000 deg. for fusion. At ordinary temperatures
air has little action on it; at higher temperatures, however, it burns
brilliantly. Nitric acid has no action on it, but hydrochloric and
dilute sulphuric acids dissolve it, liberating hydrogen.
~Compounds containing chromium as a base-forming element.~ While chromium
forms two series of salts, chromous salts are difficult to prepare and
are of little importance. The most important of the chromic series are
the following:
Chromic hydroxide Cr(OH)_{3}.
Chromic chloride CrCl_{3}.6H_{2}O.
Chromic sulphate Cr_{2}(SO_{4})_{3}.
Chrome alums
~Chromic hydroxide~ (Cr(OH)_{3}). This substance, being insoluble, can be
obtained by precipitating a solution of the chloride or sulphate with a
soluble hydroxide. It is a greenish substance which, like aluminium
hydroxide, dissolves in alkalis, forming soluble salts.
~Dehydration of chromium hydroxide.~ When heated gently chromic
hydroxide loses a part of its oxygen and hydrogen, forming the
substance CrO.OH, which, like the corresponding aluminium
compound, has more pronounced acid properties than the
hydroxide. It forms a series of salts very similar to the
spinels; chromite is the ferrous salt of this acid, having the
formula Fe(CrO_{2})_{2}. When heated to a higher temperature
chromic hydroxide is completely dehydrated, forming the
trioxide Cr_{2}O_{3}. This resembles the corresponding oxides
of aluminium and iron in many respects. It is a bright green
powder, and when ignited strongly becomes almost insoluble in
acids, as is also the case with aluminium oxide.
~Chromic sulphate~ (Cr_{2}(SO_{4})_{3}). This compound is a violet-colored
solid which dissolves in water, forming a solution of the same color.
This solution, however, turns green on heating, owing to the formation
of basic salts. Chromic sulphate, like ferric and aluminium sulphates,
unites with the sulphates of the alkali metals to form alums, of which
the best known are potassium chrome alum (KCr(SO_{4})_{2}.12H_{2}O) and
ammonium chrome alum (NH_{4}Cr(SO_{4})_{2}.12H_{2}O).
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