An Introduction to Chemical Science
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Title: An Introduction to Chemical Science
Author: Rufus P. Williams
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Language: English
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An Introduction to Chemical Science
by R.P. Williams, A.M.,
CONTENTS
PREFACE, BY R.P. WILLIAMS
TABLE OF CONTENTS
AN INTRODUCTION TO CHEMICAL SCIENCE
APPENDIX
TEXTBOOK ADVERTISEMENTS THAT APPEARED IN THE ORIGINAL EDITION
INFO ABOUT THIS E-TEXT EDITION
PREFACE, BY R.P. WILLIAMS
The object held constantly in view in writing this book has been to
prepare a suitable text-book in Chemistry for the average High
School,--one that shall be simple, practical, experimental, and
inductive, rather than a cyclopaedia of chemical information.
For the accomplishment of this purpose the author has endeavored
to omit superfluous matter, and give only the most useful and
interesting experiments, facts and theories.
In calling attention, by questions, and otherwise, to the more
important phenomena to be observed and facts to be learned, the
best features of the inductive system have been utilized.
Especially is the writing of equations, which constitute the
multum in parvo of chemical knowledge, insisted upon. As soon as
the pupil has become imbued with the spirit and meaning of
chemical equations, he need have little fear of failing to
understand the rest. To this end Chapters IX., XI., and XVI.
should be studied with great care.
In the early stages of the work the equations may with advantage
be memorized, but this can soon be discontinued. Whenever symbols
are employed, pupils should be required to give the corresponding
chemical names, or, better, both names and symbols.
The classification of chemical substances into acids, bases and
salts, and the distinctions and analogies between each of these
classes, have been brought into especial prominence. The general
relationship between the three classes, and the general
principles prevailing in the preparation of each, must be fully
understood before aught but the merest smattering of chemical
science can be known.
Chapters XV.-XXI. should be mastered as a key to the subsequent
parts of the book.
The mathematical and theoretical parts of Chemistry it has been
thought best to intersperse throughout the book, placing each
where it seemed to be especially needed; in this way, it is hoped
that the tedium which pupils find in studying consecutively many
chapters of theories will be avoided, and that the arrangement
will give an occasional change from the discussion of facts and
experiments to that of principles. In these chapters additional
questions should be given, and the pupil should be particularly
encouraged to make new problems of his own, and to solve theta.
It is needless to say that this treatise is primarily designed to
be used in connection with a laboratory. Like all other text-
books on the subject, it can be studied without such an
accessory; but the author attaches very little value to the study
of Chemistry without experimental work. The required apparatus
and chemicals involve but little expense, and the directions for
experimentation are the result of several years' experience with
classes as large as are to be found in the laboratory of any
school or college in the country.
During the present year the author personally supervises the work
of more than 180 different pupils in chemistry. This enables him
not only to assure himself that the experiments of the book are
practical, but that the directions for performing them are ample.
It is found advisable to perform most of the experiments, with
full explanation, in presence of the class, before requiring the
pupils either to do the work or to recite the lesson. In the
laboratory each pupil has a locker under his table, furnished
with apparatus, as specified in the Appendix. Each has also the
author's "Laboratory Manual," which contains on every left-hand
page full directions for an experiment, with observations to be
made, etc. The right-hand page is blank, and on that the pupil
makes a record of his work. These notes are examined at the time,
or subsequently, by the teacher, and the pupil is not allowed to
take the book from the laboratory; nor can he use any other book
on Chemistry while experimenting. By this means he learns to make
his own observations and inferences.
For the benefit of the science and the added interest in the
study, it is earnestly recommended that teachers encourage pupils
to fit up laboratories of their own at home. This need not at
first entail a large outlay. A small attic room with running
water, a very few chemicals, and a little apparatus, are enough
to begin with; these can be added to from time to time, as new
material is wanted. In this way the student will find his love
for science growing apace.
While endeavoring, by securing an able corps of critics, and in
all other ways possible, to reduce errors to a minimum, the
author disclaims any pretensions to a work entirely free from
mistakes, holding himself alone responsible for any shortcomings,
and trusting to the leniency of teachers and critics.
The manuscript has been read by Prof. Henry Carmichael, Ph.D., of
Boston, and to his broad and accurate scholarship, as well as to
his deep personal interest in the work, the author is indebted
for much valuable and original matter. The following persons have
generously read the proof, as a whole or in part, and made
suggestions regarding it, and to them the author would return his
thanks, as well as acknowledge his obligation: Prof. E. J.
Bartlett, Dartmouth College, N.H.; Prof. F. C. Robinson, Bowdoin
College, Me.; Prof. H. S. Carhart, Michigan University; Prof. B.
D. Halsted, Iowa Agricultural College; Prof. W. T. Sedgwick,
Institute of Technology, Boston; Pres. M. E. Wadsworth, Michigan
Mining School; Prof. George Huntington, Carleton College, Minn.;
Prof. Joseph Torrey, Iowa College; Mr. C. J. Lincoln, East Boston
High.School; Mr. W. H. Sylvester, English High School, Boston;
Mr. F. W. Gilley, Chelsea, Mass., High School; the late D. S.
Lewis, Chemist of the Boston Gas Works, and others.
R. P. W.
BOSTON, January 3, 1888.
TABLE OF CONTENTS
CHAPTER I.
THE METRIC SYSTEM.
Length.--Volume.--Weight
CHAPTER II.
DIVISIBILITY OF MATTER.
Mass.-Molecule.--Atom.--Element.--Compound.--Mixture.--
Analysis.--Synthesis.--Metathesis.--Chemism
CHAPTER III.
MOLECULES AND ATOMS.
Synthesis
CHAPTER IV.
ELEMENTS AND BINARIES.
Symbols.--Names.--Coefficients.--Exponents.--Table of elements
CHAPTER V.
MANIPULATION.
To prepare and cut glass, etc.
CHAPTER VI.
OXYGEN.
Preparation.--Properties.--Combustion of carbon; sulphur;
phosphorus; iron.
Chapter VII
NITROGEN
Separation--Properties
CHAPTER VIII
HYDROGEN
Preparation--Properties--Combustion--Oxy-hydrogen blowpipe
CHAPTER IX
UNION BY WEIGHT
Meaning of equations--Problems
CHAPTER X
CARBON
Preparation--Allotropic forms: diamond, graphite, amorphous
carbon, coke, mineral coal.--Carbon a reducing agent, a
decolorizer, disinfectant, absorber of gases
CHAPTER XI
VALENCE
Poles of attraction--Radicals
CHAPTER XII
ELECTRO-CHEMICAL RELATION OF ELEMENTS
Deposition of silver; copper; lead--Table of metals and non-
metals, and discussion of their differences
CHAPTER XIII.
ELECTROLYSIS.
Decomposition of water and of salts--Conclusions CHAPTER XIV.
UNION BY VOLUME.
Avogadro's law and its applications.
CHAPTER XV.
ACIDS AND BASES.
Characteristics of acids and bases.--Anhydrides.--Naming of
acids.--Alkalies
CHAPTER XVI.
SALTS.
Preparation from acids and bases.--Naming of salts.--Occurrence
CHAPTER XVII
CHLORHYDRIC ACID.
Preparation and tests.--Bromhydric, iodhiydric, and fluorhydric
acids.--Etching glass
CHAPTER XVIII.
NITRIC ACID.
Preparation, properties, tests, and uses.--Aqua regia:
preparation and action
CHAPTER XIX.
SULPHURIC ACID.
Preparation, tests, manufacture, and importance.-Fuming sulphuric
acid
CHAPTER XX.
AMMONIUM HYDRATE.
Preparation of bases.--Formation, preparation, tests, and uses of
ammonia.
Chapter XXI.
SODIUM HYDRATE.
Preparation and properties.--Potassium hydrate and calcium
hydrate
CHAPTER XXII
OXIDES OF NITROGEN.
Nitrogen monoxide, dioxide, trioxide, tetroaide, pentoxide.
CHAPTER XXIII.
LAWS OF DEFINITE AND OF MULTIPLE PROPORTION, and their
application
CHAPTER XXIV.
CARBON PROTOXIDE and water gas.
CHAPTER XXV.
CARBON DIOXIDE.
Preparation and tests.--Oxidation in the human system.--Oxidation
in water.--Deoxidation in plants
CHAPTER XXVI.
OZONE.
Description, preparation, and test
CHAPTER XXVII
CHEMISTRY OF THE ATMOSPHERE.
Constituents of the air.--Air a mixture.--Water, carbon dioxide,
and other ingredients of the atmosphere
CHAPTER XXVIII.
THE CHEMISTRY OF WATER.
Distillation of water.--Three states.--Pure water, sea-water,
river-water, spring-water CHAPTER XXIX.
THE CHEMISTRY OF FLAME.
Candle flame.--Bunsen flame.--Light and heat.--Temperature of
combustion.--Oxidizing and reducing flames.--Combustible and
supporter.--Explosive mixture of gases.--Generalizations
CHAPTER XXX.
CHLORINE.
Preparation.--Chlorine water.--Bleaching properties.--
Disinfecting power.--A supporter of combustion.--Sources and uses
CHAPTER XXXI.
BROMINE.
Preparation.--Tests.--Description.--Uses
CHAPTER XXXII.
IODINE.
Preparation.--Tests.--Iodo-starch paper.--Occurrence.--Uses.--
Fluorine
CHAPTER XXXIII.
THE HALOGENS.
Comparison.--Acids, oxides, and salts
CHAPTER XXXIV.
VAPOR DENSITY AND MOLECULAR WEIGHT.
Gaseous weights and volumes.--Vapor density defined.--Vapor
density of oxygen
CHAPTER XXXV.
ATOMIC WEIGHT.
Definition.--Atomic weight of oxygen.--Molecular symbols.--
Molecular and atomic volumes CHAPTER XXXVI.
DIFFUSION AND CONDENSATION OF GASES.
Diffusion of gases.--Law of diffusion.--Cause.--Liquefaction and
solidification of gases
CHAPTER XXXVIL
SULPHUR.
Separation.--Crystals from fusion.--Allotropy.--Solution.--
Theory of Allotropy.--Occurrence and purification.--Uses.---
Sulphur dioxide
CHAPTER XXXVIII.
HYDROGEN SULPHIDE.
Preparation.--Tests.--Combustion.--Uses.--An analyzer of metals.-
-Occurrence and properties
CHAPTER XXXIX.
PHOSPHORUS.
Solution and combustion.--Combustion under water.--Occurrence.--
Sources.--Preparation of phosphates and phosphorus.---
Properties.--Uses.--Matches.--Red phosphorus.---Phosphene
CHAPTER XL.
ARSENIC.
Separation.--Tests.--Expert analysis.--Properties and
occurrence.-- Atomic volume.--Uses of arsenic trioxide
CHAPTER XLI.
SILICON, SILICA, AND SILICATES.
Comparison of silicon and carbon.--Silica.--Silicates.--Formation
of silica.
Chapter XLII
GLASS AND POTTERY.
Glass an artificial silicate.--Manufacture.--Importance.--
Porcelain and pottery.
CHAPTER XLIII.
METALS AND THEIR ALLOYS.
Comparison of metals and non-metals.--Alloys.--Low fusibility. --
Amalgams
CHAPTER XLIV.
SODIUM AND ITS COMPOUNDS.
Order of derivation.--Occurrence and preparation of sodium
chloride; uses.--Sodium sulphate: manufacture and uses. --Sodium
carbonate: occurrence, manufacture, and uses.-- Sodium:
preparation and uses.--Sodium hydrate: preparation and use.--
Hydrogen sodium carbonate.--Sodium nitrate
CHAPTER XLV.
POTASSIUM AND AMMONIUM.
Occurrence and preparation of potassium.--Potassium chlorate and
cyanide.--Gunpowder.--Ammonium compounds
CHAPTER XLVI.
CALCIUM COMPOUNDS.
Calcium carbonate.--Lime and its uses.--Hard water.--Formation of
caves.--Calcium sulphate
CHAPTER XLVII.
MAGNESIUM, ALUMINIUM, AND ZINC.
Occurrence and preparation of magnesium.--Compounds of aluminium:
reduction; properties, and uses.--Compounds, uses, and reduction
of zinc CHAPTER XLVIII.
IRON AND ITS COMPOUNDS.
Ores of iron.--Pig-iron.--Steel.--Wrought-iron.--Properties. --
Salts of iron.--Change of valence and of color
CHAPTER XLIX.
LEAD AND TIN.
Distribution of lead.--Poisonous properties.--Some lead
compounds.-- Tin
CHAPTER L.
COPPER, MERCURY, AND SILVER.
Occurrence and uses of copper.--Compounds and uses of mercury.--
Occurrence, reduction, and salts of silver
CHAPTER LI.
PHOTOGRAPHY.
Description.
CHAPTER LII.
PLATINUM AND GOLD.
Methods of obtaining, and uses
CHAPTER LIII.
CHEMISTRY OF ROCKS.
Classification.--Composition.--Importance of siliceous rocks.--
Soils.--Minerals.--The earth's interior.--Percentage of elements
CHAPTER LIV.
ORGANIC CHEMISTRY.
Comparison of organic and inorganic compounds.--Molecular
differences.--Synthesis of organic compounds.--Marsh-gas.
series.---Alcohols.--Ethers.--Other substitution products. --
Olefines and other series.
CHAPTER LV.
ILLUMINATING GAS.
Source, preparation, purification, and composition.--Natural gas
CHAPTER LVI.
ALCOHOL.
Fermented and distilled liquors.--Effect on the system.--Affinity
for water.--Purity
CHAPTER LVII
OILS, FATS, AND SOAPS.
Sources and kinds of oils and fats.--Saponification.--Manufacture
and action of soap.--Glycerin, nitro-glycerin, and dynamite. --
Butter and oleomargarine.
CHAPTER LVIII
CARBO-HYDRATES.
Sugars.--Glucose.--Starch.--Cellulose.--Gun-cotton.--Dextrin. --
Zylonite
CHAPTER LIX.
CHEMISTRY OF FERMENTATION.
Ferments.--Alcoholic, acetic, and lactic fermentation.--
Putrefaction.--Infectious diseases
CHAPTER LX.
CHEMISTRY OF LIFE.
Growth of minerals and of organic life.--Food of plants and of
man.--Conservation of energy and of matter
CHAPTER LXI.
THEORIES.
The La Place theory--Theory of evolution--New theory of chemistry
CHAPTER LXII
GAS VOLUMES AND WEIGHTS.
Quantitative experiments with oxygen and hydrogen--Problems
AN INTRODUCTION TO CHEMICAL SCIENCE
CHAPTER I.
THE METRIC SYSTEM.
1. The Metric System is the one here employed. A sufficient
knowledge of it for use in the study of this book may be gained
by means of the following experiments, which should be performed
at the outset by each pupil.
2. Length.
Experiment 1.--Note the length of 10 cm. (centimeters) on a
metric ruler, as shown in Figure 1. Estimate by the eye alone
this distance on the cover of a book, and then verify the result.
Do the same on a t.t. (test-tube). Try this several times on
different objects till you can carry in mind a tolerably accurate
idea of 10 cm. About how many inches is it?
In the same way estimate the length of 1 cm, verifying each
result. How does this compare with the distance between two blue
lines of foolscap? Measure the diameter of the old nickel five-
cent piece.
Next, try in the same way 5 cm. Carry each result in mind, taking
such notes as may be necessary.
(Fig. 1)
3. Capacity.
Experiment 2.--Into a graduate, shown in Figure 2, holding 25 or
50 cc. (cubic centimeters) put 10 cc. of water; then pour this into
a t.t. Note, without marking, what proportion of the latter is
filled; pour out the water, and again put into the t.t. the same
quantity as nearly as can be estimated by the eye. Verify the
result by pouring the water back into the graduate. Repeat
several times until your estimate is quite accurate with a t.t.
of given size. If you wish, try it with other sizes. Now estimate
1 cc. of a liquid in a similar way. Do the same with 5 cc.
A cubic basin 10 cm on a side holds a liter. A liter contains
1,000 cc. If filled with water, it weighs, under standard
conditions, 1,000 grams. Verify by measurement.
4. Weight.
Experiment 3.--Put a small piece of paper on each pan of a pair
of scales. On one place a 10 g. (gram) weight. Balance this by
placing fine salt on the other pan. Note the quantity as nearly
as possible with the eye, then remove. Now put on the paper what
you think is 10 g. of salt. Verify by weighing. Repeat, as before,
several times. Weigh 1 g., and estimate as before. Can 1 g. of
salt be piled on a one-cent coin? Experiment with 5 g.
5. Resume--Lengths are measured in centimeters, liquids in cubic
centimeters, solids in grams. In cases where it is not convenient
to measure a liquid or weigh a solid, the estimates above will be
near enough for most experiments herein given. Different solids
of the same bulk of course differ in weight, but for one gram
what can be piled on a one-cent piece may be called a
sufficiently close estimate. The distance between two lines of
foolscap is very nearly a centimeter. A cubic centimeter is seen
in Figure 1. Temperatures are recorded in the centigrade scale.
CHAPTER II.
WHAT CHEMISTRY IS.
6. Divisibility of Matter.
Experiment 4.--Examine a few crystals of sugar, and crush them
with the fingers. Grind them as fine as convenient, and examine
with a lens. They are still capable of division. Put 3 g. of
sugar into a t.t., pour over it 5 cc. of water, shake well, boil
for a minute, holding the t.t. obliquely in the flame, using for
the purpose a pair of wooden nippers (Fig. 3). If the sugar does
not disappear, add more water. When cool, touch a drop of the
liquid to the tongue. Evidently the sugar remains, though in a
state too finely divided to be seen. This is called a solution,
the sugar is said to be soluble in water, and water to be a
solvent of sugar.
(Fig 3.)
Now fold a filter paper, as in Figure 4, arrange it in a funnel
(Fig. 5), and pour the solution upon it, catching what passes
through, which is called the filtrate, in another t.t. that rests
in a receiver (Fig. 5). After filtering, notice whether any
residue is left on the filter paper. Taste a drop of the
filtrate. Has sugar gone through the filter? If so, what do you
infer of substances in solution passing through a filter? Save
half the filtrate for Experiment 5, and dilute the other half
with two or three times its own volume of water. Shake well, and
taste.
(Fig 4.)
(Fig 5.) We might have diluted the sugar solution many times
more, and still the sweet taste would have remained. Thus the
small quantity of sugar would be distributed through the whole
mass, and be very finely divided.
By other experiments a much finer subdivision can be made. A
solution of.00000002 g. of the red coloring matter, fuchsine, in
1 cc. of alcohol gives a distinct color.
Such experiments would seem to indicate that there is no limit to
the divisibility of matter. But considerations which we cannot
discuss here lead to the belief that such a limit does exist;
that there are particles of sugar, and of all substances, which
are incapable of further division without entirely changing the
nature of the substance. To these smallest particles the name
molecules is given.
A mass is any portion of a substance larger than a molecule; it
is an aggregation of molecules.
A molecule is the smallest particle of a substance that can exist
alone.
A substance in solution may be in a more finely divided state
than otherwise, but it is not necessarily in its ultimate state
of division.
7. A Chemical Change.--Cannot this smallest particle of sugar,
the molecule, be separated into still smaller particles of
something else? May it not be a compound body, and will not some
force separate it into two or more substances? The next
experiment will answer the question.
Experiment 5.--Take the sugar solution saved from Experiment 4,
and add slowly 4 cc.of strong sulphuric acid. Note any change of
color, also the heat of the t.t. Add more acid if needed.
A substance entirely different in color and properties has been
formed. Now either the sugar, the acid, or the water has
undergone a chemical change. It is, in fact, the sugar. But the
molecule is the smallest particle of sugar possible. The acid
must have either added something to the sugar molecules, or
subtracted something from them. It was the latter. Here, then, is
a force entirely different from the one which tends to reduce
masses to molecules. The molecule has the same properties as the
mass. Only a physical force was used in dissolving the sugar, and
no heat was liberated. The acid has changed the sugar into a
black mass, in fact into charcoal or carbon, and water; and heat
has been produced. A chemical change has been brought about.
From this we see that molecules are not the ultimate divisions of
matter. The smallest sugar particles are made up of still smaller
particles of other things which do not resemble sugar, as a word
is composed of letters which alone do not resemble the word. But
can the charcoal itself be resolved into other substances, and
these into still others, and so on? Carbon is one of the
substances from which nothing else has been obtained. There are
about seventy others which have not been resolved. These are
called elements; and out of them are built all the compounds--
mineral, vegetable, and animal--which we know.
8. An element is a chemically indivisible substance, or one from
which nothing else can be extracted.
A compound is a substance which is made up of elements united in
exact proportions by a force called chemism, or chemical
affinity.
A mixture is composed of two or more elements or compounds
blended together, but not held by any chemical attraction.
To which of these three classes does sugar belong? Carbon? The
solution of sugar in water?
Carbon is an element; we call its smallest particle an atom.
An atom is the smallest particle of an element that can enter
into combination. Atoms are indivisible and usually do not exist
alone. Both elements and compounds have molecules.
The molecule of an element usually contains two atoms; that of a
compound may have two, or it may have hundreds. For a given
compound the number is always definite.
Chemism is the force that binds atoms together to form molecules.
The sugar molecule contains atoms, forty-five in all, of three
different elements: carbon, hydrogen, and oxygen. That of salt
has two atoms: one of sodium, one of chlorine. Should we say "an
atom of sugar"? Why? Of what is a mass of sugar made up? A
molecule? A mass of carbon? A molecule? Did the chemical affinity
of the acid break up masses or molecules? In this respect it is a
type of all chemical action. The distinction between physics and
chemistry is here well shown. The molecule is the unit of the
physicist, the atom that of the chemist. However large the masses
changed by chemical action, that action is always on the
individual molecule, the atoms of which are separated. If the
molecule were an indivisible particle, no science of chemistry
would be possible. The physicist finds the properties of masses
of matter and resolves them into molecules, the chemist breaks up
the molecule and from its atoms builds up other compounds.
Analysis is the separation of compounds into their elements.
Synthesis is the building up of compounds from their elements.
Of which is the sugar experiment an example? Metathesis is an
exchange of atoms in two different compounds; it gives rise to
still other compounds.
A chemical change may add something to a substance, or subtract
something from it, or it may both subtract and add, making a new
substance with entirely different properties. Sulphur and carbon
are two stable solids. The chemical union of the two forms a
volatile liquid. A substance may be at one time a solid, at
another a liquid, at another a gas, and yet not undergo any
chemical change, because in each case the chemical composition is
identical.
State which of these are chemical changes: rusting of iron,
falling of rain, radiation of heat, souring of milk, evaporation
of water, decay of vegetation, burning of wood, breaking of iron,
bleaching of cloth. Give any other illustrations that occur to
you.
Chemistry treats of matter in its simplest forms, and of the
various combinations of those simplest forms.
CHAPTER III.
MOLECULES AND ATOMS.
9. Molecules are Extremely Small.--It has been estimated that a
liter of any gas at 0 degrees and 760 mm. pressure contains 10^24
molecules, i.e. one with twenty-four ciphers.
Thomson estimates that if a drop of water were magnified to the
size of the earth, and its molecules increased in the same
proportion, they would be larger than fine shot, but not so large
as cricket balls.
A German has recently obtained a deposit of silver two-millionths
of a millimeter thick, and visible to the naked eye. The computed
diameter of the molecule is only one and a half millionths of a
millimeter.
By a law of chemistry there is the same number of molecules in a
given volume of every gas, if the temperature and pressure are
the same. Hence, all gaseous molecules are of the same size,
including, of course, the surrounding space. They are in rapid
motion, and the lighter the gas the more rapid the motion. This
gives rise to diffusion. See page 114.
10. We Know Nothing Definite of the Form of Molecules.--In this
book they will always be represented as of the same size, that of
two squares. A molecule is itself composed of atoms,--from two to
several hundred. The size of the atom of most elements we
represent by one square.11. Atoms.--If the gaseous molecules be
of the same size, it is clear that either the atoms themselves
must be condensed, or the spaces between them must be smaller
than before. We suppose the latter to be the case, and that they
do not touch one another, the same thing being true of molecules.
Atoms composing sugar must be crowded nearer together than those
of salt. These atoms are probably in constant motion in the
molecule, as the latter is in the mass. If we regard this square
as a mass of matter, the dots may represent molecules; if we call
it a molecule, the dots may be called atoms, though many
molecules have no more than two or three atoms.
The following experiments illustrate the union of atoms to form
molecules, and of elements to form compounds.
12. Union of Atoms.
Experiment 6.--Mix, on a paper, 5 g. of iron turnings, and the
same bulk of powdered sulphur, and transfer them to an ignition
tube, a tube of hard glass for withstanding high temperatures.
Hold the tube in the flame of a burner till the contents have
become red-hot. After a minute break it by holding it under a jet
of water. Put the contents into an evaporating-dish, and look for
any uncombined iron or sulphur. Both iron and sulphur are
elements. Is this an example of synthesis or of analysis? Why? Is
the chemical union between masses of iron and sulphur, or between
molecules, or between atoms? Is the product a compound, an
element, or a mixture?
Experiment 7.--Try the same experiment, using copper instead of
iron. The full explanation of these experiments is given on page
13.
CHAPTER IV.
ELEMENTS AND BINARIES.
13. About Seventy Different Elements are now recognized, half of
which have been discovered within little more than a century.
These differ from one another in (1) atomic weight, (2) physical
and chemical properties, (3) mode of occurrence, etc. Page 12
contains the most important elements.
The symbol of an element is usually the initial letter or letters
of its Latin name, and stands for one atom of the element. C is
the symbol for carbon, and represents one atom of it. O means one
atom of oxygen.[The symbols of elements will also be used in this
book to stand for an indefinite quantity of them; e.g. O will be
used for oxygen in general as well as for one atom. The text will
readily decide when symbols have a definite meaning, and when
they are used in place of words.] Write, explain, and memorize
the symbols of the elements in heavy type.
14. The Atomic Weight of an element is the weight of its atom
compared with that of hydrogen. H is taken as the standard
because it has the least atomic weight. The atomic weight of O is
16, which means that its atom weighs 16 times as much as the H
atom. Every symbol, then, stands for a definite weight of the
element, i.e. its atomic weight, as well as for its atom.
How much bromine by weight does Br stand for? What do these
symbols mean--As, Na, N, P? If O represents one atom, how much
does O2 or 2 O stand for? How much by weight? Most elements have
two atoms in the molecule. How many molecules in 6 H? 10 N? S8?
I20?
The symbol of a compound is formed by writing in succession the
symbols of the elements of which it is composed. How many atoms
in the following molecules, and how many of each element: C2H60?
HNO3? PbSO4? MgCl2? (Hg2(NO3)2?)
15. The Simplest Compounds are Binaries.--A binary is a substance
composed of two elements; e.g. common salt, which is a compound
of sodium and chlorine. Its symbol is NaCl, its chemical name
sodium chloride. The ending ide is applied to the last name of
binaries. How many parts by weight of Na and of Cl in NaCl? What
is the molecular weight, i.e. the weight of its molecule? Name
KCl. How many atoms in its molecule? Parts by weight of each
element? Molecular weight? Does the symbol stand for more than
one molecule? How many molecules in 4 NaCl? How many atoms of Na
and of Cl? Name these: HCl, NaBr, NaI, KBr, AgCl, AgI, HBr, HI,
HF, HgO, ZnO, ZnS, MgO, CaO. Compute the proportion by weight of
each element in the last three.
A coefficient before the symbol of a compound includes all the
elements of the symbol, and shows the number of molecules. How
many in these: 6 KBr? 3 Sn0? 12 NaCl? How many atoms of each
element in the above?
An exponent, always written below, applies only to the element
after which it is written, and shows the number of atoms. Explain
these: AuCl3, ZnCl2, Hg2Cl2.
Write symbols for four molecules of sodium bromide, one of silver
iodide (always omit coefficient one), eight of potassium bromide,
ten of hydrogen chloride; also for one molecule of each of these:
hydrogen fluoride, potassium iodide, silver chloride.
In all the above cases the elements have united atom for atom.
Some elements will not so unite. In CaCl2 how many atoms of each
element? Parts by weight of each? Give molecular weight. Is the
size of the molecule thereby changed? Name these, give the number
of atoms of each element in the molecule, and the proportion by
weight, also their molecular weights: AuCl3, ZnCl2, MnCl2, Na2O,
K2S, H3P, H4C.
Principal Elements.
Name. Sym. At. Wt. Valence. Vap.D. At.Vol. Mol.Vol. State.
Aluminium Al 27. II, IV ... ... ... Solid
Antimony Sb 120. III, V. ... ... ... "
Arsenic As 75. III, V 150. "
Barium Ba 137. II ... ... ... "
Bismuth Bi 210. III, V ... ... ... "
Boron B 11. III ... ... ... "
Bromine Br 80. I, (V) 80. Liquid
Cadmium Cd 112. II 56. Solid
Calcium Ca 40. II ... ... ... "
Carbon C 12. (II), IV ... ... ... "
Chlorine Cl 35.5 I, (V) 35.5 Gas
Chromium Cr 52. (II),IV,VI ... ... ... Solid
Cobalt Co 59. II, IV ... ... ... Gas
Copper Cu 63. I, II ... ... ... "
Fluorine F 19. I, (V) ... ... ... Gas
Gold Au 196. (I), III ... ... ... Solid
Hydrogen H 1. I 1. Gas
Iodine I 127. I, (V) 127. ... ... Solid
Iron Fe 56. II,IV,(VI) ... ... ... "
Lead Pb 206. II, IV ... ... ... "
Lithium Li 7. I ... ... ... "
Magnesium Mg 24. II ... ... ... "
Manganese Mn 55. II, IV, VI ... ... ... "
Mercury Hg 200. I, II 100. Liquid
Nickel Ni 59. II, IV ... ... ... Solid
Nitrogen N 14. (I),III,V 14. Gas
Oxygen O 16. II 16. "
Phosphorus P 31. (I),III, V 62. Solid
Platinum Pt 197. (II), IV ... ... ... "
Potassium K 39. I ... ... ... "
Silicon Si 28. IV ... ... ... "
Silver Ag 108. I ... ... ... "
Sodium Na 23. I ... ... ... "
Strontium Sr 87. II ... ... ... "
Sulphur S 32. II,IV,(VI) 32(96) "
Tin Sn 118. II, IV ... ... ... "
Zinc Zn 65. II 32.5 "
If more than one atom of an element enters into the composition
of a binary, a prefix is often used to denote the number. SO2 is
called sulphur dioxide, to distinguish it from SO3, sulphur
trioxide. Name these: CO2, SiO2, MnO2. The prefixes are: mono or
proto, one; di or bi, two; tri or ter, three; tetra, four; pente,
five; hex, six; etc. Diarsenic pentoxide is written, As2O5.
Symbolize these: carbon protoxide, diphosphorus pentoxide,
diphosphorus trioxide, iron disulphide, iron protosulphide. Often
only the prefix of the last name is used.
16. An Oxide is a Compound of Oxygen and Some Other Element, as
HgO. What is a chloride? Define sulphide, phosphide, arsenide,
carbide, bromide, iodide, fluoride.
In Experiment 6, where S and Fe united, the symbol of the product
was FeS. Name it. How many parts by weight of each element? What
is its molecular weight? To produce FeS a chemical union took
place between each atom of the Fe and of the S. We may express
this reaction, i.e. chemical action, by an equation:--
Iron + Sulphur = Iron Sulphide
Or, using symbols Fe + S = FeS
Using atomic weights, 56 32 = 88.
These equations are explained by saying that 56 parts by weight
of iron unite chemically with 32 parts by weight of sulphur to
produce 88 parts by weight of iron sulphide. This, then,
indicates the proportion of each element which combines, and
which should be taken for the experiment. If 56 g. of Fe be used,
32 g. of S should be taken. If we use more than 56 parts of Fe
with 32 of S, will it all combine? If more than 32 of S with 56
of Fe? There is found to be a definite quantity of each element
in every chemical compound. Symbols would have no meaning if this
were not so.
Write and explain the equation for the experiment with copper and
sulphur, using names, symbols, and weights, as above.
CHAPTER V.
MANIPULATION.
17. To Break Glass Tubing.
Experiment 8.--Lay the tubing on a flat surface, and draw a sharp
three-cornered file two or three times at right angles across it
where it is to be broken, till a scratch is made. Take the tube
in the hands, having the two thumbs nearly opposite the scratch,
and the fingers on the other side. Press outward quickly with the
thumbs, and at the same time pull the hands strongly apart, and
the tubing should break squarely at the scratch.
To break large tubing, or cut off bottles, lamp chimneys, etc.,
first make a scratch as before; then heat the handle of a file,
or a blunt iron--in a blast-lamp flame by preference--till it is
red-hot, and at once press it against the scratch till the glass
begins to crack. The fracture can be led in any direction by
keeping the iron just in front of it. Re-heat the iron as often
as necessary.
18. To Make Ignition-Tubes.
Experiment 9.--Hold the glass tubing between the thumb and
forefinger of each hand, resting it against the second finger.
Heat it in the upper flame, slowly at first, then strongly, but
heat only a very small portion in length, and keep it in constant
rotation with the right hand. Hold it steadily, and avoid
twisting it as the glass softens. The yielding is detected by the
yellow flame above the glass and by an uneven pressure on the
hands. Pull it a little as it yields, then heat a part just at
one side of the most softened portion. Rotate constantly without
twisting, and soon it can be separated into two closed tubes. No
thread should be attached; but if there be one, it can be broken
off and the end welded. The bottom can be made more symmetrical
by heating it red-hot, then blowing, gradually, into the open
end, this being inserted in the mouth. The parts should be
annealed by holding above the flame for a short time, to cool
slowly.
For hard glass--Bohemian--or large tubes, the blast-lamp or
blowpipe is needed. In the blast-lamp air is forced out with
illuminating gas. This gives a high degree of heat. Bulbs can be
made in the same way as ignition-tubes, and thistle-tubes are
made by blowing out the end of a heated bulb, and rounding it
with charcoal.
19. To Bend Glass Tubing.
Experiment 10.--Hold the tube in the upper flame. Rotate it so as
to heat all parts equally, and let the flame spread over 3 or 4
cm. in length. When the glass begins to yield, without removing
from the flame slowly bend it as desired. Avoid twisting, and be
sure to have all parts in the same plane; also avoid bending too
quickly, if you would have a well-rounded joint. Anneal each bend
as made. Heated glass of any kind should never be brought in
contact with a cool body. For making O, H, etc., a glass tube --
delivery-tube--50 cm. long should have three bends, as in Figure
6. The pupil should first experiment with short pieces of glass,
10 or 15 cm. long. An ordinary gas flame is the best for bending
glass.
20. To Cut Glass.
Experiment 11.--Lay the glass plate on a flat surface, and draw a
steel glass-cutter--revolving wheel--over it, holding this
against a ruler for a guide, and pressing down hard enough to
scratch the glass. Then break it by holding between the thumb and
fingers, having the thumbs on the side opposite to the scratch,
and pressing them outward while bending the ends of the glass
inward. The break will follow the scratch.
Holes can be bored through glass and bottles with a broken end of
a round file kept wet with a solution of camphor in oil of
turpentine.
21. To Perforate Corks.
Experiment 12.--First make a small hole in the cork with the
pointed handle of a round--rat-tail--file. Have the hole
perpendicular to the surface of the cork. This can be done by
holding the cork in the left hand and pressing against the larger
surface, or upper part, of the cork, with the file in the right
hand. Only a mere opening is made in this way, which must be
enlarged by the other end of the file. A second or third file of
larger size may be employed, according to the size of the hole to
be made, which must be a little smaller than the tube it is to
receive, and perfectly round.
CHAPTER VI.
OXYGEN.
22. To Obtain Oxygen.
Experiment 13.--Take 5 g. of crystals of potassium chlorate
(KClO3) and, without pulverizing, mix with the same weight of
pure powdered manganese dioxide (MnO2). Put the mixture into a
t.t., and insert a d.t.--delivery-tube--having the cork fit
tightly. Hang it on a r.s.--ring-stand,-- as in Figure 7, having
the other end of the d.t.
(Fig 7.)
under the shelf, in a pneumatic trough, filled with water just
above the shelf. Fill three or more receivers--wide-mouthed
bottles--with water, cover the mouth of each with a glass plate,
invert it with its mouth under water, and put it on the shelf of
the trough, removing the plate. No air should be in the bottles.
Have the end of the d.t. so that the gas will rise through the
orifice. Hold a lighted lamp in the hand, and bring the flame
against the mixture in the t.t. Keep
the lamp slightly in motion, with the hand, so as not to break
the t.t. by over-heating in one place. Heat the mixture strongly,
if necessary. The upper part of the t.t. is filled with air:
allow this to escape for a few seconds; then move a receiver over
the orifice, and fill it with gas. As soon as the lamp is taken
away, remove the d.t. from the water. The gas contracts, on
cooling, and if not removed, water will be drawn over, and the
t.t. will be broken. Let the t.t. hang on the r.s. till cool.
With glass plates take out the receivers, leaving them covered,
mouth upward (Fig. 8), with little or no water inside. When cool,
the t.t. may be cleaned with water, by covering its mouth with
the thumb or hand, and shaking it vigorously.
What elements, and how many, in KClO3? In Mn02? It is evident
that each of these compounds contains O. Why, then, could we not
have taken either separately, instead of mixing the two? This
could have been done at a sufficiently high temperature. Mu02
requires a much higher temperature for dissociation, i.e.
separation into its elements, than KClO3, while a mixture of the
two causes O to come off from KClO3 at a lower temperature than
if alone. It is not known that Mn02 suffers any change.
Each molecule of potassium chlorate undergoes the following
change:--
Potassium Chlorate = Potassium Chloride + Oxygen
KClO3 = KCl + 3 O.
Is this analysis or synthesis? Complete the equation, by using
weights, and explain it. Notice whether the right- hand member of
the equation has the same number of atoms as the left. Has
anything been lost or gained? What element has heat separated?
Does the experiment show whether O is very soluble in water? How
many grams of O are obtainable from 122.58 g. KCIO3? PROPERTIES.
23. Combustion of Carbon.
OXYGEN Experiment 14.--Examine the gas in one of the receivers.
Put a lighted splinter into the receiver, sliding along the glass
cover. Remove it, blow it out, and put in again while glowing. Is
it re-kindled? Repeat till it will no longer burn. Is the gas a
supporter of combustion? How did the combustion compare with that
in air? Is it probable that air is pure O? Why did the flame at
last go out? Has the O been destroyed, or chemically united with
something else?
Wood is in part C. CO2 is formed by the combustion; name it. The
equation is C + 2O = CO2. Affix the names and weights. Is CO2 a
supporter of combustion? Note that when C is burned with plenty
of O, CO2 is always formed, and that no matter how great the
conflagration, the union is atom by atom. Combustion, as here
shown, is only a rapid union of O with some other substance, as C
or H.
24. Combustion of Sulphur.
Experiment 15.--Hollow out one end of a piece of electric-light
pencil, or of crayon, 3 cm. long, and attach it to a Cu wire
(Fig. 9). Put into this a piece of S as large as a pea, ignite it
by holding in the flame, and then hold it in a receiver of O.
Note the color and brightness of the flame, and compare with the
same in the air. Also note the color and odor of the product. The
new gas is SO2. Name it, and write the equation for its
production from S and O. How do you almost daily perform a
similar experiment? Is the product a supporter of combustion?
25. Combustion of Phosphorus.
Experiment 16.--With forceps, which should always be used in
handling this element, put a bit of P, half as large as the S
above,into the crayon, called a deflagrating-spoon. Heat another
wire, touch it to the P, and at once lower the latter into a
receiver of O. Notice the combustion, the color of the flame and
of the product. After removing, be sure to burn every bit of P by
holding it in a flame, as it is liable to take fire if left. The
product of the combustion is a union of what two elements? Is it
an oxide? Its symbol is P2O5. Write the equation, using symbols,
names, and weights. Towards the close of the experiment, when the
O is nearly all combined, P2O3 is formed, as it is also when P
oxidizes at a low temperature. Name it and write the equation.
26. Combustion of Iron.
Experiment 17.--Take in the forceps a piece of iron picture-cord
wire 6 or 8cm long, hold one end in the flame for an instant,
then dip it into some S. Enough S will adhere to be set on fire
by holding it in the flame again. Then at once dip it into a
receiver of O with a little water in the bottom. The iron will
burn with scintillations. Is this analysis or synthesis? What
elements combine? A watch-spring, heated to take out the temper,
may be used, but picture-wire is better.
The product is Fe3O4. Write the equation. How much Fe by weight
in the formula? How much O? What per cent by weight of Fe in the
compound? Multiply the fractional part by 100. What per cent of
0? Whatper cent of C0 .is C? O2? Find the percentage composition
of SO2. P2O5.
From the last five experiments what do you infer of the tendency
of O to unite with other elements?
27. Oxygen is a Gas without Color, Odor, or Taste.
It is chemically a very active element; that is, it unites with
almost everything. Fluorine is the only element with which it
will not combine. When oxygen combines with a single element,
what is the compound called? We have found that O makes up a
certain portion of the air; later, we shall see how large the
proportion is. Its tendency to combine with almost everything is
a reason for the decay, rust, and oxidation of so many
substances, and for conflagrations, great and small. New
compounds are thusformed, of which O constitutes one factor.
Water, H2O, is only a chemical union of O and H. Iron rust, Fe2O3
and H2O, is composed of O, Fe, and water. The burning of wood or
of coal gives rise to carbon dioxide, CO2, and water. Decay of
animal and vegetable matter is hastened by this all-pervading
element. O forms a portion of all animal and vegetable matter, of
almost all rocks and minerals, and of water. It is the most
abundant of all elements, and makes up from one-half to two-
thirds of the earth's surface. Compute the proportion of it, by
weight, in water, H2O. It is the union of O in the air with C and
H in our blood that keeps up the heat of the body and supports
life. See page 81.
There are many ways of preparing this element besides the one
given above. It may be obtained from water (Experiment 38) and
from many other compounds, e.g. by heating mercury oxide,
HgO.
CHAPTER VII.
NITROGEN.
28. Separation.
Experiment 18.--Fasten a piece of electric-light pencil, or of
crayon, to a wire, as in Experiment 15, and bend the wire so it
will reach half-way to the bottom of a receiver. Using forceps,
put into the crayon a small piece of phosphorus. Pass the wire up
through the orifice in the shelf of a p.t. (pneumatic trough),
having water at least l cm. above the shelf. Heat another wire,
touch it to the P, and quickly invert an empty receiver over the
P, having the mouth under water, so as to admit no air (Fig. 10).
Let the P burn as long as it will, then remove the wire and the
crayon, letting in no air. Note the color of the product, and
leave till it is tolerably clear, then remove the receiver with a
glass plate, leaving the water in the bottom.
Do the fumes resemble those of Experiment 16? Does it seem likely
(Fig 10.) that part of the air is O? Why a part only? Find what
proportion of the receiver is filled with water by measuring the
water with a graduate; then fill it with water and measure that;
compute the percentage which the former is of the latter. What
proportion of the air, then, is O? What was the only means of
escape for the P2O6, and P2O2 formed? These products are solids.
Are they soluble in water? Compute the percentage composition,
always by weight, of P2O2 and P2O5.
The gas left in the receiver is evidently not O. Experiment 19
will prove this conclusively, and show the properties of the new
gas.
29. Properties.
Experiment 19.--When the white cloud has disappeared, slide the
plate along, and insert a burning stick; try one that still
glows.
See whether the P and S on the end of a match will burn. Is the
gas a supporter of combustion? Since it does not unite with C,
S, or P, is it an active or a passive element? Compare it with
O. Air is about 14 1/2 times as heavy as H. Which is heavier, air
or N? See page 12. Air or O?
Write out the chief properties, physical and chemical, of N, as
found in this experiment.
30. Inactivity of N.--N will scarcely unite chemically except on
being set free from compounds. It has, however, an intense
affinity for boron, and will even go through a carbon crucible to
unite with it. It is not combined with O in the air; but the two
form a mixture (page 86), of which N makes up four-fifths, its
use being to dilute the O. What would be the effect, in case of a
fire, if air were pure O? What effect on the human system?
Growing plants need a great deal of N, but they are incapable of
making use of that in the air, on account of the chemical
inactivity of the element. Their supply comes from compounds in
earth, water, and air. By reason of its inertness N is very
easily set free from its compounds. For this reason it is a
constituent of most explosives, as gunpowder, nitro-glycerine,
dynamite, etc. These solids, by heat or concussion, are suddenly
changed to gases, which thereby occupy much more space, causing
an explosion.
Nitrogen exists in many compounds, such as the nitrates; but the
great source of it all is the atmosphere. See page 85.
CHAPTER VIII.
HYDROGEN.
31. Preparation.
Experiment 20.--Prepare apparatus as for making O. Be sure that
the cork perfectly fits both d.t. and t.t., or the H will escape.
Cover 5 g. granulated Zn, in the t.t., with 10 cc. H2O, and add 5
cc. chlorhydric acid, HCl. Adjust as for O (Fig. 7), except that
no heat is to be applied. If the action is not brisk enough, add
more HCl. Collect several receivers of the gas over water, adding
small quantities of HCl when necessary. Observe the black
floating residuum; it is carbon, lead, etc. With a glass plate
remove the receivers, keeping them inverted (Fig. 11), or the H
will escape.
32. The Chemical Change is as follows:--
Zinc + hydrogen chloride = zinc chloride + hydrogen.
Zn + 2 HCl = ZnCl2 + 2H.
Complete by adding the weights, and explain. Notice that the
water does not take part in the change; it is added to dissolve
the ZnCl2 formed, and thus keep it from coating the Zn and
preventing further action of the acid. Note also that Zn has
simply changed places with H, one atom of the former having
driven off two atoms of the latter. The H, having nothing to
unite with, is set free as a gas, and collected over water. Of
course Zn must have a stronger chemical affinity for Cl than H
has, or the change could not have taken place. Why one Zn atom
replaces two H atoms will be explained later, asfar as an
explanation is possible. This equation, should be studied
carefully, as a type of all equations. The left-hand member shows
what were taken, i.e. the factors; the right-hand shows what were
obtained, i.e. the products. H2SO4 might have been used instead
of HCl. In that case the reaction, or equation, would have been:
--
Zinc + hydrogen sulphate = zinc sulphate + hydrogen.
Zn + H2SO4 = ZnSO4 + 2H.
Iron might have been used instead of zinc, in which case the
reactions would have been:--
Iron + hydrogen chloride = iron chloride + hydrogen.
Fe + 2 HCl = FeCl2 + 2 H.
Iron + hydrogen sulphate = iron sulphate + hydrogen.
Fe + H2SO4 = FeSO4 + 2 H.
Write the weights and explain the equations. The latter should be
memorized.
33. Properties.
Experiment 21.--Lift with the left hand a receiver of H, still
inverted, and insert a burning splinter with the right (Fig. 12).
Does the splinter continue to burn? Does the gas burn? If so,
where? Is the light brilliant? Note the color of the flame. Is
there any explosion? Try this experiment with several receivers.
Is the gas a supporter of combustion? i.e. will carbon burn in
it? Is it combustible? i.e. does it burn? If so, it unites with
some part of the air. With what part?34. Collecting H by Upward
Displacement.
Experiment 22.--Pass a d.t. from a H generator to the top of a
receiver or t.t. (Fig. 13). The escaping H being so much lighter
than air will force the latter down. To obtain the gas unmixed
with air, the d.t. should tightly fit a cardboard placed under
the mouth of the receiver. When filled, the receiver can be
removed, inverted as usual, and the gas tested. In this and other
experiments for generating H, a thistle-tube, the end of which
dips under the liquid, can be used for pouring in acid, as in
Figure 13.
35. Philosopher's Lamp and Musical Flame.
Experiment 23.--Fit to a cork a piece of glass tubing 10 or 15
cm. long, having the outer end drawn out to a point with a small
opening, and insert it in the H generator. Before igniting the
gas at the end of the tube take the, precaution to collect a t.t.
of it by upward displacement, and bring this in contact with a
flame. If a sharp explosion ensues, air is not wholly expelled
from the generator, and it would be dangerous to light the gas.
When no sound, or very little, follows, light the escaping gas.
The generation of H must not be too rapid, neither should the
t.t. be held under the face, as the cork is liable to be forced
out by the pressure of H. A safety-tube, similar to the thistle-
tube above, will prevent this. This apparatus is called the
"philosopher's lamp." Thrust the flame into a long glass tube 1-
1/2 to 3 cm. in diameter, as shown in Figure 14, and listen for a
musical note.
36. Product of Burning H in Air.
Experiment 24.--Fill a tube 2 or 3 cm. in diameter with calcium
chloride, CaCl2, and connect one end with a generator of H (Fig.
15). At the other end have a philosopher's lamp-tube.Observing
the usual precautions, light the gas and hold over it a receiver,
till quite a quantity of moisture collects. All water was taken
from the gas by the dryer, CaCl2. What is, therefore, the product
of burning H in air? Complete this equation and explain it: 2H +
O = ? Figure 16 shows a drying apparatus arranged to hold CaCl2.
[Fig. 15][Fig. 16]
37. Explosiveness of H.
Experiment 25. -- Fill a soda-water bottle of thick glass with
water, invert it in a pneumatic trough, and collect not over 1/4
full of H. Now remove the bottle, still inverted, letting air in
to fill the other 3/4. Mix the air and H by covering the mouth of
the bottle with the hand, and shaking well; then hold the mouth
of the bottle, slightly inclined, in a flame. Explain the
explosion which follows. If 3/4 was air, what part was O? What
use did the N serve? Note any danger in exploding H mixed with
pure O. What proportions of O and H by volume would be most
dangerously explosive? What proportion by weight?
By the rapid union of the two elements, the high temperature
suddenly expanded the gaseous product, which immediately
contracted; both expansion and contraction produced the noise of
explosion.
38. Pure H Is a Gas without Color, Odor, or Taste.
--It is the lightest of the elements, 14 1/2 times as light
asair. It occurs uncombined in coal-mines, and some other places,
but the readiness with which it unites with other elements,
particularly O, prevents its accumulation in large quantities. It
constitutes two-thirds of the volume of the gases resulting from
the decomposition of water, and one-ninth of the weight. Compute
the latter from its symbol. It is a constituent of plants and
animals, and some rocks. Considering the volume of the ocean, the
total amount of H is large. It can be separated from H2O by
electrolysis, or by C, as in the manufacture of water gas.
When burned with O it forms H2O. Pure O and H when burning give
great heat, but little light. The oxy-hydrogen blow-pipe (Fig.
17) is a device for producing the highest temperatures of
combustion. It has O in the inner tube and H in the outer. Why
would it not be better the other way? These unite at the end, and
are burned, giving great heat. A piece of lime put into the flame
gives the brilliant Drummond or calcium light.
Chapter IX. UNION BY WEIGHT.
39. In the Equation --
Zn + 2 HCl = ZnCl2 + 2 H
65 + 73 = 136 + 2
65 parts by weight of Zn are required to liberate 2 parts by weight of
H; or, by using 65 g Zn with 73 g HCl, we obtain 2 g H. If twice as
much Zn (130 g) were used, 4 g H could be obtained, with, of course,
twice as much HCl. With 260 g. Zn, how much H could be liberated?
A proportion may be made as follows:--
Zn given : Zn required :: H given : H required.
65 : 260 :: 2 : x.
[footnote: Given, as here used, means the weight called for by the
equation; required means that called for by the question.]
Solving, we have 8 g H.
How much H is obtainable by using 5 g Zn, as in the experiment?
To avoid error in solving similar problems, the best plan is as
follows:--
Zn + 2HCl = ZnCl2 + 2 H | 65:5::2:x
65 2 | 65 x = 10
5 x | x = 10/65 = 2/13 Ans. 2/13 g.
The equation should first be written; next, the atomic or molecular
weights which you wish to use, and only those, to avoid confusion;
then, on the third line, the quantity of the substance to be used, with
underneath the substance wanted. The example above will best
how this. This plan will prevent the possibility of error. The proportion
will then be:--
a given : a required :: b given : b required.
How much Zn is required to produce 30 g. H?
Zn + 2HCl = ZnCl2 + 2H | 2:30::65:x
65 2 | 2x = 1950
x 30 | x = 975 Ans. 975 g. Zn.
Solve:--
(1) How much Zn is necessary for 14 g. H?
(2) How many pounds of Zn are necessary for 3 pounds of H?
(3) How many grams of H from 17 g. of Zn?
(4) How many tons of H from 1/2 ton of Zn?
Suppose we wish to find how much chlorhydric acid--pure gas--
will give 12 g. H. The question involves only HCl and H. Arrange
as follows:--
Zn + 2HCl = ZnCl2 + 2 H | H giv. : H req. :: HCl giv. : HCl req.
73 2 | 2 : 12 :: 73 x
x 12 | 2x=876 x=438
Ans. 438 g. HCl.
Solve:--
(1) How much HCl is needed to produce 100 g. H?
(2) How much H in 10 g. HCl?
(3) How much ZnCl2 is formed by using 50 g. HCl? The question
is now between HCl and ZnCl2.
Zn + 2HC1 = ZnCl2 + 2H
73 136 | Arrange the proportion, and solve.
50 x
Suppose we have generated H by using H2S04: the equation is
Zn + H2S04 = ZnSO4 + 2 H. There is the same relation as before
between the quantities of Zn and of H, but the H2S04 and ZnS04 are
different.
How much H2SO4 is needed to generate 12 g. H?
Zn + H2SO4 = ZnS04 + 2 H
98 2 | Make the proportion, and solve
x 12
Solve:--
(1) How much H in 200 g. H2S04?
(2) How much ZnS04 is produced from 200 g. H2S04?
(3) How much H2S04 is needed for 7 1/2 g H?
(4) How much Zn will 40 g. H2SO4 combine with?
(5) How much Fe will 40 g. H2SO4 combine with?
(6) How much H can be obtained by using 75 g Fe?
These principles apply to all reactions. Suppose, for example, we
wish to get l0 g. of O: how much KClO3 will it be necessary to use?
The reaction is:--
KClO3 = KCl + O3 | 48 : 10 :: 122.5 : x
122.5 48 |
x 10 | Ans. 25.5+ g. KClO3.
The pupil should be required to make up problems of his own,
using various reactions, and to solve them.
CHAPTER X.
CARBON.
Examine graphite, anthracite coal, bituminous coal, cannel coal,
wood, gas carbon, coke.
40. Preparation of C.
Experiment 26.--Hold a porcelain dish or a plate in the flame of
a candle, or of a Bunsen burner with the openings at the bottom
closed. After a minute examine the deposit. It is carbon, i.e.
lamp- black or soot, which is a constituent of gas, or of the
candle. Open the valve at the base of the Bunsen burner, and hold
the deposit in the flame. Does the C gradually disappear? If so,
it has been burned to CO2. C + 2 O = CO2. Is C a combustible
element?
Experiment 27.--Ignite a splinter, and observe the combustion and
the smoke, if any. Try to collect some C in the same way as
before.
With plenty of O and high enough temperature, all the C is burned
to CO2, whether in gas, candle, or wood. CO2 is an invisible gas.
The porcelain, when held in the flame, cools the C below the
point at which it burns, called the kindling-point, and hence it
is deposited. The greater part of smoke is unburned carbon.
Experiment 28.--Hold an inverted dry t.t. or receiver over the
flame of a burning candle, and look for any moisture (H2O). What
two elements are shown by these experiments to exist in the
candle? The same two are found in wood and in gas. Experiment
29.--Put into a small Hessian crucible (Fig. 18) some pieces of
wood 2 or 3 cm long, cover with sand, and heat the crucible
strongly. When smoking stops, cool the crucible, remove the
contents, and examine the charcoal. The gases have been driven
off from the wood, and the greater part of what is left is C.
Experiment 30.--Put 1 g. of sugar into a porcelain crucible, and
heat till the sugar is black. C is left. See Experiment 5. Remove
the C with a strong solution of sodium hydrate (page 208).
41. Allotropic Forms.--Carbon is peculiar in that it occurs in at
least three allotropic, i.e. different, forms, all having
different properties. These are diamond, graphite, and amorphous
--not crystalline--carbon. The latter includes charcoal, lamp-
black, bone-black, gas carbon, coke, and mineral coal. All these
forms of C have one property in common; they burn in O at a high
temperature, forming CO2. This proves that each is the element C,
though it is often mixed with some impurities.
Allotropy, or allotropism, is the quality which an element often
has of appearing under various forms, with different properties.
The forms of C are a good illustration.
42. Diamond is the purest C; but even this in burning leaves a
little ash, showing that it is not quite pure. It is a rare
mineral, found in India, South Africa, and Brazil, and is the
hardest and most highly refractive to light of all minerals.
Boron is harder. [Footnote: B, not occurring free, is not a
mineral.] When heated in the electric arc, at very high
temperatures, diamond swells and turns black. 43. Graphite, or
Plumbago, is One of the Softest Minerals.--It is black and
infusible, and oxidizes only at very high temperatures, higher
than the diamond. It contains from 95 to 98 per cent C. Graphite
is found in the oldest rock formations, in the United States and
Siberia. It is artificially formed in the iron furnace. Graphite
is employed for crucibles where great heat is required, for a
lubricant, for making metal castings, and, mixed with clay, for
lead-pencils. It is often called black-lead.
44. Amorphous Carbon comprises the following varieties.
Charcoal is made by heating wood, for a long time, out of contact
with the air. The volatile gases are thus driven off from the
wood; what is left is C, and a small quantity of mineral matter
which remains as ash when the coal is burned.
45. Lamp-black is prepared as in Experiment 26, or by igniting
turpentine (C1OH16), naphtha, and various oils, and collecting
the C of the smoke. It is used for making printers' ink, India
ink, etc. A very pure variety is obtained from natural gas.
Bone-black, or animal charcoal, is obtained by distilling bones,
i.e. by heating them in retorts into which no air is admitted.
The C is the charred residue.
Gas Carbon is formed in the retorts of the gas-house. See page
182. It is used to some extent in electrical work.
46. Coke is the residue left after distilling soft coal. It is
tolerably pure carbon, with some ash and a little volatile
matter. It burns without flame. 47. Mineral Coal is fossilized
wood or other vegetable matter. Millions of years ago trees and
other vegetation covered the earth as they do to-day. In certain
places they slowly sank, together with the land, into the
interior of the earth, were covered with sand, rock, and water,
and heated from the earth's interior. A slow distillation took
place, which drove off some of the gases, and converted vegetable
matter into coal. All the coal dug from the earth represents
vegetable life of a former period. Millions of years were
required for the transformation; but the same change is in
progress now, where peat beds are forming from turf.
Coal is found in all countries, the largest beds being in the
United States. From the nature of its formation, coal varies much
in purity.
Anthracite, or hard coal, is purest in carbon, some varieties
having from 90 to 95 per cent. This represents most complete
distillation in the earth; i.e. the gases have mostly been driven
off. It is much used in New England.
48. Bituminous, or soft coal, crocks the hands, and burns rapidly
with much flame and smoke. The greater part of the coal in the
earth is bituminous. It represents incomplete distillation.
Hence, by artificially distilling it, illuminating gas is made.
See page 180. It is far less pure C than anthracite.
49. Cannel Coal is a variety of bituminous coal which can be
ignited like a candle. This is because so many of the gases are
still left, and it shows cannel to be less pure C than bituminous
coal.
50. Lignite, Peat, Turf, etc., are still less pure varieties of
C. Construct a table of the naturally occurring forms of this
element, in the order of their purity. Carbon forms the basis of
all vegetable and animal life; it is found in many rocks, mineral
oils, asphaltum, natural gas, and in the air as CO2.
51. C a Reducing Agent.
Experiment 31.--Put into a small ignition-tube a mixture of 4 or
5 g. of powdered copper oxide (CuO), with half its bulk of
powdered charcoal. Heat strongly for ten or fifteen minutes.
Examine the contents for metallic copper. With which element of
CuO has C united? The reaction may be written: Cu0 + C = CO + Cu.
Complete and explain.
A Reducing, or Deoxidizing, Agent is a substance which takes away
oxygen from a compound. C is the most common and important
reducing agent, being used for this purpose in smelting iron and
other ores, making water-gas, etc.
An Oxidizing Agent is a substance that gives up its O to a
reducing agent. What oxidizing agent in the above experiment?
52. C a Decolorizer.
Experiment 32.--Put 3 or 4 g. of bone-black into a receiver, and
add 10 or 15 cc.of cochineal solution. Shake this thoroughly,
covering the bottle with the hand. Then pour the whole on a
filter paper, and examine the filtrate. If all the color is not
removed, filter again. What property of C is shown by this
experiment? Any other coloring solution may be tried.
The decolorizing power of charcoal is an important
characteristic. Animal charcoal is used in large quantities for
decolorizing sugar. The coloring matter is taken out mechanically
by the C, there being no chemical action. 53. C a Disinfectant.
Experiment 33.--Repeat the previous experiment, adding a solution
of H2S3 i.e. hydrogen sulphide, in water, instead of cochineal
solution. See page 120. Note whether the bad odor is removed. If
not, repeat.
Charcoal has the property of absorbing large quantities of many
gases. Ill-smelling and noxious gases are condensed in the pores
of the C; O is taken in at the same time from the air, and these
gases are there oxidized and rendered odorless and harmless. For
this reason charcoal is much used in hospitals and sick-rooms, as
a disinfectant. This property of condensing O, as well as other
gases, is shown in the experiment below.
54. C an Absorber of Gases and a Retainer of Heat.
Experiment 34.--Put a piece of phosphorus of the size of a pea,
and well dried, on a thick paper. Cover it well with bone-black,
and look for combustion after a while. O has been condensed from
the air, absorbed by the C, and thus communicated to the P. Burn
all the P at last.
VALENCE.
55. The Symbols NaCl and MgCl2 differ in two ways.--What are
they? Let us see why the atom of Mg unites with two Cl atoms,
while that of Na takes but one. If the atoms of two elements
attract each other, there must be either a general attraction all
over their surfaces, or else some one or more points of
attraction. Suppose the latter to be true, each atom must have
one or more poles or bonds of attraction, like the poles of a
magnet. Different elements differ in their number of bonds. Na
has one, which may be written graphically Na-; Cl has one, -Cl.
When Na unites with Cl, the bonds of each element balance, as
follows: Na-Cl. The element Mg, however, has two such bonds, as
Mg= or -Mg-. When Mg unites with Cl, in order to balance, or
saturate, the bonds, it is evident that two atoms of Cl must be
used, as Cl-Mg-Cl, or MgCl2.
A compound or an element, in order to exist, must have no free
bonds. In organic chemistry the exceptions to this rule are very
numerous, and, in fact, we do not know that atoms have bonds at
all; but we can best explain the phenomena by supposing them, and
for a general statement we may say that there must be no free
bonds. In binaries the bonds of each element must balance.
56. The Valence, Quantivalence, of an Element is its Combining
Power Measured by Bonds.--H, having the least number of bonds,
one, is taken as the unit. Valence has always to be taken into
account in writing the symbol of a compound. It is often written
above and after the elements [i.e. written like an exponent], as
K^I, Mg^II.
An element having a valence of one is a monad; of two, a dyad;
three, a triad; four, tetrad; five, pentad; six, hexad, etc. It
is also said to be monovalent, di- or bivalent, etc. This theory
of bonds shows why an atom cannot exist alone. It would have free
or unused bonds, and hence must combine with its fellow to form a
molecule, in case of an element as well as in that of a compound.
This is illustrated by these graphic symbols in which there are
no free bonds: H-H, O=O, N[3-bond symbol]N, C[4-bond symbol]C. A
graphic symbol shows apparent molecular structure.
After all, how do we know that there are twice as many Cl atoms
in the chloride of magnesium as in that of sodium? The compounds
have been analyzed over and over again, and have been found to
correspond to the symbols MgCl2 and NaCl. This will be better
understood after studying the chapter on atomic weights. In
writing the symbol for the union of H with O, if we take an atom
of each, the bonds do not balance, H-=O, the former having one;
the latter, two. Evidently two atoms of H are needed, as H-O-H,
or
H
= O , or H2O. In the union of Zn and O, each has two bonds;
H
hence they unite atom with atom, Zn = O, or ZnO.
Write the grapbic and the common symbols for the union of H^I and
Cl^I; of K^I and Br^I; Ag^I and O^II; Na^I and S^II; H^I and
P^III. Study valences. It will be seen that some elements have a
variable quantivalence. Sn has either 2 or 4; P has 3 or 5. It
usually varies by two for a given element, as though a pair of
bonds sometimes saturated each other;. e.g. =Sn=, a quantivalence
of 4, and |Sn=, a quantivalence of 2. There are, therefore, two
oxides of tin, SnO and SnO2, or Sn=O and O=Sn=O. Write symbols
for the two chlorides of tin; two oxides of P; two oxides of
arsenic.
The chlorides of iron are FeCl2 and Fe2Cl6. In the latter, it
might be supposed that the quantivalence of Fe is 3, but the
graphic symbol shows it to be 4. It is called a pseudo-triad, or
false triad. Cr and Al are also pseudo-triads.
Cl Cl | | Cl--Fe--Fe--Cl | | Cl Cl
Write formulae for two oxides of iron; the oxide of Al.
57. A Radical is a Group of Elements which has no separate
existence, but enters into combination like a single atom; e.g.
(NO3) in the compounds HNO3 or KNO3; (SO4) in H2SO4. In HNO3 the
radical has a valence of 1, to balance that of H, H-NO3). In
H2SO4, what is the valence of (SO4)? Give it in each of these
radicals, noting first that of the first element: K(NO3),
Na2(SO4), Na2(CO3), K(ClO3), H3(PO4), Ca3(PO4)2, Na4(SiO4).
Suppose we wish to know the symbol for calcium phosphate. Ca and
PO4 are the two parts. In H3(PO4) the radical is a triad, to
balance H3. Ca is a dyad, Ca==(P04). The least common multiple of
the bonds (2 and 3) is 6, which, divided by 2 (no. Ca bonds),
gives 3 (no. Ca atoms to be taken). 6 / 3 (no. (PO4) bonds) gives
2 (no. PO4 radicals to be taken). Hence the symbol Ca3(P04)2.
Verify this by writing graphically.
Write symbols for the union of Mg and (SO4), Na and (PO4), Zn and
(NO3), K and (NO3), K and (SO4), Mg and (PO4), Fe and (SO4) (both
valences of Fe), Fe and (NO3), taking the valences of the
radicals from HNO3, H2SO4, H3PO4.
Chapter XII.
ELECTRO-CHEMICAL RELATION OF ELEMENTS.
58. Examine untarnished pieces of iron, silver, nickel, lead,
etc.; also quartz, resin, silk, wood, paper. Notice that from the
first four light is reflected in a different way from that of the
others. This property of reflecting light is known as luster.
Metals have a metallic luster which is peculiar to themselves;
and this, for the present, may be regarded as their chief
characteristic. Are they at the positive or negative end of the
list? See page 43. How is it with the non-metals? This
arrangement has a significance in chemistry which we must now
examine. The three appended experiments show how one metal can be
withdrawn from solution by a second, this second by a third, the
third by a fourth, and so on. For expedition, three pupils can
work together for the three following experiments, each doing
one, and examining the results of the others.
59. Deposition of Silver.
Experiment 35.--Put a ten-cent Ag coin into an evaporating-dish,
and pour over it a mixture of 5 cc. HNO3 and 10 cc. H2O. Warm
till all, or nearly all, the Ag dissolves. Remove the lamp. 3 Ag
+ 4 HNO3 = 3 AgNO3 + 2 H2O + NO. Then add 10 cc. H2O, and at once
put in a short piece of Cu wire, or a cent. Leave till quite a
deposit appears, then pour off the liquid, wash the deposit
thoroughly, and remove it from the coin. See whether the metal
resembles Ag. 2 AgNO3 + Cu =?60. Deposition of Copper.
Experiment 36.--Dissolve a cent or some Cu turnings in dilute
HNO3, as in Experiment 35, and dilute the solution. 3 Cu + 8 HN09
- 3 Cu (NOA+4 H2O+2 NO.)
Then put in a clean strip of Pb, and set aside as before,
examining the deposit finally. Cu(NO3), + Pb - ?
61. Deposition of Lead.
Experiment 37.--Perform this experiment in the same manner as the
two previous ones, dissolving a small piece of Pb, and using a
strip of Zn to precipitate the Pb. 3 Pb + 8 HNO3 - 3 Pb (NO4)2 +
4 Ha0 + 2 NO. Pb (NO3) 2 + Zn = ? h.
62. Explanation. -These experiments show that Cu will replace Ag
in a solution of AgNO3, that Pb will replace and deposit Cu from
a similar compound, and that Zn will deposit Pb in the same way.
They show that the affinity of Zn for (NO3) is stronger than
either Ag, Cu, or Pb. We. express this affinity by saying that Zn
is the most positive of the four metals, while Ag is the most
nega- tive. Cu is positive to Ag, but negative to Pb and Zn.
Which of the four elements are positive to Pb, and which
negative? Mg would withdraw Zn from a similar solution, and be in
its turn withdrawn by Na. The table on page 43 is founded on this
relation. A given element is positive to every element above it
in the list, and negative to all below it.
Metals are usually classed as positive, non-metals as negative.
Each in union with O and 1=I gives rise to a very important class
of compounds,=--the negative to acids, the positive to bases.
In the following, note whether the positive or the negative
element is written first:--HCl, Na20,-As2S3, -MgBr2, Ag2S. Na2SO4
is made up of two parts, Na2 being positive, the radical SO4
negative. Like elements, radicals are either positive or
negative. In the following, separate the positive element from
the negative radical by a vertical line: Na2CO3, NaNO3, ZnSO4,
KClO3.
The most common positive radical is NH4, ammonium, as in NH4Cl.
It always deports itself as a metal. The commonest radical is the
negative OH, called hydroxyl, from hydrogen- oxygen. Take away H
from the symbol of water, H-O-H, and hydroxyl --(OH) with one
free bond is left. If an element takes the place of H, i.e.
unites with OH, the compound is called a hydrate. KOH is
potassium hydrate. Name NaOH, Ca(OH)2, NH4OH, Zn(OH)2, Al2(OH)6.
Is the first part of each symbol above positive or negative?
H has an intermediate place in the list. It is a constituent of
both acids and bases, and of the neutral substance, water.
ORDER.
--
Negative or Non-Metallic Elements.
Acid-forming with H(usually OH).
Oxygen
Sulphur
Nitrogen
Fluorine
Chlorine
Bromine
Iodine
Phosphorus
Arsenic
Carbon
Silicon
Hydrogen
Positive or Metallic Elements.
Base-forming with OH.
Gold
Platinum
Mercury
Silver
Copper
Tin
Lead
Iron
Zinc
Aluminium
Magnesium
Calcium
Sodium
Potassium
CHAPTER XIII.
ELECTROLYSIS.
The following experiment is to be performed only by the teacher,
but pupils should make drawings and explain.
63. Decomposition of Water.
Experiment 38.--Arrange "in series" two or more cells of a Bunsen
battery (Physics, page 164), [References are made in this book to
Gage's Introduction to Physical Science.] and attach the terminal
wires to an electrolytic apparatus (Fig. 19) filled with water
made slightly acid with H2SO4. Construct a diagram of the
apparatus, marking the Zn in the liquid +, since it is positive,
and the C, or other element, -. Mark the electrode attached to
the Zn -, and that attached to the C +; positive electricity at
one end of a body commonly implies negative at the other.
Opposites attract, while like electricities repel each other.
These analogies will aid the memory. At the + electrode is the -
element of H2O, and at the - electrode the + element. Note, page
43, whether H or O is positive with reference to the other, and
write the symbol for each at the proper electrode. Compare the
diagram with the apparatus, to verify your conclusion. Why does
gas collect twice as fast at one electrode as at the other? What
does this prove of the composition of water? When filled, test
the gases in each tube, for O and H, with a burning stick.
Electrical analysis is called electrolysis.
If a solution of NaCl be electrolyzed, which element will go to
the + pole? Which, if the salt were K2SO4? Explain these
reactions in the electrolysis of that salt. K2SO4 = K2 + S03 + O.
SO4 is unstable, and breaks up into SO3 and O. Both K and SO3
have great affinity for water. K2 + 2 H2O = 2 KOH + H2. S03 + H2O
= H2SO4.
The base KOH would be found at the - electrode, and the acid
H2SO4 at the + electrode.
The positive portion, K, uniting with H2O forms a base; the
negative part, S03, with H2O forms an acid. Of what does this
show a salt to be composed?
64. Conclusions.--These experiments show (1) that at the +
electrode there always appears the negative element, or radical,
of the compound, and at the - electrode the positive element; (2)
that these elements unite with those of water, to make, in the
former case, acids, in the latter, bases; (3) that acids and
bases differ as negative and positive elements differ, each being
united with O and H, and yet producing compounds of a directly
opposite character; (4) that salts are really compounded of acids
and bases. This explains why salts are usually inactive and
neutral in character, while acids and bases are active agents.
Thus we see why the most positive or the most negative elements
in general have the strongest affinities, while those
intermediate in the list are inactive, and have weak affinities;
why alloys of the metals are weak compounds; why a neutral
substance, like water, has such a weak affinity for the salts
which it holds in solution; and why an aqueous solution is
regarded as a mechanical mixture rather than a chemical compound.
In this view, the division line between chemistry and physics is
not a distinct one. These will be better understood after
studying the chapters on acids, bases and salts.
Chapter XIV.
UNION BY VOLUME.
66. Avogadro's Law of Gases.--Equal volumes of all gases, the
temperature and pressure being the same, have the same number of
molecules. This law is the foundation of modern chemistry. A
cubic centimeter of O has as many molecules as a cubic centimeter
of H, a liter of N the same number as a liter of steam, under
similar conditions. Compare the number of molecules in 5 l. of
N2O with that in 10 l. Cl. 7 cc. vapor of I to 6 cc. vapor of S.
The half-molecules of two gases have, of course, the same
relation to each other, and in elements the half-molecule is
usually the atom.
The molecular volumes--molecules and the surrounding space--of
all gases must therefore be equal, as must the half-volumes.
Notice that this law applies only to gases, not to liquids or
solids. Let us apply it to the experiment for the electrolysis of
water. In this we found twice as much H by volume as O.
Evidently, then, steam has twice as many molecules of H as of O,
and twice as many half-molecules, or atoms. If the molecule has
one atom of O, it must have two of H, and the formula will be
H2O.
Suppose we reverse the process and synthesize steam, which can be
done by passing an electric spark through a mixture of H and O in
a eudiometer over mercury; we should need to take twice as much H
as O. Now when 2 cc. of H combine thus with 1 cc. of O, only 2
cc.of steam are produced. Three volumes are condensed into two
volumes, and of course three molecular volumes into two, three
atomic volumes into two. This may be written as follows:--
H + H + O = H2O.
This is a condensation of one-third.
If 2 l. of chlorhydric acid gas be analyzed, there will result 1
l. of H and 1 l. of Cl. The same relation exists between the
molecules and the atoms, and the reaction is:--
HCl = H + Cl.
Reverse the process, and 1 l. of H unites with 1 l. of Cl to
produce 2 l. of the acid gas; there is no condensation, and the
symbol is HCl. In seven volumes HCl how many of each constituent?
The combination of two volumes of H with one volume of S is found
to produce two volumes of hydrogen sulphide. Therefore two atoms
of H combine with one of S to form a molecule whose symbol is
H2S.
H + H + S = H2S.
What is the condensation in this case?
PROBLEMS.
(1) How many liters of S will it take to unite with 4 l. of H?
How much H2S will be formed?
(2) How many liters of H will it take to combine with 5 l. of S?
How much H2S results?
(3) In 6 l. H2S how many liters H, and how much S? Prove.
(4) In four volumes H2S how many volumes of each constituent?
(5) If three volumes of H be mixed with two volumes of S, so as
to make H2S, how much will be formed? How much of either element
will be left? An analysis of 2 cc. of ammonia gives 1 cc. N and 3
cc. H. The symbol must then be NH3, the reaction,--
NH3 = N + H + H + H.
What condensation in the synthesis of NH3?
In 12 cc. NH3 how many cubic centimeters of each element? In 2
1/2 cc? How much H by volume is required to combine with nine
volumes of N? How many volumes of NH3 are produced?
In elements that have not been weighed in the gaseous state, as
C, the evidence of atomic volume is not direct, but we will
assume it. Thus two volumes of marsh gas would separate into one
of C and four of H. What is its symbol and supposed condensation?
Two volumes of alcohol vapor resolve into two of C, six of H, and
one of O. What is its symbol? its condensation?
The symbol itself of a compound will usually show what its
condensation is; e.g. HCl, HBr, HF, etc., have two atoms; hence
there will be no shrinkage. In H2O, SO2, CO2, the molecule has
three atoms condensed into the space of two, or one-third
shrinkage. In NH3 four volumes are crowded into the space of two,
a condensation of one-half.
P, As, Hg, Zn, have exceptional atomic volumes.
Chapter XV.
ACIDS AND BASES.
66. What Acids Are.
Experiment 39.--Pour a few drops of chlorhydric acid, HCl, into a
clean evaporating-dish. Add 5 cc. H2O, and stir. Touch a drop to
the tongue, noting the taste. Dip into it the end of a piece of
blue litmus paper, and record the result. Thoroughly wash the
dish, then pour in a few drops of nitric acid, HNO3, and 5 cc.
H2O, and stir. Taste, and test with blue litmus. Test in the same
way sulphuric acid, H2SO4. Name two characteristics of an acid.
In a vertical line write the formulae of the acids above. What
element is common to them all? Is the rest of the formula
positive or negative?
67. An Acid is a substance composed of H and a negative element
or radical. It has usually a sour taste, and turns blue litmus
red. Litmus is a vegetable extract obtained from lichens in
Southern Europe. Acids have the same action on many other
vegetable pigments. Are the following acid formulae, and why?
H2SO3, HBr, HNO2, H3PO3, H4SiO4. Most acids have O as well as H.
Complete the symbols for acids in the following list, and name
them, from the type given:--
HCl, chlorhydric acid. HN03, nitric acid.
?Br, ? ?Cl? ?
?I, ? ?Br? ?
?F, ? ?I? ?
H3PO4, phosphoric acid. H3PO3, phosphorous acid.
?As? ? ?As? ?
Complete these equations:--
H2SO3 - H2O = ? | 2 HN03 - H2O = ?
H2SO4 - H2O = ? | 2 HNO2 - H2O = ?
H2CO3 - H2O = ? | 2 H3AsO4 - 3 H2O = ?
Are the products in each case metallic or non-metallic oxides?
They are called anhydrides. Notice that each is formed by the
withdrawal of water from an acid. Reverse the equations; as, SO3
+ H2O = ?
68. An Anhydride is what remains after water has been removed
from an acid; or, it is the oxide of a non- metallic element,
which, united with water, forms an acid. SO2 is sulphurous
anhydride, SO2 sulphuric anhydride, the ending ic meaning more O,
or negative element, than ous. Name the others above.
Anhydrides were formerly called acids,--anhydrous acids, in
distinction from hydrated ones, as CO2 even now is often called
carbonic acid.
Experiment 40.--Hold a piece of wet blue litmus paper in the
fumes of SO2, and note the acid test. Try the same with dry
litmus paper.
Experiment 41.--Burn a little S in a receiver of air containing
10 cc. H2O, and loosely covered, as in the O experiment. Then
shake to dissolve the SO2. H2O + SO2 = H2SO3. Apply test paper.
69. Naming Acids.--Compare formulae H2SO3 and H2SO4. Of two acids
having the same elements, the name of the one with least O, or
negative element, ends in ous, the other in ic. H2SO3 is
sulphurous acid, H2SO4, sulphuric acid. Name H3PO4 and H3PO3;
H3AsO3 and H3ASO4; HNO2 and HNO3.
If there are more than two acids in a series, the prefixes hypo,
less, and per, more, are used. The following is such a series:
HClO, HClO2, HClO3, HClO4.
HClO3 is chloric acid; HClO2, chlorous; HClO, hypochlorous; HClO4
perchloric. Hypo means less of the negative element than ous; per
means more of the negative element than ic. Name: H3PO4 (ic),
H3PO3, H3PO2. Also HBrO (HBrO2 does not exist), HBrO3 (ic),
HBrO4.
What are the three most negative elements? Note their occurrence
in the three strongest and most common acids. Hereafter note the
names and symbols of all the acids you see.
70. What Bases Are.
Experiment 42.--Put a few drops of NH4OH into an evaporating-
dish. Add 5 cc. H2O, and stir. Taste a drop. Dip into it a piece
of red litmus paper, noting the effect. Cleanse the dish, and
treat in the same way a few drops NaOH solution, recording the
result. Do the same with KOH. Acid stains on the clothing, with
the exception of those made by HNO3, maybe removed by NH4OH.
H2SO4, however, rapidly destroys the fiber of the cloth.
Name two characteristics of a base. In the formulae of those
bases, what two common elements? Name the radical. Compare those
symbols with the symbol for water, HOH. Is (OH) positive or
negative? Is the other part of each formula positive or negative?
What are two constituents, then, of a base? Bases are called
hydrates. Write in a vertical line five positive elements. Note
the valence of each, and complete the formula for its base. Affix
the names. Can you see any reason why the three bases above given
are the strongest?
Taking the valences of Cr and Fe, write symbols for two sets of
hydrates, and name them. Try to recognize and name every base
hereafter met with.
A Base is a substance which is composed of a metal, or positive
radical, and OH. It generally turns red litmus blue, and often
has an acrid taste.
An Alkali is a base which is readily soluble in water. The three
principal alkalies are NH4OH, KOH, and NaOH.
Alkali Metals are those which form alkalies. Name three.
An Alkaline Reaction is the turning of red litmus blue.
An Acid Reaction is the turning of blue litmus red.
Experiment 43.--Pour 5 cc. of a solution of litmus in water, into
a clean t.t. or small beaker. Pour 2 or 3 cc. of HCl into an
evaporating-dish, and the same quantity of NH4OH into another
dish. Take a drop of the HCl on a stirring-rod and stir the
litmus solution with it. Note the acid reaction. Clean the rod,
and with it take a drop (or more if necessary) of NH4OH, and add
this to the red litmus solution, noting the alkaline reaction.
Experiment in the same way with the two other principal acids and
the two other alkalies.
Litmus paper is commonly used to test these reactions, and
hereafter whenever the term LITMUS is employed in that sense, the
test-paper should be understood. This paper can be prepared by
dipping unglazed paper into a strong aqueous solution of
litmus.
CHAPTER XVI.
SALTS.
71. Acids and Bases are usually Opposite in Character.--When two
forces act in opposition they tend to neutralize each other. We
may see an analogy to this in the union of the two opposite
classes of compounds, acids and bases, to form salts.
72. Neutralization.
Experiment 44.--Put into an evaporating-dish 5 cc. of NaOH
solution. Add HCl to this from a t.t., a few drops at a time,
stirring the mixture with a glass rod (Fig. 20), and testing it
with litmus paper, until the liquid is neutral, i.e. will not
turn the test paper from blue to red, or red to blue. Test with
both colors. If it turns blue to red, too much acid has been
added; if red to blue, too much base. When it is very nearly
neutral, add the reagent, HCl or NaOH, a drop at a time with the
stirring-rod. It must be absolutely neutral to both colors.
Evaporate the water by heating the dish over asbestus paper, wire
gauze, or sand, in an iron plate (Fig. 21) till the residue
becomes dry and white. Cool the residue, taste, and name it. The
equation is: HCl + NaOH = NaCl + HOH or H2O. Note which elements,
positive or negative, change places. Why was the liquid boiled?
The residue is a type of a large class of compounds, called
salts.
(Fig. 20) (Fig. 21)
Experiment 45. -- Experiment in the same way with KOH solution
and H2SO4, applying the same tests. H2SO4 + 2 KOH = K2SO4 + 2
HOH. What is the solid product?
Experiment 46.--Neutralize NH4OH with HNO3, evaporate, apply the
tests, and write the equation. Write equations for the
combination of NaOH and H2SO4; NaOH and HNO3; KOH and HCl; KOH
and HNO3; NH4OH and HCl; NH4OH and H2SO4. Describe the experiment
represented by each equation, and be sure you can perform it if
asked to do so. What is the usual action of a salt on litmus? How
is a salt made? What else is formed at the same time? Have all
salts a saline taste? Does every salt contain a positive element
or radical? A negative?
73. A Salt is the product of the union of a positive and a
negative element or radical; it may be made by mixing a base and
an acid.
The salt KI represents what acid? What base, or hydrate? Write
the equation for making KI from its acid and base. Describe the
experiment in full. Classify, as to acids, bases, or salts: KBr,
Fe(OH)2, HI, NaBr, HNO2, Al2(OH)6, KClO3, HClO3, H2S, K2S, H2S03,
K2SO4, Ca(OH)2, CaCO3, NaBr03, CaSO4, H2CO3, K2CO3, Cu(OH)2,
Cu(NO3)2, PbSO4, H3P04, Na2P04. In the SALTS above, draw a light
vertical line, separating the positive from the negative part of
the symbol. Now state what acid each represents. What base. Write
the reaction in the preparation of each salt above from its acid
and base; then state the experiment for producing it.
74. Naming Salts.--(NO3) is the nitrate radical; KNO3 is
potassium nitrate. From what acid? (NO2) is the nitrite radical;
KN02 is potassium nitrite. From what acid? Note that the endings
of the acids are OUS and IC; also that the names of their salts
end in ITE and ATE. From which acid--IC or OUS--is the salt
ending in ATE derived? That ending in ITE?
Name these salts, the acids from which they are derived, and the
endings of both acids and salts: NaNO3, NaNO2, K2SO4, K2SO3,
CaSO4, CaSO3, KClO3, KClO2, KClO, KClO4 (use prefixes HYPO and
PER, as with acids), Ca3(PO4)2, Ca3(P03)2, CuSO4, CuSO3, AgNO3,
Cu(NO3)2. FeS, FeS2, are respectively FERROUS SULPHIDE and FERRIC
SULPHIDE. Name: HgCl, HgCl2, FeCl2, Fe2Cl6, FeSO4, Fe2(SO4)3.75.
Acid Salts.--Write symbols for nitric, sulphuric, phosphoric
acids. How many H atoms in each? Replace all the H in the symbol
of each with Na, and name the products. Again, in sulphuric acid
replace one atom of H with Na; then in phosphoric replace first
one, then two, and finally three H atoms with Na. HNaSO4 is
hydrogen sodium sulphate; HNa2P04 is hydrogen di-sodium
phosphate. Name the other salts symbolized. Name HNaNH4P04.
Though these products are all salts, some contain replaceable H,
and are called acid salts. Those which have all the H replaced by
a metal are normal salts. Name and classify, as to normal or acid
salts: Na2CO3, HNaCO3, K2SO4, HKSO4, (NH4)2SO4, HNH4SO4, Na3P04,
HNa2P04, H2NaP04.
The BASICITY of an acid is determined by the number of
replaceable H atoms in its molecule. It is called MONOBASIC if it
has one; DIBASIC if two; TRI- if three, etc. Note the basicity of
each acid named above. How many possible salts of H2SO4 with Na?
Of H3P04 with Na? Which are normal and which acid? What is the
basicity of H4Si04?
Some normal, as well as acid, salts change litmus. Na2CO3,
representing a strong base and a weak acid, turns it blue. There
are other modes of obtaining salts, but this is the only one
which we sball consider.
76. Salts Occur Abundantly in Nature, such as NaCl, MgSO4, CaCO3.
Acids and bases are found in small quantities only. Why is this?
Why are there not springs of H2SO4 and NH4OH? We have seen that
acids and bases are extremely active, have opposite characters,
and combine to form relatively inactive salts. If they existed in
the free state, they would soon combine by reason of their strong
affinities. This is what in all ages of the world has taken
place, and this is why salts are common, acids and bases rare.
Active agents rarely exist in the free state in large quantities.
Oxygen seems to be an exception, but this is because there is a
superabundance of it. While vast quantities are locked up in
compounds in rocks, water, and salts of the earth, much remains
with which there is nothing to combine.
CHAPTER XVII.
CHLORHYDRIC ACID.
77. We have seen that salts are made by the union of acids and
bases. Can these last be obtained from salts?
78. Preparation of HCl.
Experiment 47.--Into a flask put 10 g. coarse NaCl, and add 20
cc. H2SO4. Connect with Woulff bottles [Woulff bottles may be
made by fitting to wide-mouthed bottles corks with three holes,
through which pass two delivery tubes, and a central safety tube
dipping into the liquid, as in Figures 22 and 23.] partly filled
with water, as in Figure 22. One bottle is enough to collect the
HCl; but in that case it is less pure, since some H2SO4 and other
impurities are carried over. Several may be connected, as in
Figure 23. The water in the first bottle must be nearly saturated
before much gas will pass into the second. Heat the mixture 15 or
20 minutes, not very strongly, to prevent too much foaming.
Notice any current in the first bottle. NaCl + H2SO4 = HNaSO4 +
HCl. Intense heat would have given: 2NaCl + H2SO4 = Na2SO4 +
2HCl. Compare these equations with those for HNO3. In which
equation above is H2SO4 used most economically? Both reactions
take place when HCl is made on the large scale.
(Fig. 22)
79. Tests. Experiment 48.--(1) Test with litmus the liquid in
each Woulffbottle. (2) Put a piece of Zn into a t.t. and cover it
with liquid from the first bottle. Write the reaction, and test
the gas. (3) To 2 cc.solution AgNO3 in a t.t. add 2 cc.of the
acid. Describe, and write the reaction. Is AgCl soluble in water?
(4) Into a t.t. pour 5 cc.Pb(NO3)2 solution, and add the same
amount of prepared acid. Give the description and the reaction.
(5) In the same way test the acid with Hg2(NO3)2 solution, giving
the reaction. (6) Drake a little HCl in a t.t., and bring the gas
escaping from the d.t. in contact with a burning stick. Does it
support the combustion of C? (7) Hold a piece of dry litmus paper
against it. [figure 23] (8) Hold it over 2 cc.of NH4OH in an
evaporating-dish. Describe, name the product, and write the
reaction. (3), (4), (5), (8), are characteristic tests for this
acid.
80. Chlorhydric, Hydrochloric or Muriatic, Acid is a Gas.--As
used, it is dissolved, in water, for which it has great affinity.
Water will hold, according to temperature, from 400 to 500 times
its volume of HCl. Hundreds of thousands of tons of the acid are
annually made, mostly in Europe, as a bye-product in Na2CO3
manufacture. The gas is passed into towers through which a spray
of water falls; this absorbs it. The yellow color in most
commercial HCl indicates impurities, some of which are Fe, S, As,
and organic matter. As, S, etc., come from the pyrites used in
making H2SO4. Chemically pure (C.P.) acid is freed from these,
and is without color. The gas may be dried by passing it through
a glass tube holding CaCl2 (Fig. 16) and collecting it over
mercury.
The muriatic acid of commerce consists of about two- thirds water
by weight. HCl can also be made by direct union of its
constituents.81. Uses.--HCl is used to make Cl, and also
bleaching- powder. Its use as a reagent in the laboratory is
illustrated by the following experiment:-- Experiment 49.--Put
into a t.t. 2 cc. AgNO3 solution, add 5 cc. H2O, then add slowly
HCl so long as a ppt. (precipitate) is formed. This ppt. is AgCl.
Now in another t.t. put 2 cc. Cu(NO3)2, solution, add 5 cc. H2O,
then a little HCl. No ppt. is formed. Now if a solution of AgNO3
and a solution of Cu(NO3)2 were mixed, and HCl added, it is
evident that the silver would be precipitated as chloride of
silver, while the copper would remain in solution. If now this be
filtered, the silver will remain on the filter paper, while in
the filtrate will be the copper. Thus we shall have performed an
analysis, or separated one metal from another. Perform it. Note,
however, that any soluble chloride, as NaCl, would produce the
same result as HCl.
BROMHYDRIC AND IODIHYDRIC ACIDS.
82. NaCl, being the most abundant compound of Cl, is the source
of commercial HCl. KCl treated in the same way would give a like
product. Theoretically HBr and HI might be made in the same way
from NaBr and NaI, but the affinity of H for Br and I is weak,
and the acids separate into their elements, when thus prepared.
83. To make HI.
Experiment 50.--Drop into a t.t. three or four crystals of I, and
add 10 cc. H2O. Hold in the water the end of a d.t. from which
H2S gas is escaping. Observe any deposit, and write the reaction.
FLUORHYDRIC ACID.
84. Preparation and Action.
Experiment 51.--Put 3 or 4 g. powdered CaF2, i.e. fluor spar or
fluorite, into a shallow lead tray, e.g. 4x5 cm, and pour over it
4 or 5 cc. H2SO4. A piece of glass large enough to cover this
should previously be warmed and covered on one side with a very
thin coat of beeswax. To distribute itevenly, warm the other side
of the glass over a flame. When cool, scratch a design (Fig. 24)
through the wax with a sharp metallic point. Lay the glass, film
side down, over the lead tray. Warm this five minutes or more by
placing it high over a small flame (Fig. 25) to avoid melting the
wax. Do not inhale the fumes. Take away the lamp, and leave the
tray and glass where it is not cold, for half an hour or more.
Then remove the wax and clean the glass with naphtha or benzine.
Look for the etching.
Two things should have occurred: (1) the generation of HF. Write
the equation for it. (2) Its etching action on glass. In this
last process HF acts on SiO2 of the glass, forming H2O and SiF4.
Why cannot HF be kept in glass bottles?
A dilute solution of HF, which is a gas, may be kept in gutta
percha bottles, the anhydrous acid in platinum only; but for the
most part, it is used as soon as made, its chief use being to
etch designs on glass-ware. Glass is also often etched by a blast
of sand (SiO2).
Notice the absence of O in the acids HF, HCI, HBr, HI, and that
each is a gas. HF is the only acid that will dissolve or act
appreciably on glass.
Chapter XVIII.
NITRIC ACID.
85. Preparation. Experiment 52.--To 10 g. KNO3 or NaNO3, in a
flask, add 15 cc. H2SO4. Securely fasten the cork of the d.t., as
HNO3 is likely to loosen it, and pass the other end to the bottom
of a t.t. held deep in a bottle of water (Fig. 26). Apply heat,
and collect 4 or 5 cc.of the liquid. The usual reaction is: KNO3
+ H2SO4 = HKSO4 + HNO3. With greater heat, 2 KNO3 + H2SO4 = K2SO4
+ 2HNO3. Which is most economical of KNO3? Of H2SO4? Instead of a
flask, a t.t. may be used if desired (Fig. 27).
86. Properties and Tests.
Experiment 53.--(1) Note the color of the prepared liquid. (2)
Put a drop on the finger; then wash it off at once. (3) Dip a
quill or piece of white silk into it; then wash off the acid.
What color is imparted to animal substances? (4) Add a little to
a few bits of Cu turnings, or to a Cu coin. Write the equation.
(5) To 2 cc.indigo solution, add 2 cc. HNO3. State the leading
properties of HNO3, from these tests.
87. Chemically Pure HNO3 is a Colorless Liquid.-- The yellow
color of that prepared in Experiment 52 is due to liquid NO2
dissolved in it. It is then called fuming HNO3, and is very
strong. NO2 is formed at a high temperature.
Commercial or ordinary HNO3, is made from NaNO3, this being
cheaper than KNO3; it is about half water.
88. Uses. HNO3 is the basis of many nitrates, as AgNO3, used for
photography, Ba(NO3)2 and Sr(NO3)2 for fire-works, and others for
dyeing and printing calico; it is employed in making aqua regia,
sulphuric acid, nitro-glycerine, gun-cotton, aniline colors,
zylonite, etc.
Enough experiments have been performed to answer the question
whether some acids can be prepared from their salts. H2SO4 is not
so made, because no acid is strong enough to act on its salts. In
making HCl, HNO3, etc., sulphuric acid was used, being the
strongest.
AQUA REGIA.
89. Preparation and Action. Experiment 54.--Into a t.t. put 2 cc.
HNO3, and 14 qcm. of either Au leaf or Pt. Warm in a flame. If
the metal is pure, no action takes place. Into another tube put 6
cc. HCl and add a similar leaf. Heat this also. There should be
no action. Pour the contents of one t.t. into the other. Note the
effect. Which is stronger, one of the acids, or the combination
of the two? Note the odor. It is that of Cl. 3HCl + HNO3 = NOCl +
2H2O + Cl2. This reaction is approximate only. The strength is
owing to nascent chlorine, which unites with Au. Au + 3Cl =
AuCl3. If Pt be used, PtCl4 is produced. No other acid except
nitro-hydrochloric will dissolve Au or Pt; hence the ancients
called it aqua regia, or king of liquids. It must be made as
wanted, since it cannot be kept and retain its strength.
CHAPTER XIX.
SULPHURIC ACID.
90. Preparation.
Experiment 55.--Having fitted a cork with four or five
perforations to a large t.t., pass a d.t. from three of these to
three smaller t.t., leaving the others open to the air, as in
Figure 28. Into one t.t. put 5 cc. H2O, into another 5 g. Cu
turnings and 10 cc. H2SO4, into the third 5 g. Cu turnings and 10
cc. dilute HNO3, half water. Hang on a ring stand, and slowly
heat the tubes containing H2O and H2SO4. Notice the fumes that
pass into the large t.t.
Trace out and apply to Figure 28 these reactions:--
(1) Cu + 2 H2SO4 = CuSO4 + 2 H2O + SO2.
(2) 3 Cu + 8 HNO3 = 3 Cu(NO3)2+ 4 H2O + 2 NO.
(3) NO + O = NO2.
(4) SO2 + H2O + NO2 =H2SO4 + NO.
(4) comes from combining the gaseous products in (1), (2), (3).
In (3), NO takes an atom of O from the air, becoming NO2, and at
once gives it up, to the H2SO3 (H2O + SO2), making H2SO4, and
again goes through the same operation of taking up O and passing
it along. NO is thus called a carrier of O. It is a reducing
agent, while NO2 is an oxidizing agent. This is a continuous
process, and very important, since it changes useless H2SO3 into
valuable H2SO4. If exposed to the air, H2SO3 would very slowly
take up O and become H2SO4.
Instead of the last experiment, this may be employed if
preferred: Burn a little S in a receiver. Put into an
evaporating-dish, 5 cc. HNO3, and dip a paper or piece of cloth
into it. Hang the paper in the receiver of SO2, letting no HNO3
drop from it. Continue this operation till a small quantity of
liquid is found in the bottle. The fumes show that HNO3 has lost
O. 2 HNO3 + SO2 = H2SO4 + 2 NO2.
91. Tests for H2SO4.
Experiment 56.--(1) Test the liquid with litmus. (2) Transfer it
to a t.t., and add an equal volume of BaCl2 solution. H2SO4 +
BaCl2 = ? Is BaSO4 soluble? (3) Put one drop H2SO4 from the
reagent bottle in 10 cc. H2O in a clean t.t., and add 1 cc. BaCl2
solution. Look for any cloudiness. This is the characteristic
test for H2SO4 and soluble sulphates, and so delicate that one
drop in a liter of H2O can be detected. (4) Instead of H2SO4, try
a little Na2SO4 solution. (5) Put two or three drops of strong
H2SO4 on writing-paper, and evaporate, high over a flame, so as
not to burn the paper. Examine it when dry. (6) Put a stick into
a t.t. containing 2 cc. H2SO4, and note the effect. (7) Review
Experiment 5. (8) Into an e.d. pour 5 cc. H2O, and then 15 cc.
H2SO4. Stir it meantime with a small t.t. containing 2 or 3 cc.
NH4OH, and notice what takes place in the latter; also note the
heat of the e.d.
The effects of (5), (6), (7), and (8) are due to the intense
affinity which H2SO4 has for H2O. So thirsty is it that it even
abstracts H and O from oxalic acid in the right proportion to
form H2O, combines them, and then absorbs the water.
92. Affinity for Water.--This acid is a desiccator or dryer, and
is used to take moisture from the air and prevent metallic
substances from rusting. In this way it dilutes itself, and may
increase its weight threefold. In diluting, the acid must always
be poured into the water slowly and with stirring, not water into
the acid, since, as H2O is lighter than H2SO4, heat enough may be
set free at the surface of contact to cause an explosion.
Contraction also takes place, as may be shown by accurately
measuring each liquid in a graduate, before mixing, and again
when cold. The mixture occupies less volume than the sum of the
two volumes. For the best results the volume of the acid should
be about three times that of the water.
93. Sulphuric Acid made on a Large Scale involves the same
principles as shown in Experiment 55, excepting that S02 is
obtained by burning S or roasting FeS2 (pyrite),
[Fig. 29.]
and HNO3 is made on the spot from NaNO3 and H2SO4. SO2 enters a
large leaden chamber, often 100 to 300 feet long, and jets of
steam and small portions of HNO3 are also forced in. The "chamber
acid" thus formed is very dilute, and must be evaporated first in
leaden pans, and finally in glass or platinum retorts, since
strong H2SO4, especially if hot, dissolves lead. See Experiment
124. Study Figure 29, and write the reactions. 2 HNO3 breaks up
into 2 NO2, H2O, and O. 94. Importance.--Sulphuric acid has been
called, next to human food, the most indispensable article known.
There is hardly a product of modern civilization in the
manufacture of which it is not directly or indirectly used.
Nearly a million tons are made yearly in Great Britain alone. It
is the basis of all acids, as Na2CO3 is of alkalies. It is the
life of chemical industry, and the quantity of it consumed is an
index of a people's civilization. Only a few of its uses can be
stated here. The two leading ones are the reduction of Ca3(PO4)2
for artificial manures and the sodium carbonate manufacture.
Foods depend on the productiveness of soils and on fertilizers,
and thus indirectly our daily bread is supplied by means of this
acid; and from sodium carbonate glass, soap, saleratus, baking-
powders, and most alkalies are made directly or indirectly. H2SO4
is employed in bleaching, dyeing, printing, telegraphy,
electroplating, galvanizing iron and wire, cleaning metals,
refining Au and Ag, making alum, blacking, vitriols, glucose,
mineral waters, ether, indigo, madder, nitroglycerine, gun-
cotton, parchment, celluloid, etc., etc.
FUMING SULPHURIC ACID.
95. Nordhausen or Fuming Sulphuric Acid, H2S207 used in
dissolving indigo and preparing coal-tar pigments, is made by
distilling FeSO4. 4FeSO4 + H2O = H2S207 + 2Fe203 + 2S02. This was
the original sulphuric acid. It is also formed when S03 is
dissolved in H2SO4. When exposed to the air, S03 escapes with
fuming.
CHAPTER XX.
AMMONIUM HYDRATE.
96. Preparation of Bases.--We have seen that many acids are made
by acting on a salt of the acid required, with a stronger acid.
This is the direct way. The following experiments will show that
bases may be prepared in a similar way by acting on salts of the
base required with other bases, which we may regard as stronger
than the ones to be obtained.
97. Preparation of NH4OH and NH3.
Experiment 57.--Powder 10 g. ammonium chloride, NH4Cl, in a mortar
and mix with 10 g. calcium hydrate, Ca(OH)2; recently slaked lime
is the best. Cover with water in a flask, and connect with Woulff
bottles, as for making HCl (Fig. 22); heat the flask for fifteen
minutes or more. The experiment may be tried on a smaller scale
with a t.t. if desired.
The reaction is: 2NH4Cl + Ca(OH)2 = CaCl2 + 2NH4OH. NH4OH is
broken up into NH3, ammonia gas, and water. NH4OH = NH3 + H2O.
These pass over into the first bottle, where the water takes up
the NH3, for which it has great affinity. One volume of water at
0° will absorb more than 1000 volumes of NH3. Thus NH4OH may be
called a solution of NH3, in H2O. Write the reaction.
Experiment 58.--Powder and mix 2 or 3 g. each of ammonium
nitrate, NH4NO3, and Ca(OH)2; put them into a t.t., and heat
slowly. Note the odor. 2NH4NO3 + Ca(OH)2 = ?
98. Tests.
Experiment 59.--(1) Generate a little of the gas in a t.t., and
note the odor. (2) Test the gas with wet red litmus paper. (3)
Put a little HCl into an e.d., and pass over it the fumes of NH3
from a d.t. Note the result, and write the equation. (4) Fill a
small t.t. with the gas by upward displacement; then, while still
inverted, put the mouth of the t.t. into water. Explain the rise
of the water. (5) How might NH4Cl be obtained from the NH4OH in
the Woulff bottles? (6) Test the liquid in each bottle with red
litmus paper. (7) Add some from the first bottle to 5 or 10 cc.
of a solution of FeSO4 or FeCl2, and look for a ppt. State the
reaction.
99. Formation.--Ammonia, hartshorn, exists in animal and
vegetable compounds, in salts, and, in small quantities, in the
atmosphere. Rain washes it from the atmosphere into the soil;
plants take it from the soil; animals extract it from plants.
Coal, bones, horns, etc., are the chief sources of it, and from
them it is obtained by distillation. It results also from
decomposing animal matter. NH3 can be produced by the direct
union of N and H, only by an electric discharge or by ozone. It
may be collected over Hg like other gases that are very soluble
in water.
100. Uses. --Ammonium hydrate, NH4OH, and ammonia, NH3, are used
in chemical operations, in making artificial ice, and to some
extent in medicine; from them also may be obtained ammonium
salts. State what you would put with NH4OH to obtain (NH4)2SO4.
To obtain NH4NO3. The use of NH4OH in the laboratory may be
illustrated by the following experiment:--
Experiment 60.--Into a t.t. put 10 cc. of a solution of ferrous
sulphate, FeSO4. Into another put 10 cc. of sodium sulphate
solution, Na2SO4. Add a little NH4OH to each. Notice a ppt. in
the one case but none in the other. If solutions of these two
compounds were mixed, the metals Fe and Na could be separated by
the addition of NH4OH, similar to the separation of Ag and Cu by
HCl. Try the experiment.
CHAPTER XXI.
SODIUM HYDRATE.
101. Preparation.
Experiment 61.--Dissolve 3 g. sodium carbonate, Na2CO3, in 10 or
15 cc. H2O in an e.d., and bring it to the boiling-point. Then
add to this a mixture of 1 or 2 g. calcium hydrate, Ca(OH)2, in 5
or 10cc. H2O. It will not dissolve. Boil the whole for five
minutes. Then pour off the liquid which holds NaOH in solution.
Evaporate if desired. This is the usual mode of preparing NaOH.
The reaction is Na2CO3 + Ca(OH)2 = 2NaOH + CaCO3. The residue is
Ca(OH)2 and CaCO3; the solution contains NaOH, which can be
solidified by evaporating the water. Sodium hydrate is an
ingredient in the manufacture of hard soap, and for this use
thousands of tons are made annually, mostly in Europe. It is an
important laboratory reagent, its use being similar to that of
ammonium hydrate. Exposed to the air, it takes up water and CO2,
forming a mixture of NaOH and Na2CO3. It is one of the strongest
alkalies, and corrodes the skin.
Experiment 62.--Put 20 cc. of H2O in a receiver. With the forceps
take a piece of Na, not larger than half a pea, from the naphtha
in which it is kept, drop it into the H2O, and at once cover the
receiver loosely with paper or cardboard. Watch the action, as
the Na decomposes H2O. HOH + Na = NaOH + H. If the water be hot
the action is so rapid that enough heat is produced to set the H
on fire. That the gas is H can be shown by putting the Na under
the mouth of a small inverted t.t., filled with cold water, in a
water-pan. Na rises to the top, and the t.t. fills with H, which
can be tested. NaOH dissolves in the water.102. Properties.
Experiment 63.--(1) Test with red litmus paper the solutions
obtained in the last two experiments. (2) To 5cc.of alum
solution, K2A12(SO4)4, add 2cc.of the liquid, and notice the
color and form of the ppt.
POTASSIUM HYDRATE.
103. KOH is made in the Same Way as NaOH.
Describe the process in full (Experiment 61), and give the
equation.
Experiment 64.--Drop a small piece of K into a receiver of H2O,
as in Experiment 62. The K must be very small, and the experiment
should not be watched at too close a range. The receiver should
not be covered with glass, but with paper. The H burns, uniting
with O of the air. The purple color is imparted by the burning,
or oxidation of small particles of K. Write the equation for the
combustion of each.
H2O might be considered the symbol of an acid, since it is the
union of H and a negative element; or, if written HOH, it might
be called a base, since it has a positive element and the (OH)
radical. It is neutral to litmus, and on this account might be
called a salt. It is better, however, to call it simply an oxide.
Potassium hydrate, caustic potash, is employed for the
manufacture of soft soap. As a chemical reagent its action is
almost precisely like that of caustic soda, though it is usually
considered a stronger base, as K is a more electro-positive
element than Na.
CALCIUM HYDRATE.
104. Calcium Hydrate, the Most Common of the Bases, is nearly as
important to them as H2SO4 is to acids. Since it is used to make
the other bases, it might be called the strongest base; as H2SO4
is often called the strongest acid. The strength of an acid or
base, however,depends on the substance to which it is applied, as
well as on itself, and for most purposes this one is classified
as a weaker base than the three previously described.
Sulphuric acid, the most useful of the acids, is not made
directly from its salts, but has to be synthesized. Calcium
hydrate is also made by an indirect process, as follows:
CaCO3, i.e. limestone, marble, etc., is burnt in kilns with C, a
process which separates the gas, CO2, according to the reaction:
CaCO3 = CaO + CO2. CaO is unslaked lime, or quick-lime. On
treating this with water, slaked lime, Ca(OH)2 is formed, with
generation of great heat. CaO + H2O = Ca(OH)2. Its affinity for
H2O is so great that it takes the latter from the air, if
exposed.
Experiment 65.--Saturate some unslaked lime with water, in an
e.d., and look for the results stated above, leaving it as long
as may be necessary.
105. Resume.--From the experiments in the last few chapters on
the three divisions of chemical compounds, acids, bases and
salts, we have seen (1) that acids and bases are the chemical
opposites of each other; (2) that salts are formed by the union
of acids and bases; (3) that some acids can be obtained from
their salts by the action of a stronger acid; (4) that some bases
can be got from salts by the similar action of other bases; (5)
that the strongest acids and bases, as well as others, may be
obtained in an indirect way by synthesis.
CHAPTER XXII.
OXIDES OF NITROGEN.
106. There are five oxides of N, only two of which are important.
NITROGEN MONOXIDE (N2O).
107. Preparation.
Experiment 66.--Put into a flask, holding 200cc, lOg of ammonium
nitrate, NH4NO3; heat it over wire gauze or asbestus in an iron
plate, having a d.t. connected with a large t.t., which is held
in a receiver of water, and from this t.t., another d.t. passing
into a pneumatic trough, so as to collect the gas over water
(Fig. 30). Have all the bearings tight. The reaction is NH4NO3 =
2H2O + N2O. The t.t. is for collecting the H2O.
[Fig. 30.]
Note the color of the liquid in the t.t.; taste a drop, and test
it with litmus. If the flask is heated too fast, some NO is
formed, and this taking O from the air makes NO2, which liquefies
and gives an acid reaction and a red color. Some NH4NO3 is also
liable to be carried over.
108. Properties.
Experiment 67.--Test the gas in the receiver with a burning stick
and a glowing one, and compare the combustion with that in O.
N20may also be tested with S and P, if desired. N is set free in
each case. Write the reactions.
Nitrogen monoxide or protoxide, the nitrous oxide of dentists,
when inhaled, produces insensibility to pain,-- anaesthesia,--
and, if continued, death from suffocation. Birds die in half a
minute from breathing it. Mixed with one-fourth O, and inhaled
for a minute or two, it produces intoxication and laughter, and
hence is called laughing gas. As made in Experiment 66, it
contains Cl and NO, as impurities, and should not be breathed.
NITROGEN DIOXIDE (NO, OR N2O2).
109. Preparation.
Experiment 68.--Into a t.t. or receiver put 5g Cu turnings, add 5
cc. H2O and 5 cc. HNO3. Collect the gas like H, over water. 3Cu +
8HNO3 = ? What two products will be left in the generator? Notice
the color of the liquid. This color is characteristic of Cu
salts. Notice also the red fumes of NO2.
110. Properties.
Experiment 69.--Test the gas with a burning stick, admitting as
little air as possible. Test it with burning S. NO is not a
supporter of C and S combustion. Put a small bit of P in a
deflagrating-spoon, and when it is vigorously burning, lower it
into the gas. It should continue to burn. State the reaction.
What combustion will NO support? Note that NO is half N, while
N2O is two-thirds N, and account for the difference in supporting
combustion.
NITROGEN TETROXIDE (NO2 or N2O4).
111. Preparation.
Experiment 70.--Lift from the water-pan a receiver of NO, and
note the colored fumes. They are NO2, or N2O4, nitrogen
tetroxide. NO + O = NO2. Is NO combustible? What is the source of
O in the experiment?OXIDES OF NITROGEN.
NITROGEN TRIOXIDE (N2O3).
112. Preparation.
Experiment 71.--Put into a t.t. 1 g. of starch and 1 cc. of HNO3.
Heat the mixture for a minute. The red fumes are N2O3 and NO2.
Nitrogen pentoxide, N2O5, is an unimportant solid. United with
water it forms HNO3. N2O5 + H2O = 2HNO3.
CHAPTER XXIII.
LAWS OF DEFINITE AND OF MULTIPLE PROPORTION.
113. Weight and Volume.--We have seen that water contains two
parts of H by volume to one part of O; or, by weight, two parts
of H to sixteen of O. These proportions are invariable, or no
symbol for water would be possible. Every compound in the same
way has an unvarying proportion of elements.
114. Law of Definite Proportion.--In a given compound the
proportion of any element by weight, or, if a gas, by volume is
always constant. Apply the law, by weight and by volume, to
these: HCl, NH3, H2S, N2O.
There is another law of equal importance in chemistry, which the
compounds of N and O well illustrate.
Weight. Volume.
N. O. N. O.
Nitrogen protoxide N2O 28 16 2 1
Nitrogen dioxide N2O2 28 32 2 2
Nitrogen trioxide. N2O3 28 48 2 3
Nitrogen tetroxide N2O4 28 64 2 4
Nitrogen pentoxide N2O5 28 80 2 5
Note that the proportion of O by weight is in each case a
multiple of the first, 16. Also that the proportion by volume of
O is a multiple of that in the first compound. In this example
the N remains the same. If that had varied in the different
compounds, it would also havevaried by a multiple of the smallest
proportion. This is true in all compounds.
115. Law of Multiple Proportion.--Whenever one element combines
with another in more than one proportion, it always combines in
some multiple, one or more, of its least combining weight, or, if
a gas, of its least combining volume.
The least combining weight of an element is its atomic weight;
and it is this fact of a least combining weight that leads us to
believe the atom to be indivisible.
Apply the law in the case of P2O, P2O3, P2O5; in HClO, HClO2,
HClO3, HClO4, arranging the symbols, weights, and volumes in a
table, as above.
The volumetric proportions of each element in the oxides of
nitrogen are exhibited below.
_ + _ + _ = __
N + N + O = N2O
_ + _ + _ + _ = __
N + N + O + O = N2O2
_ + _ + _ + _ + _ = __
N + N + O + O + O = N2O3
_ + _ + _ + _ + _ + _ = __
N + N + O + O + O + O = N2O4
_ + _ + _ + _ + _ + _ + _ = __
N + N + O + O + O + O + O = N2O5
CHAPTER XXIV.
CARBON PROTOXIDE.
116. Preparation.
Experiment 72.--Put into a flask, of 200 cc., 5 g. of oxalic acid
crystals, H2C2O4, and 25 cc. H2SO4. Have the d.t. pass into a
solution of NaOH in a Woulff bottle (Fig. 31), and collect
the gas over water. Heat the flask slowly, and avoid inhaling the
gas.
117. Tests.
Experiment 73.--Remove a receiver of the gas, and try to light
the latter with a splinter. Is it combustible, or a supporter of
(C) combustion? What is the color of the flame? When the
combustion ceases, shake up a little lime water with the gas left
in the receiver. What gas has been formed by the combustion, as
shown by the test? See page 80. Give the reaction for the
combustion.
We have seen that H2SO4 has great affinity for H2O. Oxalic acid
consists of H, C, O in the right proportion to form H2O, CO2, and
CO. H2SO4 withdraws H and O in the right proportion to form
water, unites them, and then absorbs the water, leaving the C and
O to combine and form CO2 and CO. NaOH solution removes CO2 from
the mixture, forming Na2CO3, and leaves CO. Write both reactions.
118. Carbon Protoxide, called also carbon monoxide, carbonic
oxide, etc., is a gas, having no color or taste, butpossessing a
faint odor. It is very poisonous. Being the lesser oxide of C, it
is formed when C is burned in a limited supply of O, whereas CO2
is always produced when O is abundant. The formation of each is
well shown by tracing the combustion in a coal fire. Air enters
at the bottom, and CO2 is first formed. C + 2O = CO2. As this gas
passes up, the white-hot coal removes one atom of O, leaving CO.
CO2 + C - 2CO. At the top, if the draft be open, a blue flame
shows the combustion of CO. CO + O = CO2. The same reduction of
CO2 takes place in the iron furnace, and whenever there is not
enough oxygen to form CO2, the product is CO.
Great care should be taken that this gas does not escape into the
room, as one per cent has proved fatal. Not all of it is burned
at the top of the coal; and when the stove door is open, the
upper drafts should be open also. It is the most poisonous of the
gases from coal; hence the danger from sleeping in a room having
a coal fire.
119. Water Gas.--CO is one of the constituents of "water gas,"
which, by reason of its cheapness, is supplanting gas made from
coal, as an illuminator, in some cities. It is made by passing
superheated steam over red-hot charcoal or coke. C unites with
the O of H2O, forming CO, and sets H free, thus producing two
inflammable gases. C + H2O --? As neither of these gives much
light, naphtha is distilled and mixed with them in small
quantities to furnish illuminating power See page 183.
CHAPTER XXV.
CARBON DIOXIDE.
120. Preparation.
Experiment 74.--Put into a t.t., or a bottle with a d.t. and a
thistle-tube, 10 or 20 g. CaCO3, marble in lumps; add as many
cubic centimeters of H2O, and half as much HCl, and collect the
gas by downward displacement (Fig. 39). Add more acid as needed.
CaCO3 + 2 HCl = CaCl2 + H2CO3. H2CO3 = H2O + CO2. H2CO3 is a very
weak compound, and at once breaks up. By some, its existence as a
compound is doubted.
121. Tests.
Experiment 75.--(1) Put a burning and a glowing stick into the
t.t. or bottle. (2) Hold the end of the d.t. directly against the
flame of a small burning stick. Does the gas support combustion?
(3) Pour a receiver of the gas over a candle flame. What does
this show of the weight of the gas? (4) Pass a little CO2 into
some H2O (Fig. 32), and test it with litmus. Give the reaction
for the solution of CO2 in H2O.
Experiment 76.--Put into a t.t. 51 cc. of clear Ca(OH)2 solution,
i.e. lime water; insert in this the end of a d.t. from a CO2
generator (Fig. 32). Notice any ppt. formed. It is CaCO3. Let the
action continue until the ppt. disappears and the liquid is
clear. Then remove the d.t., boil the clear liquid for a minute,
and notice whether the ppt. reappears.
122. Explanation.
Ca(OH)2 + CO2 = CaCO3 + H2O. The curious phenomena of this
experiment are explained by the solubility of CaCO3 in water
containing CO2, and its insolu-bility in water, having no CO2.
When all the Ca(OH)3 is combined, or changed to CaCO3, the excess
of CO2 unites with H2O, forming the weak acid H2CO3, which
dissolves the precipitate, CaCO3, and gives a clear liquid. On
heating this, H2CO3 gives up its CO2, and CaCO3 is
reprecipitated, not being soluble in pure water.
Lime water, Ca(OH)2 solution, is therefore a test for the
presence of CO2. To show that carbon dioxide is formed in
breathing, and in the combustion of C, and that it is present in
the air, perform the following experiment:
Experiment 77.--(1) Put a little lime water into a t.t., and blow
into it through a piece of glass tubing. Any turbidity shows
what? (2) Burn a candle for a few minutes in a receiver of air,
then take out the candle and shake up lime water with the gas.
(3) Expose some lime water in an e.d. to the air for some time.
133. Oxidation in the Human System.--Carbon dioxide, or carbonic
anhydride, carbonic acid, etc., CO2, is a heavy gas, without
color or odor. It has a sharp, prickly taste, and is commonly
reckoned as poisonous if inhaled in large quantities, though it
does not chemically combine with the blood as CO does. Ten per
cent in the air will sometimes produce death, and five per cent
produces drowsiness. It exists in minute portions in the
atmosphere, and often accumulates at the bottom of old wells and
caverns, owing to its slow diffusive power. Before going down
into one of these, the air should always be tested by lowering a
lighted candle. If this is extinguished, there is danger. CO2 is
the deadly "choke damp" after a mine explosion, CH4 being
converted into CO2 and H2O; a great deal is liberated during
volcanic eruptions, and it is formed in breathing by the union of
O in the air with C in the system. This union of C and O takes
place in the lungs and in all the tissues of the body, even on
the surface. Oxygen is taken into the lungs, passes through the
thin membrane into the blood, forms a weak chemical union with
the red corpuscles, and is conveyed by them to all parts of the
system. Throughout the body, wherever necessary, C and H are
supplied for the O, and unite with it to form CO2 and H2O. These
are taken up by the blood though they do not form a chemical
union with it, are carried to the lungs, and pass out, together
with the unused N and surplus O. The system is thus purified, and
the waste must be supplied by food. The process also keeps up the
heat of the body as really as the combustion of C or P in O
produces heat. The temperature of the body does not vary much
from 99 degrees F., any excess of heat passing off through
perspiration, and being changed into other forms of energy.
If, as in some fevers, the temperature rises above about 105
degrees F., the blood corpuscles are killed, and the person dies.
During violent exercise much material is consumed, circulation is
rapid, and quick breathing ensues. Oxygen is necessary for life.
A healthy person inhales plentifully; and this element is one of
nature's best remedies for disease. Deep and continued
inhalations in cold weather are better than furnace fires to heat
the system. All animals breathe O and exhale CO2. Fishes and
other aquatic animals obtain it, not by decomposing H2O, but from
air dissolved in water. Being cold-blooded, they need relatively
little; but if no fresh water is supplied to those in captivity,
they soon die of O starvation.
124. Oxidation in Water.--Swift-running streams are clear and
comparatively pure, because their organic impurities are
constantly brought to the surface and oxidized, whereas in
stagnant pools these impurities accumulate. Reservoirs of water
for city supply have sometimes been freed from impurities by
aeration, i.e. by forcing air into the water.
125. Deoxidation in Plants.--Since CO2 is so constantly poured
into the atmosphere, why does it not accumulate there in large
quantity? Why is there not less free O in the air to-day than
there was a thousand years ago? The answer to these questions is
found in the growth of vegetation. In the leaf of every plant are
thousands of little chemical laboratories; CO2 diffused in small
quantities in the air passes, together with a very little H2O,
into the leaf, usually from its under side, and is decomposed by
the radiant energy of the sun. The C is built into the woody
fiber of the tree, and the O is ready to be re-breathed or burned
again. CO2 contributes to the growth of plants, O to that of
animals; and the constituents of the atmosphere vary little from
one age to another. The compensation of nature is here well
shown. Plants feed upon what animals discard, transforming it
into material for the sustenance of the latter, while animals
prepare food for plants. All the C in plants is supposed to come
from the CO2 in the atmosphere. Animals obtain their supply from
plants. The utility of the small percentage of CO2 in the air is
thus seen.
126. Uses.--CO2 is used in making "soda-water," and in chemical
engines to put out fires in their early stages. In either case it
may be prepared by treating Na2CO3 or CaCO3 with H2SO4. Give the
reactions. On a small scale CO2 is made from HNaCO3. CO2 has a
very weak affinity for water, but probably forms with it H2CO3.
Much carbon dioxide can be forced into water under pressure. This
forms soda-water, which really contains no soda. The
justification for the name is the material from which it is
sometimes made. Salts from H2CO3, called carbonates, are
numerous, Na2CO3 and CaCO3 being the most important.
Chapter XXVI.
OZONE.
127. Preparation.
Experiment 78.--Scrape off the oxide from the surface of a piece
of phosphorus 2 cm long, put it into a wide-mouthed bottle, half
cover the P with water, cover the bottle with a glass, and leave
it for half an hour or more.
128. Tests.
Experiment 79.--Remove the glass cover, smell the gas, and hold
in it some wet iodo-starch paper. Look for any blue color. Iodine
has been set free, according to the reaction, 2 KI + 03= K20 + O2
+ I2, and has imparted a blue color to the starch, and ordinary
oxygen has been formed. Why will not oxygen set iodine free from
KI?. What besides ozone will liberate it?
129. Ozone, oxidized oxygen, active oxygen, etc., is an
allotropic form of O. Its molecule is 03, while that of ordinary
oxygen is 02.
Three atoms of oxygen are condensed into the space of two atoms
of ozone, or three molecules of O are condensed into two
molecules of ozone, or three liters of O are condensed into two
liters of ozone. Ozone is thus formed by oxidizing ordinary
oxygen. 02 + O = 03. This takes place during thunder storms and
in artificial electrical discharges. The quantity of ozone
produced is small, five per cent being the maximum, and the usual
quantity is far less than that.
Ozone is a powerful oxidizing agent, and will change S, P, and As
into their ic acids. Cotton cloth was formerly bleached, and
linen is now bleached, by spreading it on the grass and leaving
it for weeks to be acted on by ozone, which is usually present in
the air in small quantities, especially in the country. Ozone is
a disinfectant, like other bleaching agents, and serves to clear
the air of noxious gases and germs of infectious diseases. So
much ozone is reduced in this way that the air of cities contains
less of it than country air. A third is consumed in uniting with
the substance which it oxidizes, while two-thirds are changed
into oxygen, as in Experiment 79.
It is unhealthful to breathe much ozone, but a little in the air
is desirable for disinfection.
Ozone will cause the inert N of the air to unite with H, to form
ammonia. No other agent capable of doing this is known, so that
all the NH3 in the air, in fact all ammonium compounds taken up
by plants from soils and fertilizers, may have been made
originally through the agency of ozone. At a low temperature
ozone has been liquefied. It is then distinctly blue.
Electrolysis of water is the best mode of preparing this
substance in quantity. When prepared from P it is mixed with
P2O3.
Chapter XXVII.
CHEMISTRY OF THE ATMOSPHERE.
130. Constituents.--The four chief constituents of the atmosphere
are N, O, H2O, CO2, in the order of their abundance. What
experiments show the presence of N, O, and CO2 in the air? Set a
pitcher of ice water in a warm room, and the moisture that
collects on the outside is deposited from the air. This shows the
presence of H2O. Rain, clouds, fog, and dew prove the same. H2SO4
and CaCl2, on exposure to air, take up water. Experiment 18 shows
that there is not far from four times as much N as O by volume in
air. Hence if the atmosphere were a compound of N and O, and the
proportion of four to one were exact, its symbol would be N4O.
131. Air not a Compound.--The following facts show that air is
not a compound, but rather a mixture of these gases.
1. The proportion of N and O in the air, though it does not vary
much, is not always exactly the same. This could not be true if
it were a compound. Why?
2. If N4O were dissolved in water, the N would be four times the
O in volume; but when air is dissolved, less than twice as much N
as O is taken up.
3. No heat or condensation takes place when four measures of N
are brought in contact with one of O. It cannot then be N4O, for
the vapor density of N4O would be 36--i.e. (14 x 4 + 16) / 2; but
that of air is 14 1/2 nearly --i.e. (14 x 4 + 16) / 5. Analysis
shows about 79 parts of N to 21 parts of O by volume in air.
132. Water.--The volume of H2O, watery vapor, in the atmosphere
is very variable. Warm air will hold more than cold, and at any
temperature air may be near saturation, i.e. having all it will
hold at that temperature, or it may have little. But some is
always present; though the hot desert winds of North Africa are
not more than 1/15 saturated. A cubic meter of air at 25 degrees,
when saturated, contains more than 22 g. of water.
133. Carbon Dioxide.--Carbon dioxide does not make up more than
three or four parts in ten thousand of the air; but, in the whole
of the atmosphere, this gives a very large aggregate. Why does
not CO2 form a layer below the O and N?
134. Other Ingredients.--Other substances are found in the air in
minute portions, e.g. NH3 constitutes nearly one-millionth. Air
is also impregnated with living and dead germs, dust particles,
unburned carbon, etc., but these for the most part are confined
to the portion near the earth's surface. In pestilential regions
the germs of disease are said sometimes to contaminate the air
for miles around.
Chapter XXVIII.
THE CHEMISTRY OF WATER.
135. Pure Water.--Review the experiments for electrolysis, and
for burning H. Pure water is obtained by distillation.
Experiment 80.--Provide a glass tube 40 or 50 cm long and 3 or 4
cm in diameter. Fit to each end a cork with two perforations,
through one of which a long tube passes the entire length of the
larger tube (Fig. 32a). Connect one end of this with a flask of
water arranged for heating; pass the other end into an open
receptacle for collecting the distilled water. Into the other
perforations lead short tubes,-- the one for water to flow into
the large tube from a jet; the other, for the same to flow out.
This condenses the steam by circulating cold water around it. The
apparatus is called a Liebig's condenser. Put water into the
flask, boil it, and notice the condensed liquid. It is
comparatively pure water; for most of the substances in solution
have a higher boiling-point than water, and are left behind when
it is vaporized.
(Fig. 32a.)
136. Test.
Experiment 81.--Test the purity of distilled water by slowly
evaporating a few drops on Pt foil in a room free from dust.
There should be no spot or residue left on the foil. Test in the
same way undistilled water. 137. Water exists in Three States,--
solid, liquid, and vaporous. It freezes at 0 degrees, suddenly
expanding considerably as it passes into the solid state. It
boils, i.e. overcomes atmospheric pressure and is vaporized, at
100 degrees (760 mm pressure). If the pressure is greater, the
boiling-point is raised, i.e. it takes a higher temperature to
overcome a greater pressure. If there be less pressure, as on a
mountain, the boiling-point is lowered below 100 degrees. Salts
dissolved in water raise its boiling-point, and lower its
freezing-point to an extent depending on the kind and quantity of
the salt. Water, however, evaporates at all temperatures, even
from ice.
Pure water has no taste or smell, and, in small quantities, no
color. It is rarely if ever found on the earth. What is taken up
by the air in evaporation is nearly pure; but when it falls as
rain or snow, impurities are absorbed from the atmosphere. Water
falling after a long rain, especially in the country, is
tolerably free from impurities. Some springs have also nearly
pure water; but to separate all foreign matter from it, water
must be distilled. Even then it is liable to contain traces of
ammonia, or some other substance which vaporizes at a lower
temperature than water.
138. Sea-Water.--The ocean is the ultimate source of all water.
From it and from lakes, rivers, and soils, water is taken into
the atmosphere, falls as rain or snow, and sinks into the ground,
reappearing in springs, or flowing off in brooks and rivers to
the ocean or inland seas. Ocean water must naturally contain
soluble salts; and many salts which are not soluble in pure water
are dissolved in sea-water. In fact, there is a probability that
all elements exist to some extent in sea-water, but many of them
in extremely minute quantities. Sodium and magnesium salts are
the two most abundant, and the bitter taste is due to MgSO4 and
MgCl2. A liter of sea- water, nearly 1000 g., holds over 37 g. of
various salts, 29 of which are NaCl. See Hard Water.
139. River Water.--River water holds fewer salts, but has a great
deal of organic matter, living and dead, derived from the regions
through which it flows. To render this harmless for drinking,
such water should be boiled, or filtered through unglazed
porcelain. Carbon filters are now thought to possess but little
virtue for separating harmful germs.
140. Spring Water.--The water of springs varies as widely in
composition as do the rocks whence it bubbles forth. Sulphur
springs contain much H2S; many geysers hold SiO2 in solution;
chalybeate waters have compounds of Fe; others have Na2SO4, MgSO4
NaCl, etc.
CHAPTER XXIX.
THE CHEMISTRY OF FLAME.
141. Candle Flame.
Experiment 82.--Examine a candle flame, holding a dark object
behind it. Note three distinct portions: (1) a colorless interior
about the wick, (2) a yellow light-giving portion beyond that,
(3) a thin blue envelope outside of all, and scarcely
discernible. Hold a small stick across the flame so that it may
lie in all three parts, and observe that no combustion takes
place in the inner portion.
142. Explanation.--A candle of paraffine, or tallow, is chiefly
composed of compounds of C and H, in the solid state. The burning
wick melts the solid; the liquid is then drawn up by the wick
till the heat vaporizes and decomposes it, and O of the air comes
in contact with the outer heated portion of gas, and burns it
completely. Air tends to penetrate the whole body of the flame,
but only N can pass through uncombined, for the O that is left
after combustion in the outer portion seizes upon the compounds
of C and H in the next, or yellow, part. There is not enough O
here for complete combustion; at this temperature H burns before
C, and the latter is set free. In that state it is of course a
solid. Now an incandescent solid, or one glowing with heat, gives
light, while the combustion of a gas gives scarcely any light,
though it may produce great heat. While C in the middle flame is
glowing, during the moment of its dissociation from H, it gives
light. In the outer flame the temperature is high enough to burn
entirely the gaseous compounds of C and H together, so that no
solid C is set free, and hence no light is given except the faint
blue. No combustion takes place in the inner blue cone, because
no O reaches there.
By packing a wick into a cylindrical tin cup 5 or 10 cm high and
4 cm in diameter, containing alcohol, and lighting it, gunpowder
can be held in the middle of the flame in a def. spoon, without
burning. This shows the low temperature of that portion. Burning
P will also be extinguished, thus showing the exclusion of O.
143. Bunsen Flame.
Experiment 83.--Examine a Bunsen burner. Unscrew the top, and
note the orifices for the admission of gas and of air. Make a
drawing. Replace the parts; then light the gas at the top,
opening the air-holes at the base. Notice that the flame burns
with very little color. Try to distinguish the three parts, as in
the candle flame. These parts can best be seen by allowing direct
sunlight to fall on the flame and observing its shadow on a white
ground. Make a drawing of the flame. Hold across it a Pt wire and
note at what part the wire glows most. Also press down on the
flame for an instant with a cardboard or piece of paper; remove
before it takes fire, and notice the charred circle. Put the end
of a match into the blue cone, and note that it does not burn.
Put the end of a Pt wire into this blue cone, and observe that it
glows when near the top of the cone. What do these experiments
show? Ascertain whether this inner portion contains a combustible
material, by holding in it one end of a small d.t., and trying to
ignite any gas escaping at the other end. It should burn. This
shows that no combustion takes place in the interior of the
flame, because sufficient free O is not present.
Next, close the air-holes, and note that the flame is yellow and
gives much light. From this we infer the presence of solid
particles in an incandescent state. But these could not come from
the air. They must be C particles which have been set free from
the C and H compounds of the gas, just as in the candle flame.
The smoke that rises proves this. Hold an e.d. in the flame and
collect some C. Try the same with the air-holes open. 144. Light
and Heat of Flame.--Which of the two flames is hotter, the one
with the air-holes open, or that with them closed? Evidently the
former; for air is drawn in and mixes with the gas as it rises in
the tube, and, on reaching the flame at the top, the two are well
mingled, and the gaseous compounds of C and H burn at so high a
temperature that solid C is not freed; hence there is little
light. On closing the air-holes, no O can reach the flame except
from the outside, and the heat is much less intense.
(Fig 33.) (Fig 34.)
The H burns first, and sets the C free, which, while glowing,
gives the light. This again illustrates the facts (1) that flame
is caused by burning gas; (2) that light is produced by
incandescent solids. Charcoal, coke, and anthracite coal burn
without flame, or with very little, because of the absence of
gases.
145. Temperature of Combustion.
Experiment 84.--Light a Bunsen flame, with the basal orifices
open, and hold over it a fine wire gauze. Notice that the flame
does not rise above the gauze. Extinguish the light, and try to
ignite the gas above the gauze, holding the latter within 5 or 6
cm of the burner tube. Notice that it does not burn below the
gauze (Fig. 33).
Gas and O are both present. Evidently, then, the only condition
wanting for combustion is a sufficiently high temperature. The
gauze cools the gas below its kindling- point.
This principle is made use of in the miner's lamp of Davy (Fig.
34). In coal mines a very inflammable gas, CH4, called fire-damp,
issues from the coal. If this collects in large quantities and
mixes with O of the air, a kindling-point is all that is needed
to make a violent explosion. An ordinary lamp would produce this,
but the gauze lamp prevents it; for, though the inside may be
filled with burning gas, CH4, the flame cannot communicate with
the outside.
(Fig 35.) (Fig 36.)
a, reducing flame b, oxidizing flame
146. Oxidizing and Reducing Flames.--The hottest part of a Bunsen
flame is just above the inner blue cone (b, Fig. 36). Evidently
there is more O at that point. If a reducing agent, i.e. a
substance which takes up O, be put into this part of the flame,
the latter will remove the O and appropriate it, forming an
oxide. Cu heated there would become copper oxide. This part is
called the oxidizing flame. The inner blue part of the Bunsen
flame is devoid of O. It ought to remove O from an oxidizing
agent, i.e. a substance which supplies O. If copper oxide be
heated there (a, Fig. 36) by means of a mouth blow-pipe (Fig.
35), the flame will appropriate the O and leave the copper. This
is called the reducing flame. Only the upper part of this blue
central cone has heat enough to act in this way. By using a
prepared piece of metal, to make the flame thin and to shut off
the air, and then blowing the flame with a blow-pipe, greater
strength can be obtained in both oxidizing and reducing flames
(Fig. 36).
147. Combustible and Supporter Interchangeable.-- H was found to
burn in O. H was the combustible, O the supporter. Would O itself
burn in H?--i.e. would the combustible become the supporter, and
the supporter the combustible? As illuminating gas consists
largely of H, and as air is part O, we may try the experiment
with gas and air. Gas will burn in air. Will air burn in gas?
Experiment 85.--Fit a cork with two holes in it to the large end
of a lamp chimney. Through each hole pass a short piece of
tubing, and connect one of these with a rubber tube leading to a
gas-jet. Pass a metallic tube, long enough to reach the top of
the chimney, through the other, so that it will move easily up
and down. Turn on the gas, and light it at the top of the
chimney. Hold the end of the tube passing through the cork in the
flame for a minute, then draw it down to the middle of the
chimney (Fig. 37, a) and finally slowly remove it (b). Note that
O from the air is burning in the gas. Which is the supporter, and
which the combustible in this case? O will burn equally well in
an atmosphere of H, as can be shown by experiment.
148. Explosive Mixture of Gases.
Experiment 86.--Slowly turn down the burning gas of a Bunsen
lamp, having the orifices open, and notice that it suddenly
explodes and goes out at the top, but now burns at the base. As
the gas was gradually turned off, more air became mixed with it,
until there was the right proportion of each gas for an
explosion. Figure 38 shows the same thing. Light the gas at the
top a, when the tube c covers the jet b. Then gradually raise the
tube c. At a certain place there is the same explosion as with
the lamp.
149. Generalizations.--These experiments show (1) that three
conditions are necessary for combustion,--a combustible, a
supporter, and a burning temperature which varies for different
substances. Given these, "a fire" always results. The conditions
for "spontaneous combustion" do not differ from those of any
combustion. See Experiments 34, 112, 113, 114. (2) That
combustible and supporter are interchangeable. If H burns in O, O
will burn in H, the product, being the same in each case. (3) For
any combustion there must be a certain proportion of combustible
and of supporter. Twenty per cent of CO2 in the air dilutes the O
to such an extent that C will not burn. Hence the utility of the
chemical engine for putting out fires. (4) When two
gases, a combustible and a supporter, are mixed in the requisite
proportion, they form an explosive mixture, needing only the
kindling temperature to unite them.
Chemical combination is always accompanied by disengagement of
heat. Chemical dissociation is always accompanied by absorption
of heat. The disengagement, or the absorption, is not always
evident to the senses.
Combustion is the chemical combination of two or more substances
with the self-evident disengagement of great heat, and usually of
light.
The temperature of ignition varies greatly with different
substances. PH3 burns spontaneously at the usual temperatures of
the air. P takes fire at 60 degrees, but even at 10 degrees it
oxidizes with rapidity enough to produce phosphorescence. The
vapor of CS2 may be set on fire by a glass rod heated to 150
degrees, but a red-hot iron will not ignite illuminating gas.
Spontaneous combustion often takes place in woolen or cotton rags
which have been saturated with oil. The oil rapidly absorbs O,
and sets fire to the cloth. This is thought to be the origin of
some very destructive fires.
CHAPTER XXX.
CHLORINE.
150. Preparation.
Experiment 87.--Put into a t.t. 5 g. of fine granular MnO2 and 10
cc. HCl. Apply heat carefully, and collect the gas by downward
displacement in a receiver loosely covered with paper (Fig. 39).
Add more HCl if needed. Have a good draft of air, and do not
inhale the gas. If you have accidentally breathed it, inhale
alcohol vapor from a handkerchief; alcohol has great affinity for
Cl. Note the color of the gas, and compare its weight with that
of air.
MnO2 + 4 HCl = MnCl2 + 2 H2O + 2 Cl. How much Cl can be separated
with 5 g. MnO2?
If preferred, a flask may be used for a generator instead of a
t.t. Cl can be obtained directly from NaCl by adding H2SO4 (which
produces HCl) and MnO2. 2 NaCl + 2 H2SO4 + MnO2 = MnSO4 + Na2SO4
+ 2 H2O + 2 Cl. Try the experiment, using a t.t. and adding
water.
151. Cl from Bleaching-Powder.
Experiment 88.--Put a few grams of bleaching- powder into a small
beaker, and set this into a larger one. Cover the latter with
pasteboard or paper, through which passes a thistle-tube reaching
into the small beaker (Fig. 40). Pour through the tube a little
H2SO4 dilated with its volume of H2O.
152. Chlorine Water.--A solution of Cl in water is often useful,
and may be made as follows:-- Experiment 89.--To 3 or 4 crystals
of KClO3 add a few drops of HCl. Heat a minute, and when the gas
begins to disengage, pour in 10 cc. H2O, which dissolves the gas.
2 KClO3 + 4 HCl = 2 KCl + Cl2O4 + 2 H2O + 2 Cl.
153. Bleaching Properties.
Experiment 90.--Put into a receiver of Cl, preferably before
generating it, two pieces of Turkey red cloth, one wet, the other
dry; a small piece of printed paper and a written one; also a red
rose or a green leaf, each wet. Note from which the color is
discharged. If it is not discharged from all, put a little H2O
into the receiver, shake it well, and state what ones are
bleached.
Experiment 91.--(1) Add 5 cc. of Cl water to 5 cc. of indigo
solution. (2) Treat in the same way 5 cc. K2Cr2O7 (potassium
dichromate) solution, and record the results.
Indigo, writing-ink, and Turkey red or madder, are vegetable
pigments; printer's ink contains C, and K2Cr2O7 is a mineral
pigment. State what coloring matters Cl will bleach.
154. Disinfecting Power.
Experiment 92.--Pass a little H2S gas from a generator into a
t.t. containing Cl water. Look for a deposit of S. Notice that
the odor of H2S disappears. H2S + 2 Cl = 2 HCl + S.
155. A Supporter of Combustion.
Experiment 93.--Sprinkle into a receiver of Cl a very little fine
powder or filings of Cu, As, or Sb, and notice the combustion.
Observe that here is a case of combustion in which O does not
take part. Chlorides of the metals are of course formed. Write
the reactions. See whether Cl will support the combustion of
paper or of a stick of wood.
Experiment 94.--Warm 2 or 3 cc. of oil of turpentine (C1OH16) in
an evaporating-dish; dip a piece of tissue paper into it, and
very quickly thrust this into a receiver of Cl. It should take
fire and deposit carbon. C1OH16 + 16 Cl = ? Test the moisture on
the sides of the receiver with litmus. Clean the receiver with a
little petroleum.
Experiment 95.--Prepare a H generator with a lamp-tube bent as in
Figure 41. Light the H, observing the cautions in Experiment 23,
and when well burning, lower the flame into a receiver of Cl.
Observe the change of color which the flame undergoes as it comes
in contact with Cl. Give the reaction for the burning. Test with
litmus any moisture on the sides of the receiver. A mixture of Cl
and H, in direct sunlight combines with explosive violence;
whereas in diffused sunlight it combines slowly, and in darkness
it does not combine. From these experiments state the chief
properties of Cl, and what combustion it will support.
[Figure 41.]
156. Sources and Uses.--The great source of Cl is NaCl, though it
is often made from HCl. Its chief use is in making bleaching-
powder, one pound of which will bleach 300 to 500 pounds of
cloth. Cl is very easily liberated from this powder by a dilute
acid, or, slowly, by taking moisture from the air. Hence its use
as a disinfectant in destroying noxious gases and the germs of
infectious diseases. Cl attacks organic matter and germs as it
does the membrane of the throat or lungs, owing to its affinity
for H.
Cl is the best bleaching agent for cotton goods. It is not
suitable for animal materials, such as silk and wool, as it
attacks their fiber. It does not discharge either mineral or
carbon colors. The chemistry of bleaching is obscure.
As dry material will not bleach, Cl seems to unite with H in H2O
and to set O free. The O then unites with some portion of the
coloring matter, oxidizing it, and breaking up its molecule.
Colors bleached by Cl cannot be restored.
Chapter XXXI.
BROMINE.
Examine bromine, potassium bromide, sodium bromide, magnesium
bromide.
157. Preparation.
Experiment 96.--Pulverize 2 or 3 g. KBr, and mix it with about
the same bulk of MnO2. After putting this into a t.t, add as much
H2SO4, mix them together by shaking, attach a d.t., and conduct
the end of it into a t.t. that is immersed in a bottle of cold
water. Slowly heat the contents of the t.t., and notice the color
of the escaping vapor, and any liquid that condenses in the
receiver. Avoid inhaling the fumes, or getting them into the
eyes.
MnO2 + 2 KBr + 2 H2SO4 = ? Compare this with the equation for
making Cl from NaCl.
158. Tests.
Experiment 97.--Try the bleaching action of Br vapor as in the
case of Cl. Bleach a piece of litmus paper, and try to restore
the color with NH4OH. Explain its bleaching and disinfecting
action. Try the combustibility of As, Sb, and Cu.
159. Description.--Bromine at usual temperatures is a liquid
element; it is the only common one except Hg; it. quickly
evaporates on exposure to air. The chemistry of its manufacture
is like that of Cl; its bleaching and disinfecting powers are
similar to the latter, though they are not quite so strong as
those of Cl. Its affinity for H and for metals is also strongly
marked. A drop of Br on the skin produces a sore slow to heal.
Bromine salts are mainly KBr, NaBr, MgBr2. These in small
quantities accompany NaCl, and are most common in brine springs.
The world's supply of Br comes chiefly from West Virginia and
Ohio, over 300,000 pounds being produced from the salt (NaCl)
wells there in 1884. The water taken from these wells is nearly
evaporated, after which NaCl crystallizes out, leaving a thick
liquid--bittern, or mother liquor--which contains the salts of
Br. The bittern is treated with H2SO4 and Mn02, as above.
For transportation in large quantities, Br has to be made into
the salts NaBr and KBr, on account of the danger attending
leakage or breakage of the receptacles for Br.
160. Uses.--Its chief uses are in photography (page 167),
medicine, as KBr, and analytical chemistry.
Chapter XXXII.
IODINE.
Examine iodine, potassium iodide.
161. Preparation of I.
Experiment 98.--Put into a t.t. 2 or 3 g. of powdered KI mixed
with an equal bulk of MnO2, add H2SO4 enough to cover well, shake
together, complete the apparatus as for making Br, and heat.
Notice the color of the vapor, and any sublimate. The direct
product of the solidification of a vapor is called a sublimate.
The process is sublimation. Observe any crystals formed. Write
the reaction, and compare the process with that for making Br and
Cl. Compare the vapor density of I with that of Br and of Cl.
With that of air. What vapor is heavier than I? What acid and
what base are represented by KI?
162. Tests.
Experiment 99.--(1) Put a crystal of I in the palm of the hand
and watch it for a minute. (2) Put 2 or 3 crystals into a t.t.,
and warm it, meanwhile holding a stirring-rod half-way down the
tube. Notice the vapor, also a sublimate on the sides of the t.t.
and rod. (3) Add to 2 or 3 crystals in a t.t. 5 cc. of alcohol,
C2H5OH; warm it, and see whether a solution is formed. If so, add
5 cc. H2O and look for a ppt. of I. Does this show that I is not
at all soluble in H2O, or not so soluble as in alcohol?
163. Starch Solution and Iodine Test.
Experiment 100.--Pulverize a gram or two of starch, put it into
an evaporating-dish, add 4 or 5 drops of water, and mix; then
heat to the boiling-point 10 cc. H2O in a t.t., and pour it over
the starch, stirring it meanwhile.
(1) Dip into this starch paste a piece of paper, hold it in the
vapor of I, and look for a change of color. (2) Pour a drop of
the starch paste into a clean t.t., and add a drop or two of the
solution of I in alcohol. Add 5 cc. H2O, note the color, then
boil, and finally cool. (3) The presence of starch in a potato or
apple can be shown by putting a drop of I solution in alcohol on
a slice of either, and observing the color. (4) Try to dissolve a
few crystals of I in 5 cc. H2O by boiling. If it does not
disappear, see whether any has dissolved, by touching a drop of
the water to starch paste. This should show that I is slightly
soluble in water.
164. Iodo-Starch Paper.
Experiment 101.--Add to some starch paste that contains no I 5
cc. of a solution of KI, and stir the mixture. Why is it not
colored blue? Dip into this several strips of paper, dry them,
and save for use. This paper is called iodo-starch paper, and is
used as a test for ozone, chlorine, etc. Bring a piece of it in
contact with the vapor of chlorine, bromine, or ozone, and notice
the blue color.
Experiment 102.--Add a few drops of chlorine water to 2cc. of the
starch and KI solution in 10 cc. H2O. This should show the same
effect as the previous experiment.
165. Explanation.--Only free I, not compounds of it, will color
starch blue. It must first be set free from KI. Ozone, chlorine,
etc., have a strong affinity for K, and when brought in contact
with KI they unite with K and set free I, which then acts on the
starch present. Com- plete the equation: KI + Cl = ?
166. Occurrence.--The ultimate source of I is sea water, of which
it constitutes far too small a percentage to be separated
artificially. Sea-weeds, or algae, especially those growing in
the deep sea, absorb its salts--NaI, KI, etc.--from the water. It
thus forms a part of the plant, and from this much of the I of
commerce is obtained. Algae are collected in the spring, on the
coasts of Ireland, Scotland, and Normandy, where rough weather
throws them up. They are dried, and finally burned or distilled;
the ashes are leached to dissolve I salts; the water is nearly
evaporated, and the residue is treated with H2SO4, and MnO2, as
in the case of Br and Cl. I also occurs in Chili, as NaI and
NaIO3, mixed with NaNO3. This is an important source of the I
supply.
167. Uses.--I is much used in medicine, and was formerly employed
in taking daguerreotypes and photographs. Its solution in alcohol
or in ether is known as tincture of iodine.
168. Fluorine.--F, Cl, Br, I, are called halogens or haloids, and
exist in compounds--salts--in sea water. F is the most active of
all elements, combining with every element except O. Until
recently it has never been isolated, for as soon as set free from
one compound it attacks the nearest substance, and seems to be as
much averse to combining with itself, or to existing in the
elementary state, as to uniting with O. It is supposed to be a
gas, and, as is claimed, has lately been isolated by electrolysis
from HF in a Pt U-tube. Fluorite (CaF2) and cryolite (Al2F6 + 6
NaF) are its two principal mineral sources. The enamel of the
teeth contains F in composition.
CHAPTER XXXIII.
THE HALOGENS.
169. Halogens Compared.--The elements F, Cl, Br, I, form a
natural group. Their properties, as well as those of their
compounds, vary in a step-by-step way, as seen below. F is
sometimes an exception. They are best remembered by comparing
them with one another. Notice:
1. Similarity of name-ending. Each name ends in ine.
2. Similarity of origin. Salt water is the ultimate source of
all, except F.
3. Similarity of valence. Each is usually a monad.
4. Similarity of preparation. Cl, Br, I, are obtained from their
salts by means of MnO2 end H2SO4.
5. Variation in occurrence. Cl occurs in sea-salt, Br in sea-
water, I in sea-weed.
6. Variation in color; F being colorless, Cl green, Br red, I
violet.
7. Gradation in sp. gr.; F 19, Cl 35.5, Br 80, I 127.
8. Gradation in state, corresponding to sp. gr.; F being a light
gas, Cl a heavy gas, Br a liquid, I a solid.
9. Corresponding gradation in their usual chemical activity; F
being most active, then Cl, Br, and I.
10. Corresponding gradation in the strength of the H acids; the
strongest being HF, the next, HCl, etc.
11. Corresponding gradation in the explosibility of their N
compounds; the strongest NCl3, the next, NBr3, etc.
12. Corresponding gradation in the number of H and O acids; Cl 4,
Br 3, I 2.
170. Compounds.--The following are some of the oxides, acids, and
salts of the halogens. Name them.
CI2O (+H2O=) 2 HClO. The salts are hypochlorites, as Ca(ClO)2.
Cl2O3 (+H20=) 2 HClO2. The salts are chlorites, as KClO2.
Cl2O4
-- HClO3 The salts are chlorates, as KClO3.
-- HClO4 The salts are perchlorates, as KClO4,
-- HBrO The salts are ? KBrO,
-- -- The salts are wanting.
-- HBrO3. The salts are ? KBrO3,
-- HBrO4. The salts are ? KBrO4,
-- -- The salts are wanting.
-- -- The salts are wanting.
I2O5 (+H2O=) 2 HIO3. The salts are ? KIO3.
-- HIO4. The salts are ? KIO4.
F forms no oxides, and no acids except HF. HF, HCl, HBr, HI, are
striking illustrations of acids with no O. HClO4 is a very strong
oxidizing agent. A drop of it will set paper on fire, or with
powdered charcoal explode violently. This is owing to the ease
with which it gives up 0. Notice why its molecule is broken up
more readily than HC103. The higher the molecular tower, or the
more atoms it contains, the greater its liability to fall. Some
organic compounds contain hundreds of atoms, and hence are easily
broken down, or, as we say, are unstable. Inorganic compounds
are, as a rule, much more stable than organic ones. It is not
always true, however, that the compound with the least number of
atoms is the most stable. SO2 is more stable than SO3, but H2SO3
is less so than H2SO4.
Chapter XXXIV.
VAPOR DENSITY AND MOLECULAR WEIGHT.
Examine a liter measure, in the form of a cube,--cubic decimeter,
--and a cubic centimeter.
171. Gaseous Weights and Volumes.--A liter of H, at 0 degrees
and 760 mm., weighs nearly 0.09 g. This weight is called a crith.
Find the weight of H in the following, in criths and in grams: 15
1., 0.07 1., 50.3 1., 0.035 1., 0.6 1..
It has been estimated that there are (10) 24. molecules of H in a
liter. Does the number vary for different gases? The weight of a
molecule of H in parts of a crith is 1/(10) 24.; in parts of a
gram .09/(10) 24.. If the H molecule is composed of 2 atoms, what
is the weight of its atom in fractions of a crith? What in
fractions of a gram? The weight of the H atom is a microcrith.
What part of a crith is a microcrith?
172. Vapor Density.--Vapor density, or specific gravity referred
to H as the standard, (Physics) is the ratio of the weight of a
given volume of a gas or vapor to the weight of the same volume
of H. A liter of steam weighs nine times as much as a liter of H.
Its vapor density is therefore nine. For convenience, a definite
volume of H is usually taken as the standard, viz., the H atom.
The volume of the H atom and that of the half-molecule of H2O, or
of any gas are identical, each being represented by one square.
If, then, the standard of vapor density is the H atom, half the
molecular weight of a gas must be its vapor density, since it is
evident that we thus compare the weights of equal volumes. The
vapor density of H2O, steam, is found from the symbol as follows:
(2 + 16) / 2 = 9. To obtain the vapor density of any compound
from the formula, we have only to divide its molecular weight by
two. Find the vapor density of HCl, N2O, NO, C12H22O11, Cl, CO2,
HF, SO2. Explain each case.
The half-molecule, instead of the whole, is taken; because our
standard is the hydrogen atom, the smallest portion of matter, by
weight, known to science.
How many criths in a liter of HCl? How many grams? Compute the
number of criths and of grams in one liter of the compounds whose
symbols appear above.
PROBLEMS.
(1) A certain volume of H weighs 0.36 g. at standard temperature
and pressure. How many liters does it contain? If one liter
weighs 0.09 g., to weigh 0.36 g. it will take 0.36 / 0.09 = 4
liters.
(2) How many liters, or criths, of H in 63 g.? 2.7 g.? 1 g.? 5
g.? 250 g.? Explain each.
(3) Suppose the gas to be twice as heavy as H, how many liters in
0.36 g.? A liter of the gas will weigh 0.18 g. (0.09 X 2). In
0.36 g. there will be 0.36 / 0.18 = 2. Answer the question for 63
g., 2.7 g., etc.
(4) How many liters of Cl in each of the above numbers of grams?
(5) How many of HCl? H2O (steam)? CO2? Explain fully every case.
Vapor density is very easily determined from the formula by the
method given above. But in practice the formula is obtained from
the vapor density, and hence the method there given has to be
reversed.
173. Vapor Density of Oxygen.--Suppose we were to obtain the
vapor density of O. We should carefully seal and weigh a given
volume, say a liter, at a noted temperature and barometric
pressure, which are reducedto 0 degrees and 760 mm, and compare
it with the weight of the same volume of H. This has been done
repeatedly, and O has been found to weigh 16 times as much as H,
volume for volume, or, more exactly, 15.96+. Now a liter of each
gas has the same number of molecules, therefore the O molecule
weighs 16 times the H molecule. The half-molecule of each has the
same proportion, and the vapor density of O is 16. Atomic weight
is obtained in a very different way.
PROBLEMS.
(1) A liter of Cl is found to weigh 3.195 g. Compute its vapor
density, and explain fully.
(2) A liter of Hg vapor, under standard conditions, weighs 9 g.
Find its vapor density, and explain.
The vapor density of only a few elements has been satisfactorily
determined. See page 12. Some cannot be vaporized; others can be,
but only under conditions which prevent weighing them. The vapor
density of very many compounds also is unknown.
(3) A liter of CO2 weighs 1.98 g. Find the vapor density, and
from that the molecular weight, remembering that the latter is
twice the former. See whether it corresponds to that obtained
from the formula, CO2. This is,in fact, the way a formula is
ascertained, if the atomic weights of its elements are known.
(4) A liter of a compound gas weighs 2.88 g. Analysis shows that
its weight is half S and half O. As the atomic weight of S is 32,
and that of O is 16, what is the symbol for the gas?
Solution. Its molecular weight is 64, i.e. (2.88=0.09) X 2, of
which 32 is S and 32 O. The atomic weight of S is 32, hence there
is one atom of S, while of O there are two atoms. The formula is
SO2.
(5) A liter of a compound gas, which is found to contain 1 C and
3 O by weight, weighs 1.26 g. What is its formula? Atomic weights
are taken from page 12. Prove your answer.
(6) A liter of a compound of N and O weighs 1.98 g. The N is
7/11; and the O 4/11. What is the gas?
(7) A compound of N and H gas weighs 0.765 g. to the liter. The N
is 14/17 of the whole, the H 3/17. What gas is it? CHAPTER XXXV.
ATOMIC WEIGHT.
174. Definition.--We have seen that the molecular weight of a
compound, as well as of most elements, is obtained from the vapor
density by doubling the latter. It remains to explain how atomic
weights are obtained. The term is rather misleading. The atomic
weight of an element is its least combining weight, the smallest
portion that enters into chemical union, which is, of course, the
weight of an atom.
175. Atomic Weight of Oxygen.--Suppose we wish to find the atomic
weight of oxygen. We must find the smallest proportion by weight
in which it occurs in any compound. This can only be done by
analyzing all the compounds of O that can be vaporized. As
illustrative of these compounds take the six following:--
Wt. of other
Names. V. d. Mol. Wt. Wt. of O. Elem. Symbol.
Carbon monoxide... 14 28 16 12 ?
Carbon dioxide.... 22 44 32 12 ?
Hydrogen monoxide... 9 18 16 2 ?
Nitrogen monoxide... 22 44 16 28 ?
Nitrogen trioxide... 38 76 48 28 ?
Nitrogen pentoxide... 54 108 80 28 ?
176. Molecular Symbols.--From the vapor density of the gases--
column 2--we obtain their molecular weight-- column 3. To find
the proportion of O, it must be separated by chemical means from
its compounds and separately weighed. These relative weights are
given in column 4. Now the smallest weight of O which unites in
any case is its atomic weight. If any compound of O should in
future be found in which its combining weight is 8 or 4, that
would be called its atomic weight. By dividing the numbers in
column 4, wt. of O, by 16, the atomic weight of O, we obtain the
number of O atoms in the molecule. Subtracting the weights of O
from the molecular weights, we have the parts of the other
elements, column 5, and dividing these by the atomic weight of
the respective elements, we have the number of atoms of those
elements, these last, combined with the number of O atoms, give
the symbol. In this way complete the last column.
Show how to get the atomic weight of Cl from these compounds,
arranging them in tabular form, and completing as above: HCl,
KCl, NaCl, ZnCl2, MgCl2; the atomic weight of N in these: N2O,
NO, NH3.
177. Molecular and Atomic Volumes.--We thus see that vapor
density and atomic weight are obtained in two quite different
ways. In the case of elements the two are usually identical, i.e.
with the few whose vapor density is known; but this is not always
true, and it leads to interesting conclusions regarding atomic
volume. In O both vapor density and atomic weight are 16. This
gives 2 atoms of O to the molecule, i.e. the molecular weight /
the atomic weight. The size of an O atom is therefore half the
gaseous molecule, and is represented by one square. S has a vapor
density and an atomic weight of 32 each. Compute the number of
atoms in the molecule. Compute for I, in which the two are
identical, 127. P has an atomic weight of 31, while its vapor
density is 62. Its molecule must consist of 4 atoms, each half
the size of the H atom, The vapor density of As is 150, the
atomic weight 75. Compute the number of atoms in its molecule,
and represent their relative size. Hg has an atomic weight of
200, a vapor density of 100. Compute as before, and compare the
results with those on page 12. Ozone has an atomic weight of 16,
a vapor density 24. Compute.
Chapter XXXVI.
DIFFUSION AND CONDENSATION OF GASES.
178. Diffusion of Gases.--Oxygen is 16 times as heavy as H. If
the two gases were mixed, without combining, in a confined space,
it might be supposed that O would settle to the bottom and H rise
to the top. This would, in fact, take place at first, but only
for an instant, for all gases tend to diffuse or become
intimately mixed. The lighter the gas the more quickly it
diffuses.
179. Law of Diffusion of Gases.--The diffusibility of gases
varies inversely as the square roots of their vapor densities.
Compare the diffusibility of H with that of O. dif. H:dif. O::
sqrt(16): sqrt(1), or dif: H: dif. O:: 4: 1.
That is to say, if H and O be set free from separate receivers in
a room, the H will become intermingled with the atmosphere four
times as quickly as the O. Compare the diffusibility of O and N;
of Cl and H. Take the atomic weights of these, since they are the
same as the vapor densities. In case of a compound gas, half the
molecular weight must be taken for the vapor density; e.g. dif.
N20: dif. O.:: sqrt(16): sqrt(22).
180. Cause.--Diffusion is due to molecular motion; the lighter
the gas the more rapid the vibration of its molecules. Compare
the diffusibility of CO2 and that of Cl; of HCl and SO2; of HF
and I.
181. Liquefaction and Solidification of Gases.--Water boils at
100 degrees, under standard pressure, though evaporating at all
temperatures; it vaporizes at a lower point if the pressure be
less, as on a mountain, and at a higher temperature if the
pressure be greater, as at points below the sea level. Alcohol
boils at 78 degrees, standard pressure, and every liquid has a
point of temperature and pressure above which it must pass into
the gaseous state. Likewise every gas has a critical temperature
above which it cannot be liquefied at any pressure.
This condition was not recognized formerly, and before 1877, O,
H, N, C4, CO, NO, etc., had not been liquefied, though put under
a pressure of more than 2,000 atmospheres. They were called
permanent gases. In 1877 Cailletet and Pictet liquefied and
solidified these and others. The lowest temperature, about -225
degrees, was produced by suddenly releasing the pressure from
solid N to 4mm, which caused it rapidly to evaporate.
Evaporation, especially under diminished pressure, always lowers
the temperature by withdrawing heat.
These low degrees are indicated by a H thermometer, or if too low
for that, by a "thermo-electric couple" of copper and German
silver.
The pupil can easily liquefy SO, by passing it through a U-tube
which is surrounded by a mixture of ice and salt in a large
receiver. At the meeting of the American Association for the
Advancement of Science in 1887, a solid brick of CO2 was seen and
handled by the members, Liquid H is steel blue.
A few results obtained under a pressure of one atmosphere are:--
Boiling Points: C2H4--102 degrees; CH4--184 degrees; O--181
degrees; N --194 degrees; CO--190 degrees; NO--154 degrees; Air--
191 degrees.
Solidifying Points: Cl -102 degrees; HCl -115 degrees; Ether -129
degrees; Alcohol -130 degrees.
Chapter XXXVII.
SULPHUR.
Examine brimstone, flowers of sulphur, pyrite, chalcopyrite,
sphalerite, galenite, gypsum, barite.
182. Separation.
Experiment 103.--To a solution of 2 g. of sodium sulphide,, Na2S2
in 10 cc. H2O add 3 or 4cc. HCl, and look for a ppt. Filter, and
examine the residue. It is lac sulphur, or milk of sulphur.
183. Crystals from Fusion.
Experiment 104.--In a beaker of 25 or 50 cc. capacity put 20 g.
brimstone. Place this over a flame with asbestos paper
interposed, and melt it slowly. Note the color of the liquid,
then let it cool, watching for crystals. When partly solidified
pour the liquid portion into an evapo- rating-dish of water, and
observe the crystals of S forming in the beaker (Fig. 42). The
hard mass may be separated from the glass by a little HNO3 and a
thin knife-blade, or by CS2.
184. Allotropy.
Experiment 105.--Place in a t.t. 15g of brimstone, then heat
slowly till it melts. Notice the thin amber-colored liquid. The
temperature is now a little above 100 degrees. As the heat
increases, notice that it grows darker till it becomes black and
so viscid that it cannot be poured out. It is now above 200
degrees. Still heat, and observe that it changes to a slightly
lighter color, and is again a thin liquid. At this time it is
above 300 degrees. Now pour a little into an evaporating dish
containing water. Examine this, noticing that it can be stretched
like rubber. Leave it in the water till it becomes hard. Continue
heating thebrimstone in the t.t. till it boils at about 450
degrees, and note the color of the escaping vapor. Just above
this point it takes fire. Cool the t.t., holding it in the light
meantime, and look for a sublimate of S on the sides.
185. Solution.
Experiment 106.--Place in an evaporating-dish a gram of powdered
brimstone, and add 5cc, CS2, carbon disulphide. Stir, and see
whether S is dissolved. Put this in a draft of air, and note the
evaporation of the liquid CS2, and the deposit of S crystals.
These crystals are different in form from those resulting from
cooling from fusion.
186. Theory of Allotropy.--The last three experiments well
illustrate allotropy. We found S to crystallize in two different
ways. Substances can crystallize in seven different systems, and
usually a given substance is found in one of these systems only;
e.g. galena is invariably cubical. An element having two such
forms is said to be dimorphous. If it crystallizes in three
systems, it is trimorphous. A crystal has a definite arrangement
of its molecules. If without crystalline form, a substance is
called amorphous. An illustration of amorphism was S after it had
been poured into water. Thus S has at least three allotropic
forms, and the gradations between these probably represent
others. Allotropy seems to be due to varied molecular structure.
We know but little of the molecular condition of solids and
liquids, since we have no law to guide us like Avogadro's in
gases; but, from the density of S vapor at different
temperatures, we infer that liquids and solids have their
molecules very differently made up from those of gases. The least
combining weight of S is 32. Its vapor density at 1,000 degrees
is 32; hence its molecular weight is 64, i.e. vapor density x 2;
and there are 2 atoms in its molecule at that temperature,
molecular weight / atomic weight. At 500 degrees, however, the
vapor density is 96and the molecular weight 192. At this degree
the molecule must contain 6 atoms. How many it has in the
allotropic forms, as a solid, is beyond our knowledge; but it
seems quite likely that allotropy is due to some change of
molecular structure.
The above experiments show two modes of obtaining crystals, by
fusion and by solution.
187. Occurrence and Purification.--Sulphur occurs both free and
combined, and is a very common element. It is found free in all
volcanic regions, but Sicily furnishes most of it. Great
quantities are thrown up from the interior of the earth during an
eruption. The heat of volcanic action probably separates it from
its compound, which may be CaSO4. Vast quantities of the
poisonous SO2 gas are also liberated during an eruption, this
being, in volume of gases evolved, next to H2O. S is crudely
separated from its earthy impurities in Sicily by piling it into
heaps, covering to prevent access of air, and igniting, when some
of the S burns, and the rest melts and is collected. After
removal from the island it is further purified by distilling in
retorts connected with large chambers where it sublimes on the
sides as flowers of sulphur (Fig. 43). This is melted and run
into molds, forming roll brimstone. S also occurs as a
constituent of animal and vegetable compounds, as in mustard,
hair, eggs, etc. The tarnishing of silver spoons by eggs is due
to the formation of silver sulphide, Ag2S. The yellow color of
eggs, however, is due to oils, not to S.
The main compounds of S are sulphides and sulphates. What acids
do they respectively represent? Metallic sulphides are as common
as oxides; e.g. FeS2, or pyrite, PbS, or galenite, ZnS, or
sphalerite, CuFeS2, or chalcopyrite, etc. The most abundant
sulphate is CaSO4, or gypsum. BaSO4, or barite, and Na2SO4, or
Glauber's salt, are others.
The only one of these compounds that is utilized for its S is
FeS2. In Europe this furnishes a great deal of the S for H2SO4. S
is obtained by roasting FeS2. 3 FeS2 = Fe3S4 + 2 S.
188. Uses. -The greatest use of S is in the manufacture of H2SO4.
A great deal is used in making gunpowder, matches, vulcanized
rubber, and the artificial sulphides, like HgS, H2S, CS2, etc.
The last is a very volatile, ill- smelling liquid, made by the
combination of two solids, S being passed over red-hot charcoal.
It dissolves S, P, rubber, gums, and many other substances
insoluble in H2O.
189. Sulphur Dioxide, SO2, has been made in many experiments. It
is a bleaching agent, a disinfectant, and a very active compound,
having great affinity for water, but it will not support
combustion. Like most disinfectants, it is very injurious to the
system. It is used to bleach silk and wool--animal substances--
and straw goods, which Cl would injure; but the color can be
restored, as the coloring molecule seems not to be broken up, but
to combine with SO2, which is again separated by reagents. Goods
bleached with SO2 often turn yellow after a time.
190. SO2 a Bleacher.
Experiment 107.-Test its bleaching power by burning S under a
receiver under which a wet rose or a green leaf is also placed.
Chapter XXXVIII.
HYDROGEN SULPHIDE.
Examine ferrous sulphide, natural and artificial.
191. Preparation.
Experiment 108.--Put a gram of ferrous sulphide (FeS) into a t.t.
fitted with a d.t., as in Figure 32. Add 10cc. H2O and 5cc.
H2SO4. H2S is formed. Write the equation, omitting H2O. What is
left in solution?
192. Tests.
Experiment 109.-(1) Take the odor of the escaping gas. (2) Pour
into a t.t. 5cc.solution AgNO3, and place the end of the d.t.
from a H2S generator into the solution and note the color of the
ppt. What is the ppt.? Write the equation. (3) Experiment in the
same way with Pb(NO3)2 solution. Write the equation. (4) Let some
H2S bubble into a t.t. of clean water. To see whether H2S is
soluble in H2O, put a few drops of the water on a silver coin.
Ag2S is formed. Describe, and write the equation. Do the same
with a copper coin. (5) Put a drop of lead acetate solution,
Pb(C2H3O2)2, on a piece of unglazed paper, and hold this before
the d.t. from which H2S is escap- ing. PbS is formed. Write the
equation. This is the characteristic test of H2S.
193. Combustion of H2S
Experiment 110.--Attach a philosopher's lamp tube to the H2S
generator, and, observing the same precautions as with H, light
the gas. What two products must be formed? State the reaction.
The color of the flame. Compute the molecular weight and the
vapor density of H2S. 194. Uses. -Hydrogen sulphide or
sulphuretted hydrogen, H2S, is employed chiefly as a reagent in
the chemical laboratory. It forms sulphides with many of the
metals, as shown in the last experiment. These are precipitated
from solution, and may be separated from other metals which are
not so precipitated, as was found in the case of HCl and NH4OH.
The subjoined experiment will illustrate this. Suppose we wished
to separate Pb from Ba, having salts of the two mixed together,
as Pb(NO3)2 and Ba(NO3)2.
195. H2S an Analyzer of Metals.
Experiment 111.--Pass Some H2S gas in to 5cc.solution Ba(NO3)2.
No ppt. is formed. Do the same with Pb(NO3)2 solution. A ppt.
appears. Now mix 5cc.of each of these solutions in a t.t. and
pass the gas from a H2S generator into the liquid. What is
precipitated, and what is unchanged? When fully saturated with
the gas, as indicated by the smell, filter. Which metal is on the
filter and which is in the filtrate? Other reagents, as Na2CO3
solution, would precipitate the latter.
196. Occurrence and Properties. -- H2S is an ill-smell- ing,
poisonous gas, formed in sewers, rotten eggs, and other decaying
albuminous matter. It is formed in the earth, probably from the
action of water on sulphides, and issues with water from sulphur
springs.
A characteristic property is the formation of metallic sulphides,
as above. A skipper one night anchored his newly painted vessel
near the Boston gas-house, where the refuse was deposited, with
its escaping H2S. In the morning, to his consternation, the craft
was found to be black. H2S had come in contact with the lead in
the white paint, forming black PbS. This gradually oxidized after
reaching the open sea, and the white color reappeared.
Chapter XXXIX.
PHOSPHORUS.
NOTE.--Phosphorus should be kept in water, and handled with
forceps, never with the fingers, except under water, as it is
liable to burn the flesh and produce ulcerating sores. Pieces not
larger than half a pea should be used, and every bit should
finally be burned.
197. Solution and Combustion. Experiment 112. -Put 1 or 2 pieces
of P into an evaporating- dish, and pour over them 5 or 10cc.CS2
carbon disulphide. This will be enough for a class. When
dissolved, dip pieces of unglazed paper into it, and hold these
in the air, looking for any combustion as they dry. The P is
finely divided in solution, which accounts for its more ready
combustion then. Notice that the paper is not destroyed. This is
an example of so-called "spontaneous combustion." The burning-
point of P, the combustible, in air, the supporter, is about 60
degrees.
198. Combustion under Water.
Experiment 113. -Put a piece of P in a t.t. which rests in a
receiver, add a few crystals KClO3 and 5cc. H2O. Now pour in
through a thistle-tube 1cc.or more of H2SO4. Look for any flame.
H2SO4 acts very strongly on KClO3. What is set free? From this
fact explain the combustion in water.
199. Occurrence.--P is very widely disseminated, but not
abundant, and is found only in compounds, the chief of which is
calcium phosphate Ca3(PO4)2. It occurs in granite and other
rocks, as the mineral apatite, in soils, in plants, particularly
in seeds and grains, and in the bones, brains, etc., of
vertebrates. From the human system it is excreted by the kidneys
as microcosmic salt, HNaNH4PO4; and when the brain is hard-
worked, more than usual is excreted. Hence brain-workers have
been said to "burn phosphorus."
200. Sources.--Rocks are the ultimate source of this element.
These, by the action of heat, rain, and frost, are disintegrated
and go to make soils. The rootlets of plants are sent through the
soil, and, among other things, soluble phosphates in the earth
are absorbed, circulated by the sap, and selected by the various
tissues. Animals feed on plants, and the phosphates are
circulated through the blood, and deposited in the osseous
tissue, or wherever needed.
Human bones contain nearly 60 per cent of Ca3(PO4)2; those of
some birds over 80 per cent.
The main sources of phosphates and P are the phosphate beds of
South Carolina, the apatite beds of Canada, and the bones of
animals.
201. Preparation of Phosphates and Phosphorus.--Bone ash,
obtained by burning or distilling bones, and grinding the
residue, is treated with H1SO4, and forms soluble H4Ca(PO4)2,
superphosphate of lime, and insoluble CaSO4.
Ca3(PO4)2 + 2 H2SO4 = H4Ca(PO04)2 + 2 CaSO4. This completes the
process for fertilizers. If P is desired, the above is filtered;
charcoal, a reducing agent, is added to the filtrate; the
substance is evaporated, then very strongly heated and distilled
in retorts, the necks of which dip under water. It is then
purified from any uncombined C by melting in hot water and
passing into molds in cold water.
The work is very dangerous and injurious, on account of the low
burning-point of P, and its poisonous properties. While its
compounds are necessary to human life, P itself destroys the
bones, particularly the jaw bones, of the workers in it.
Between 1,000 and 2,000 tons are made yearly, mostly for matches,
but almost all at two factories, one in England, and one in
France. 202. Properties.--P is a colorless, transparent solid,
when pure; the impure article is yellowish, translucent, and
waxy. It is insoluble in water, slightly soluble in alcohol and
ether, and it readily dissolves in CS2, oil of turpentine, etc.
Fumes, having a garlic odor, rise when it is exposed to the air,
and in the dark it is phosphorescent, emitting a greenish light.
203. Uses. -The uses of this element and its compounds are for
fertilizers, matches, vermin poisons, and chemical operations.
204. Matches.-The use of P for matches depends on its low
burning-point. Prepared wood is dipped into melted S, and the end
is then pressed against a stone slab having on it a paste of P,
KClO3, and glue. KNO3 is often used instead of KClO3. In either
case the object is to furnish O to burn P. Matches containing
KClO3 snap on being scratched, while those having KNO3 burn
quietly. The friction from scratching a match generates heat
enough to ignite the P, that enough to set the S on fire, and the
S enough to burn the wood. Give the reaction for each. Paraffine
is much used instead of S. Safety matches have no P, and must be
scratched on a surface of red P and Sb2S3, or on glass.
205. Red Phosphorus.-Two or three allotropic forms of P are
known, the principal one being red. If heated between 230 degrees
and 260 degrees, away from air, the yellow variety changes to
red, which can be kept at all temperatures below 260 degrees.
Above that it changes back. Red P is not poisonous, ignites only
at a high temperature, and is not phosphorescent, like the
yellow. 206. Spontaneous Combustion of Phosphene, or Hydrogen
Phosphide, PH3.
Experiment 114.--Put into a 20cc.flask 1 g. P and 50cc.saturated
solution NaOH or KOH. Connect with the p.t. by a long d.t., as in
Figure 44, the end of which must be kept under water. Pour 3 or
4cc.of ether into the flask, to drive out the air. It is
necessary to exclude all air, as a dangerously explosive mixture
is formed with it. Heat the mixture, and as the gas passes over
and into the air, it takes fire spontaneously, and rings of smoke
successively rise. It will do no harm if, on taking away the
lamp, the water is drawn back into the flask; but in that case
the flask should be slightly lifted to prevent breakage by the
sudden rush of water. On no account let the air be drawn over.
The experiment has no practical value, but is an interesting
illustration of the spontaneous combustion of PH3 and of vortex
rings. What are the products of the combustion? An admixture of
another compound of P and H causes the combustion.
Chapter XL.
ARSENIC.
Examine metallic arsenic, realgar, orpiment, arsenopyrite,
arsenic trioxide, copper arsenite.
The compounds of arsenic are very poisonous if taken into the
system, and must be handled with care.
207. Separation. Experiment 115.--Draw out into two parts in the
Bunsen flame a piece of glass tubing 20cm long and 1 or 2cm in
diameter. Into the end of one of the ignition tubes thus formed,
when it is cool, put one-fourth of a gram of arsenic trioxide,
As2O3, using paper to transfer it. Now put into the tube a piece
of charcoal, and press it down to within 2 or 3cm of the AS2O3
(Fig. 45). Next heat the coal red-hot, and then at once heat the
As203. Continue this process till you see a metallic sublimate-
metallic mirror-on the tube above the coal. Break the tube and
examine the sublimate. It is As. Heat vaporizes the As2O;3.
Explain the chemical action. What is the agency of C in the
experiment? Of As2O3? 2 As2O3 + 3 C = ?
208. Tests.-Experiments 115 and 116 are used as tests for the
presence of arsenic.
Experiment 116.--Prepare a H generator, - a flask with a thistle-
tube and a philosopher's lamp tube (Fig. 46), put in some
granulated Zn, water, and HCl. Test the purity of the escaping
gas (Experiment 23), and when pure, light the jet of H. H is now
burning in air. To be sure that there is no As in the ingredients
used, hold the inside of a porcelain evaporating-dish directly
against the flame for a minute. If no silvery-white mirror is
found, the chemicals are free from As. Then pour through the
thistle-tube, while the lamp is still burning, 1cc.solution of
AS2O3 in HCl or H2O a bit of As2O3 not larger than a grain of
wheat in 10 cc. HCl.
See whether the color of the flame changes; then hold the
evaporating-dish once more in the flame, and notice a metallic
deposit of As. Set away the apparatus under the hood and leave
the light burning.
This experiment must not be performed unless all the cautions are
observed, since the gas in the flask (AsH3) is the most poisonous
known, and a single bubble of it inhaled is said to have killed
the discoverer. By confining the gas inside the flask there is no
danger.
Instead of using As2O3 solution, a little Paris green, wall paper
suspected of containing arsenic, green silk, or green paper
labels, etc., may be soaked in HCl, and tested.
209. Explanation.--The chemical changes are as follows: The
compounds of As, in this case As2O3, in presence of nascent H,
are immediately converted into the deadly hydrogen arsenide
(arsine, arseniuretted hydrogen), AsH3. As2O3 + 12 H = 2 AsH3 + 3
H2O. The AsH3 mixed with excess of H tends to escape and is
burned to As2O3 and H2O, and thus is rendered comparatively
harmless as it passes into the air. This is why the flame must be
burning when the arsenic compound is introduced. 2 AsH3 + 6 O =
As2O3 + 3 H2O.
In the combustion of AsH3, H burns at a lower point than As. The
introduction of a cold body like porcelain cools the flame below
the kindling-point of As, and this is deposited, while H burns,
in exactly the same way as lamp- black was collected in
Experiment 26.
210. Expert Analysis.--A modification of this experiment is
employed by experts to test for AS2O3 poisoning. The organs.--
stomach or liver--are cut into small pieces dissolved by nascent
Cl, or HClO, made from KC1O3 and HCl, and the solution is
introduced into a H generator, as above. AS2O3 preserves the
tissues it comes in contact with, for a long time, and the test
can be made years after death. All the chemicals must be pure,
since As is found in small quantities in most ores, and the Zn,
HCl, and H2SO4 of commerce are very likely to contain it. The
above is called Marsh's test, and is so delicate that a mere
trace of arsenic can be detected.
211. Properties and Occurrence.--As is a grayish white solid, of
metallic luster, while a few of its characters are non-metallic.
It is very widely distributed, being sometimes found native, and
sometimes combined, as AsS, realgar, As2S8, orpiment, and FeAsS,
arsenopyrite. Its chief source is the last, the fine powder of
which is strongly heated, when As separates and sublimes. It has
the odor of garlic, as may be observed by heating a little on
charcoal with the blow-pipe.
212. Atomic Volume.--As is peculiar in that its atomic volume, so
far as the volume can be determined, is only half that of the H
atom. Its vapor density is 150, which gives 300 for the molecular
weight, while its least combining or atomic weight is 75. 300,
the molecular weight = 75, the atomic weight =4, the number of
atoms in the molecule. All gaseous molecules being of the same
size, represented by two squares, the atomic volume of As must be
one-fourth of this size, represented by half of one square. Of
what other element is this true? 213. Uses of As2O3.-Arsenic is
used in shot-manufacture, for hardening the metal. Its most
important compound is As2O3, arsenic trioxide, called also
arsenious anhydride, arsenious acid, white arsenic, etc. So
poisonous is this that enough could be piled on a one-cent piece
to kill a dozen persons. Taken in too large quantities it acts as
an emetic. The antidote is ferric hydrate Fe2(OH)6 and a mustard
emetic, followed by oil or milk.
The vapor density of this compound shows that its symbol should
be As4O6, but the improper one, As2O3, is likely to remain in
use. Another oxide, As2O5, arsenic pentoxide, exists, but is less
important. Show how the respective acid formulae are obtained
from these anhydrides. See page 50.
AS2O3 is used in making Paris green; in many green coloring
materials, in which it exists as copper arsenite; in coloring
wall papers, and in fly and rat poisons. It is employed for
preserving skins, etc. Fashionable women sometimes eat it for the
purpose of beautifying the complexion, to which it imparts a
ghastly white, unhealthy hue. Mountaineers in some parts of
Europe eat it for the greater power of endurance which it is
supposed to give them. By beginning with small doses these
arsenic-eaters finally consume a considerable quantity of the
poison with apparent impunity; but as soon as the habit is
stopped, all the pangs of arsenic-poisoning set in. Wall paper
containing arsenic is said to be injurious to some people, while
apparently harmless to others.
Chapter XLI.
SILICON, SILICA, AND SILICATES.
214. Comparison of Si and C.--The element Si resembles carbon in
valence and in allotropic forms. It occurs in three forms like C,
a diamond form, a graphite, and an amorphous. C forms the basis
of the vegetable and animal world; Si, of the mineral. Most soils
and rocks, except limestone, are mainly compounds of O, Si, and
metals. While O is estimated to make up nearly one- half of the
known crust of the earth, Si constitutes fully a third. The two
are usually combined, as silica, SiO2, or silicates, SiO2
combined with metallic oxides. This affinity for O is so strong
that Si is not found uncombined, and is separated with great
difficulty and only at the highest tempera
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