An Elementary Study of Chemistry
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Title: An Elementary Study of Chemistry
Author: William McPherson
William Edwards Henderson
Release date: March 18, 2007 [eBook #20848]
Language: English
Other information and formats: www.gutenberg.org/ebooks/20848
Credits: Produced by Elaine Walker, Josephine Paolucci and the
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*** START OF THE PROJECT GUTENBERG EBOOK AN ELEMENTARY STUDY OF CHEMISTRY ***
Produced by Elaine Walker, Josephine Paolucci and the
Online Distributed Proofreading Team at http://www.pgdp.net
AN ELEMENTARY
STUDY OF CHEMISTRY
BY
WILLIAM McPHERSON, PH.D.
PROFESSOR OF CHEMISTRY, OHIO STATE UNIVERSITY
AND
WILLIAM EDWARDS HENDERSON, PH.D.
ASSOCIATE PROFESSOR OF CHEMISTRY, OHIO STATE UNIVERSITY
_REVISED EDITION_
GINN & COMPANY
BOSTON * NEW YORK * CHICAGO * LONDON
COPYRIGHT, 1905, 1906, BY
WILLIAM MCPHERSON AND WILLIAM E. HENDERSON
ALL RIGHTS RESERVED
The Athenæum Press
GINN & COMPANY * PROPRIETORS * BOSTON * U.S.A.
Transcriber's note:
For Text: A word surrounded by a cedilla such as ~this~ signifies that
the word is bolded in the text. A word surrounded by underscores like
_this_ signifies the word is italics in the text. The italic and bold
markup for single italized letters (such as variables in equations) and
"foreign" abbreviations are deleted for easier reading.
For numbers and equations: Parentheses have been added to clarify
fractions. Underscores before bracketed numbers in equations denote a
subscript. Superscripts are designated with a caret and brackets, e.g.
11.1^{3} is 11.1 to the third power.
Appendix A and B have been moved to the end of the book.
Minor typos have been corrected.
PREFACE
In offering this book to teachers of elementary chemistry the authors
lay no claim to any great originality. It has been their aim to prepare
a text-book constructed along lines which have become recognized as best
suited to an elementary treatment of the subject. At the same time they
have made a consistent effort to make the text clear in outline, simple
in style and language, conservatively modern in point of view, and
thoroughly teachable.
The question as to what shall be included in an elementary text on
chemistry is perhaps the most perplexing one which an author must
answer. While an enthusiastic chemist with a broad understanding of the
science is very apt to go beyond the capacity of the elementary student,
the authors of this text, after an experience of many years, cannot help
believing that the tendency has been rather in the other direction. In
many texts no mention at all is made of fundamental laws of chemical
action because their complete presentation is quite beyond the
comprehension of the student, whereas in many cases it is possible to
present the essential features of these laws in a way that will be of
real assistance in the understanding of the science. For example, it is
a difficult matter to deduce the law of mass action in any very simple
way; yet the elementary student can readily comprehend that reactions
are reversible, and that the point of equilibrium depends upon, rather
simple conditions. The authors believe that it is worth while to
present such principles in even an elementary and partial manner because
they are of great assistance to the general student, and because they
make a foundation upon which the student who continues his studies to
more advanced courses can securely build.
The authors have no apologies to make for the extent to which they have
made use of the theory of electrolytic dissociation. It is inevitable
that in any rapidly developing science there will be differences of
opinion in regard to the value of certain theories. There can be no
question, however, that the outline of the theory of dissociation here
presented is in accord with the views of the very great majority of the
chemists of the present time. Moreover, its introduction to the extent
to which the authors have presented it simplifies rather than increases
the difficulties with which the development of the principles of the
science is attended.
The oxygen standard for atomic weights has been adopted throughout the
text. The International Committee, to which is assigned the duty of
yearly reporting a revised list of the atomic weights of the elements,
has adopted this standard for their report, and there is no longer any
authority for the older hydrogen standard. The authors do not believe
that the adoption of the oxygen standard introduces any real
difficulties in making perfectly clear the methods by which atomic
weights are calculated.
The problems appended to the various chapters have been chosen with a
view not only of fixing the principles developed in the text in the mind
of the student, but also of enabling him to answer such questions as
arise in his laboratory work. They are, therefore, more or less
practical in character. It is not necessary that all of them should be
solved, though with few exceptions the lists are not long. The answers
to the questions are not directly given in the text as a rule, but can
be inferred from the statements made. They therefore require independent
thought on the part of the student.
With very few exceptions only such experiments are included in the text
as cannot be easily carried out by the student. It is expected that
these will be performed by the teacher at the lecture table. Directions
for laboratory work by the student are published in a separate volume.
While the authors believe that the most important function of the
elementary text is to develop the principles of the science, they
recognize the importance of some discussion of the practical application
of these principles to our everyday life. Considerable space is
therefore devoted to this phase of chemistry. The teacher should
supplement this discussion whenever possible by having the class visit
different factories where chemical processes are employed.
Although this text is now for the first time offered to teachers of
elementary chemistry, it has nevertheless been used by a number of
teachers during the past three years. The present edition has been
largely rewritten in the light of the criticisms offered, and we desire
to express our thanks to the many teachers who have helped us in this
respect, especially to Dr. William Lloyd Evans of this laboratory, a
teacher of wide experience, for his continued interest and helpfulness.
We also very cordially solicit correspondence with teachers who may find
difficulties or inaccuracies in the text.
The authors wish to make acknowledgments for the photographs and
engravings of eminent chemists from which the cuts included in the text
were taken; to Messrs. Elliott and Fry, London, England, for that of
Ramsay; to The Macmillan Company for those of Davy and Dalton, taken
from the Century Science Series; to the L. E. Knott Apparatus Company,
Boston, for that of Bunsen.
THE AUTHORS
OHIO STATE UNIVERSITY
COLUMBUS, OHIO
CONTENTS
CHAPTER PAGE
I. INTRODUCTION 1
II. OXYGEN 13
III. HYDROGEN 28
IV. WATER AND HYDROGEN DIOXIDE 40
V. THE ATOMIC THEORY 59
VI. CHEMICAL EQUATIONS AND CALCULATIONS 68
VII. NITROGEN AND THE RARE ELEMENTS IN THE ATMOSPHERE 78
VIII. THE ATMOSPHERE 83
IX. SOLUTIONS 94
X. ACIDS, BASES, AND SALTS; NEUTRALIZATION 106
XI. VALENCE 116
XII. COMPOUNDS OF NITROGEN 122
XIII. REVERSIBLE REACTIONS AND CHEMICAL EQUILIBRIUM 137
XIV. SULPHUR AND ITS COMPOUNDS 143
XV. PERIODIC LAW 165
XVI. THE CHLORINE FAMILY 174
XVII. CARBON AND SOME OF ITS SIMPLER COMPOUNDS 196
XVIII. FLAMES,--ILLUMINANTS 213
XIX. MOLECULAR WEIGHTS, ATOMIC WEIGHTS, FORMULAS 223
XX. THE PHOSPHORUS FAMILY 238
XXI. SILICON, TITANIUM, BORON 257
XXII. THE METALS 267
XXIII. THE ALKALI METALS 274
XXIV. THE ALKALINE-EARTH FAMILY 300
XXV. THE MAGNESIUM FAMILY 316
XXVI. THE ALUMINIUM FAMILY 327
XXVII. THE IRON FAMILY 338
XXVIII. COPPER, MERCURY, AND SILVER 356
XXIX. TIN AND LEAD 370
XXX. MANGANESE AND CHROMIUM 379
XXXI. GOLD AND THE PLATINUM FAMILY 390
XXXII. SOME SIMPLE ORGANIC COMPOUNDS 397
INDEX 421
APPENDIX A Facing back cover
APPENDIX B Inside back cover
LIST OF FULL-PAGE ILLUSTRATIONS
PAGE
ANTOINE LAURENT LAVOISIER _Frontispiece_
JOSEPH PRIESTLEY 14
JOHN DALTON 60
WILLIAM RAMSAY 82
DMITRI IVANOVITCH MENDELÉEFF 166
HENRI MOISSAN 176
SIR HUMPHRY DAVY 276
ROBERT WILHELM BUNSEN 298
AN ELEMENTARY STUDY OF CHEMISTRY
CHAPTER I
INTRODUCTION
~The natural sciences.~ Before we advance very far in the study of nature,
it becomes evident that the one large study must be divided into a
number of more limited ones for the convenience of the investigator as
well as of the student. These more limited studies are called the
_natural sciences_.
Since the study of nature is divided in this way for mere convenience,
and not because there is any division in nature itself, it often happens
that the different sciences are very intimately related, and a thorough
knowledge of any one of them involves a considerable acquaintance with
several others. Thus the botanist must know something about animals as
well as about plants; the student of human physiology must know
something about physics as well as about the parts of the body.
~Intimate relation of chemistry and physics.~ Physics and chemistry are
two sciences related in this close way, and it is not easy to make a
precise distinction between them. In a general way it may be said that
they are both concerned with inanimate matter rather than with living,
and more particularly with the changes which such matter may be made to
undergo. These changes must be considered more closely before a
definition of the two sciences can be given.
~Physical changes.~ One class of changes is not accompanied by an
alteration in the composition of matter. When a lump of coal is broken
the pieces do not differ from the original lump save in size. A rod of
iron may be broken into pieces; it may be magnetized; it may be heated
until it glows; it may be melted. In none of these changes has the
composition of the iron been affected. The pieces of iron, the
magnetized iron, the glowing iron, the melted iron, are just as truly
iron as was the original rod. Sugar may be dissolved in water, but
neither the sugar nor the water is changed in composition. The resulting
liquid has the sweet taste of sugar; moreover the water may be
evaporated by heating and the sugar recovered unchanged. Such changes
are called _physical changes_.
DEFINITION: _Physical changes are those which do not involve a change in
the composition of the matter._
~Chemical changes.~ Matter may undergo other changes in which its
composition is altered. When a lump of coal is burned ashes and
invisible gases are formed which are entirely different in composition
and properties from the original coal. A rod of iron when exposed to
moist air is gradually changed into rust, which is entirely different
from the original iron. When sugar is heated a black substance is formed
which is neither sweet nor soluble in water. Such changes are evidently
quite different from the physical changes just described, for in them
new substances are formed in place of the ones undergoing change.
Changes of this kind are called _chemical changes_.
DEFINITION: _Chemical changes are those which involve a change in the
composition of the matter._
~How to distinguish between physical and chemical changes.~ It is not
always easy to tell to which class a given change belongs, and many
cases will require careful thought on the part of the student. The test
question in all cases is, Has the composition of the substance been
changed? Usually this can be answered by a study of the properties of
the substance before and after the change, since a change in composition
is attended by a change in properties. In some cases, however, only a
trained observer can decide the question.
~Changes in physical state.~ One class of physical changes should be noted
with especial care, since it is likely to prove misleading. It is a
familiar fact that ice is changed into water, and water into steam, by
heating. Here we have three different substances,--the solid ice, the
liquid water, and the gaseous steam,--the properties of which differ
widely. The chemist can readily show, however, that these three bodies
have exactly the same composition, being composed of the same substances
in the same proportion. Hence the change from one of these substances
into another is a physical change. Many other substances may, under
suitable conditions, be changed from solids into liquids, or from
liquids into gases, without change in composition. Thus butter and wax
will melt when heated; alcohol and gasoline will evaporate when exposed
to the air. _The three states--solid, liquid, and gas--are called the
three physical states of matter._
~Physical and chemical properties.~ Many properties of a substance can be
noted without causing the substance to undergo chemical change, and are
therefore called its _physical properties_. Among these are its physical
state, color, odor, taste, size, shape, weight. Other properties are
only discovered when the substance undergoes chemical change. These are
called its _chemical properties_. Thus we find that coal burns in air,
gunpowder explodes when ignited, milk sours when exposed to air.
~Definition of physics and chemistry.~ It is now possible to make a
general distinction between physics and chemistry.
DEFINITION: _Physics is the science which deals with those changes in
matter which do not involve a change in composition._
DEFINITION: _Chemistry is the science which deals with those changes in
matter which do involve a change in composition._
~Two factors in all changes.~ In all the changes which matter can undergo,
whether physical or chemical, two factors must be taken into account,
namely, _energy_ and _matter_.
~Energy.~ It is a familiar fact that certain bodies have the power to do
work. Thus water falling from a height upon a water wheel turns the
wheel and in this way does the work of the mills. Magnetized iron
attracts iron to itself and the motion of the iron as it moves towards
the magnet can be made to do work. When coal is burned it causes the
engine to move and transports the loaded cars from place to place. When
a body has this power to do work it is said to possess energy.
~Law of conservation of energy.~ Careful experiments have shown that when
one body parts with its energy the energy is not destroyed but is
transferred to another body or system of bodies. Just as energy cannot
be destroyed, neither can it be created. If one body gains a certain
amount of energy, some other body has lost an equivalent amount. These
facts are summed up in the law of conservation of energy which may be
stated thus: _While energy can be changed from one form into another, it
cannot be created or destroyed._
~Transformations of energy.~ Although energy can neither be created nor
destroyed, it is evident that it may assume many different forms. Thus
the falling water may turn the electric generator and produce a current
of electricity. The energy lost by the falling water is thus transformed
into the energy of the electric current. This in turn may be changed
into the energy of motion, as when the current is used for propelling
the cars, or into the energy of heat and light, as when it is used for
heating and lighting the cars. Again, the energy of coal may be
converted into energy of heat and subsequently of motion, as when it is
used as a fuel in steam engines.
Since the energy possessed by coal only becomes available when the coal
is made to undergo a chemical change, it is sometimes called _chemical
energy_. It is this form of energy in which we are especially interested
in the study of chemistry.
~Matter.~ Matter may be defined as that which occupies space and possesses
weight. Like energy, matter may be changed oftentimes from one form into
another; and since in these transformations all the other physical
properties of a substance save weight are likely to change, the inquiry
arises, Does the weight also change? Much careful experimenting has
shown that it does not. The weight of the products formed in any change
in matter always equals the weight of the substances undergoing change.
~Law of conservation of matter.~ The important truth just stated is
frequently referred to as the law of conservation of matter, and this
law may be briefly stated thus: _Matter can neither be created nor
destroyed, though it can be changed from one form into another._
~Classification of matter.~ At first sight there appears to be no limit to
the varieties of matter of which the world is made. For convenience in
study we may classify all these varieties under three heads, namely,
_mechanical mixtures_, _chemical compounds_, and _elements_.
[Illustration: Fig. 1]
~Mechanical mixtures.~ If equal bulks of common salt and iron filings are
thoroughly mixed together, a product is obtained which, judging by its
appearance, is a new substance. If it is examined more closely, however,
it will be seen to be merely a mixture of the salt and iron, each of
which substances retains its own peculiar properties. The mixture tastes
just like salt; the iron particles can be seen and their gritty
character detected. A magnet rubbed in the mixture draws out the iron
just as if the salt were not there. On the other hand, the salt can be
separated from the iron quite easily. Thus, if several grams of the
mixture are placed in a test tube, and the tube half filled with water
and thoroughly shaken, the salt dissolves in the water. The iron
particles can then be filtered from the liquid by pouring the entire
mixture upon a piece of filter paper folded so as to fit into the
interior of a funnel (Fig. 1). The paper retains the solid but allows
the clear liquid, known as the _filtrate_, to drain through. The iron
particles left upon the filter paper will be found to be identical with
the original iron. The salt can be recovered from the filtrate by
evaporation of the water. To accomplish this the filtrate is poured into
a small evaporating dish and gently heated (Fig. 2) until the water has
disappeared, or _evaporated_. The solid left in the dish is identical in
every way with the original salt. Both the iron and the salt have thus
been recovered in their original condition. It is evident that no new
substance has been formed by rubbing the salt and iron together. The
product is called a _mechanical mixture_. Such mixtures are very common
in nature, almost all minerals, sands, and soils being examples of this
class of substances. It is at once apparent that there is no law
regulating the composition of a mechanical mixture, and no two mixtures
are likely to have exactly the same composition. The ingredients of a
mechanical mixture can usually be separated by mechanical means, such as
sifting, sorting, magnetic attraction, or by dissolving one constituent
and leaving the other unchanged.
[Illustration: Fig. 2]
DEFINITION: _A mechanical mixture is one in which the constituents
retain their original properties, no chemical action having taken place
when they were brought together._
~Chemical compounds.~ If iron filings and powdered sulphur are thoroughly
ground together in a mortar, a yellowish-green substance results. It
might easily be taken to be a new body; but as in the case of the iron
and salt, the ingredients can readily be separated. A magnet draws out
the iron. Water does not dissolve the sulphur, but other liquids do, as,
for example, the liquid called carbon disulphide. When the mixture is
treated with carbon disulphide the iron is left unchanged, and the
sulphur can be obtained again, after filtering off the iron, by
evaporating the liquid. The substance is, therefore, a mechanical
mixture.
If now a new portion of the mixture is placed in a dry test tube and
carefully heated in the flame of a Bunsen burner, as shown in Fig. 3, a
striking change takes place. The mixture begins to glow at some point,
the glow rapidly extending throughout the whole mass. If the test tube
is now broken and the product examined, it will be found to be a hard,
black, brittle substance, in no way recalling the iron or the sulphur.
The magnet no longer attracts it; carbon disulphide will not dissolve
sulphur from it. It is a new substance with new properties, resulting
from the chemical union of iron and sulphur, and is called iron
sulphide. Such substances are called _chemical compounds_, and differ
from mechanical mixtures in that the substances producing them lose
their own characteristic properties. We shall see later that the two
also differ in that the composition of a chemical compound never varies.
[Illustration: Fig. 3]
DEFINITION: _A chemical compound is a substance the constituents of
which have lost their own characteristic properties, and which cannot be
separated save by a chemical change._
~Elements.~ It has been seen that iron sulphide is composed of two
entirely different substances,--iron and sulphur. The question arises,
Do these substances in turn contain other substances, that is, are they
also chemical compounds? Chemists have tried in a great many ways to
decompose them, but all their efforts have failed. Substances which have
resisted all efforts to decompose them into other substances are called
_elements_. It is not always easy to prove that a given substance is
really an element. Some way as yet untried may be successful in
decomposing it into other simpler forms of matter, and the supposed
element will then prove to be a compound. Water, lime, and many other
familiar compounds were at one time thought to be elements.
DEFINITION: _An element is a substance which cannot be separated into
simpler substances by any known means._
~Kinds of matter.~ While matter has been grouped in three classes for the
purpose of study, it will be apparent that there are really but two
distinct kinds of matter, namely, compounds and elements. A mechanical
mixture is not a third distinct kind of matter, but is made up of
varying quantities of either compounds or elements or both.
~Alchemy.~ In olden times it was thought that some way could be found to
change one element into another, and a great many efforts were made to
accomplish this transformation. Most of these efforts were directed
toward changing the commoner metals into gold, and many fanciful ways
for doing this were described. The chemists of that time were called
_alchemists_, and the art which they practiced was called _alchemy_. The
alchemists gradually became convinced that the only way common metals
could be changed into gold was by the wonderful power of a magic
substance which they called the _philosopher's stone_, which would
accomplish this transformation by its mere touch and would in addition
give perpetual youth to its fortunate possessor. No one has ever found
such a stone, and no one has succeeded in changing one metal into
another.
~Number of elements.~ The number of substances now considered to be
elements is not large--about eighty in all. Many of these are rare, and
very few of them make any large fraction of the materials in the
earth's crust. Clarke gives the following estimate of the composition of
the earth's crust:
Oxygen 47.0% Calcium 3.5%
Silicon 27.9 Magnesium 2.5
Aluminium 8.1 Sodium 2.7
Iron 4.7 Potassium 2.4
Other elements 1.2%
A complete list of the elements is given in the Appendix. In this list
the more common of the elements are marked with an asterisk. It is not
necessary to study more than a third of the total number of elements to
gain a very good knowledge of chemistry.
~Physical state of the elements.~ About ten of the elements are gases at
ordinary temperatures. Two--mercury and bromine--are liquids. The others
are all solids, though their melting points vary through wide limits,
from cæsium which melts at 26° to elements which do not melt save in the
intense heat of the electric furnace.
~Occurrence of the elements.~ Comparatively few of the elements occur as
uncombined substances in nature, most of them being found in the form of
chemical compounds. When an element does occur by itself, as is the case
with gold, we say that it occurs in the _free state_ or _native_; when
it is combined with other substances in the form of compounds, we say
that it occurs in the _combined state_, or _in combination_. In the
latter case there is usually little about the compound to suggest that
the element is present in it; for we have seen that elements lose their
own peculiar properties when they enter into combination with other
elements. It would never be suspected, for example, that the reddish,
earthy-looking iron ore contains iron.
~Names of elements.~ The names given to the elements have been selected in
a great many different ways. (1) Some names are very old and their
original meaning is obscure. Such names are iron, gold, and copper. (2)
Many names indicate some striking physical property of the element. The
name bromine, for example, is derived from a Greek word meaning a
stench, referring to the extremely unpleasant odor of the substance. The
name iodine comes from a word meaning violet, alluding to the beautiful
color of iodine vapor. (3) Some names indicate prominent chemical
properties of the elements. Thus, nitrogen means the producer of niter,
nitrogen being a constituent of niter or saltpeter. Hydrogen means water
former, signifying its presence in water. Argon means lazy or inert, the
element being so named because of its inactivity. (4) Other elements are
named from countries or localities, as germanium and scandium.
~Symbols.~ In indicating the elements found in compounds it is
inconvenient to use such long names, and hence chemists have adopted a
system of abbreviations. These abbreviations are known as _symbols_,
each element having a distinctive symbol. (1) Sometimes the initial
letter of the name will suffice to indicate the element. Thus I stands
for iodine, C for carbon. (2) Usually it is necessary to add some other
characteristic letter to the symbol, since several names may begin with
the same letter. Thus C stands for carbon, Cl for chlorine, Cd for
cadmium, Ce for cerium, Cb for columbium. (3) Sometimes the symbol is an
abbreviation of the old Latin name. In this way Fe (ferrum) indicates
iron, Cu (cuprum), copper, Au (aurum), gold. The symbols are included in
the list of elements given in the Appendix. They will become familiar
through constant use.
~Chemical affinity the cause of chemical combination.~ The agency which
causes substances to combine and which holds them together when combined
is called _chemical affinity_. The experiments described in this
chapter, however, show that heat is often necessary to bring about
chemical action. The distinction between the cause producing chemical
action and the circumstances favoring it must be clearly made. Chemical
affinity is always the cause of chemical union. Many agencies may make
it possible for chemical affinity to act by overcoming circumstances
which stand in its way. Among these agencies are heat, light, and
electricity. As a rule, solution also promotes action between two
substances. Sometimes these agencies may overcome chemical attraction
and so occasion the decomposition of a compound.
EXERCISES
1. To what class of changes do the following belong? (a) The melting
of ice; (b) the souring of milk; (c) the burning of a candle; (d)
the explosion of gunpowder; (e) the corrosion of metals. What test
question must be applied in each of the above cases?
2. Give two additional examples (a) of chemical changes; (b) of
physical changes.
3. Is a chemical change always accompanied by a physical change? Is a
physical change always accompanied by a chemical change?
4. Give two or more characteristics of a chemical change.
5. (a) When a given weight of water freezes, does it absorb or evolve
heat? (b) When the resulting ice melts, is the total heat change the
same or different from that of freezing?
6. Give three examples of each of the following: (a) mechanical
mixtures; (b) chemical compounds; (c) elements.
7. Give the derivation of the names of the following elements: thorium,
gallium, selenium, uranium. (Consult dictionary.)
8. Give examples of chemical changes which are produced through the
agency of heat; of light; of electricity.
CHAPTER II
OXYGEN
~History.~ The discovery of oxygen is generally attributed to the English
chemist Priestley, who in 1774 obtained the element by heating a
compound of mercury and oxygen, known as red oxide of mercury. It is
probable, however, that the Swedish chemist Scheele had previously
obtained it, although an account of his experiments was not published
until 1777. The name oxygen signifies acid former. It was given to the
element by the French chemist Lavoisier, since he believed that all
acids owe their characteristic properties to the presence of oxygen.
This view we now know to be incorrect.
~Occurrence.~ Oxygen is by far the most abundant of all the elements. It
occurs both in the free and in the combined state. In the free state it
occurs in the air, 100 volumes of dry air containing about 21 volumes of
oxygen. In the combined state it forms eight ninths of water and nearly
one half of the rocks composing the earth's crust. It is also an
important constituent of the compounds which compose plant and animal
tissues; for example, about 66% by weight of the human body is oxygen.
~Preparation.~ Although oxygen occurs in the free state in the atmosphere,
its separation from the nitrogen and other gases with which it is mixed
is such a difficult matter that in the laboratory it has been found more
convenient to prepare it from its compounds. The most important of the
laboratory methods are the following:
1. _Preparation from water._ Water is a compound, consisting of 11.18%
hydrogen and 88.82% oxygen. It is easily separated into these
constituents by passing an electric current through it under suitable
conditions. The process will be described in the chapter on water. While
this method of preparation is a simple one, it is not economical.
2. _Preparation from mercuric oxide._ This method is of interest, since
it is the one which led to the discovery of oxygen. The oxide, which
consists of 7.4% oxygen and 92.6% mercury, is placed in a small, glass
test tube and heated. The compound is in this way decomposed into
mercury which collects on the sides of the glass tube, forming a silvery
mirror, and oxygen which, being a gas, escapes from the tube. The
presence of the oxygen is shown by lighting the end of a splint,
extinguishing the flame and bringing the glowing coal into the mouth of
the tube. The oxygen causes the glowing coal to burst into a flame.
In a similar way oxygen may be obtained from its compounds with
some of the other elements. Thus manganese dioxide, a black
compound of manganese and oxygen, when heated to about 700°,
loses one third of its oxygen, while barium dioxide, when
heated, loses one half of its oxygen.
3. _Preparation from potassium chlorate (usual laboratory method)._
Potassium chlorate is a white solid which consists of 31.9% potassium,
28.9% chlorine, and 39.2% oxygen. When heated it undergoes a series of
changes in which all the oxygen is finally set free, leaving a compound
of potassium and chlorine called potassium chloride. The change may be
represented as follows:
/potassium\
| | (potassium / potassium \ (potassium
{ chlorine } = { } + oxygen
| | chlorate) \ chlorine / chloride)
\oxygen /
[Illustration: JOSEPH PRIESTLEY (English) (1733-1804)
School-teacher, theologian, philosopher, scientist; friend of Benjamin
Franklin; discoverer of oxygen; defender of the phlogiston theory; the
first to use mercury in a pneumatic trough, by which means he first
isolated in gaseous form hydrochloric acid, sulphur dioxide, and
ammonia]
The evolution of the oxygen begins at about 400°. It has been found,
however, that if the potassium chlorate is mixed with about one fourth
its weight of manganese dioxide, the oxygen is given off at a much lower
temperature. Just how the manganese dioxide brings about this result is
not definitely known. The amount of oxygen obtained from a given weight
of potassium chlorate is exactly the same whether the manganese dioxide
is present or not. So far as can be detected the manganese dioxide
undergoes no change.
[Illustration: Fig. 4]
~Directions for preparing oxygen.~ The manner of preparing oxygen from
potassium chlorate is illustrated in the accompanying diagram (Fig. 4).
A mixture consisting of one part of manganese dioxide and four parts of
potassium chlorate is placed in the flask A and gently heated. The
oxygen is evolved and escapes through the tube B. It is collected by
bringing over the end of the tube the mouth of a bottle completely
filled with water and inverted in a vessel of water, as shown in the
figure. The gas rises in the bottle and displaces the water. In the
preparation of large quantities of oxygen, a copper retort (Fig. 5) is
often substituted for the glass flask.
[Illustration: Fig. 5]
In the preparation of oxygen from potassium chlorate and manganese
dioxide, the materials used must be pure, otherwise a violent explosion
may occur. The purity of the materials is tested by heating a small
amount of the mixture in a test tube.
~The collection of gases.~ The method used for collecting oxygen
illustrates the general method used for collecting such gases as are
insoluble in water or nearly so. The vessel C (Fig. 4), containing the
water in which the bottles are inverted, is called a _pneumatic trough._
~Commercial methods of preparation.~ Oxygen can now be purchased stored
under great pressure in strong steel cylinders (Fig. 6). It is prepared
either by heating a mixture of potassium chlorate and manganese dioxide,
or by separating it from the nitrogen and other gases with which it is
mixed in the atmosphere. The methods employed for effecting this
separation will be described in subsequent chapters.
[Illustration: Fig. 6]
~Physical properties.~ Oxygen is a colorless, odorless, tasteless gas,
slightly heavier than air. One liter of it, measured at a temperature of
0° and under a pressure of one atmosphere, weighs 1.4285 g., while under
similar conditions one liter of air weighs 1.2923 g. It is but slightly
soluble in water. Oxygen, like other gases, may be liquefied by applying
very great pressure to the highly cooled gas. When the pressure is
removed the liquid oxygen passes again into the gaseous state, since its
boiling point under ordinary atmospheric pressure is -182.5°.
~Chemical properties.~ At ordinary temperatures oxygen is not very active
chemically. Most substances are either not at all affected by it, or the
action is so slow as to escape notice. At higher temperatures, however,
it is very active, and unites directly with most of the elements. This
activity may be shown by heating various substances until just ignited
and then bringing them into vessels of the gas, when they will burn with
great brilliancy. Thus a glowing splint introduced into a jar of oxygen
bursts into flame. Sulphur burns in the air with a very weak flame and
feeble light; in oxygen, however, the flame is increased in size and
brightness. Substances which readily burn in air, such as phosphorus,
burn in oxygen with dazzling brilliancy. Even substances which burn in
air with great difficulty, such as iron, readily burn in oxygen.
The burning of a substance in oxygen is due to the rapid combination of
the substance or of the elements composing it with the oxygen. Thus,
when sulphur burns both the oxygen and sulphur disappear as such and
there is formed a compound of the two, which is an invisible gas, having
the characteristic odor of burning sulphur. Similarly, phosphorus on
burning forms a white solid compound of phosphorus and oxygen, while
iron forms a reddish-black compound of iron and oxygen.
~Oxidation.~ The term _oxidation_ is applied to the chemical change which
takes place when a substance, or one of its constituent parts, combines
with oxygen. This process may take place rapidly, as in the burning of
phosphorus, or slowly, as in the oxidation (or rusting) of iron when
exposed to the air. It is always accompanied by the liberation of heat.
The amount of heat liberated by the oxidation of a definite weight of
any given substance is always the same, being entirely independent of
the rapidity of the process. If the oxidation takes place slowly, the
heat is generated so slowly that it is difficult to detect it. If the
oxidation takes place rapidly, however, the heat is generated in such a
short interval of time that the substance may become white hot or burst
into a flame.
~Combustion; kindling temperature.~ When oxidation takes place so rapidly
that the heat generated is sufficient to cause the substance to glow or
burst into a flame the process is called _combustion_. In order that any
substance may undergo combustion, it is necessary that it should be
heated to a certain temperature, known as the _kindling temperature._
This temperature varies widely for different bodies, but is always
definite for the same body. Thus the kindling temperature of phosphorus
is far lower than that of iron, but is definite for each. When any
portion of a substance is heated until it begins to burn the combustion
will continue without the further application of heat, provided the heat
generated by the process is sufficient to bring other parts of the
substance to the kindling temperature. On the other hand, if the heat
generated is not sufficient to maintain the kindling temperature,
combustion ceases.
~Oxides.~ The compounds formed by the oxidation of any element are called
_oxides_. Thus in the combustion of sulphur, phosphorus, and iron, the
compounds formed are called respectively oxide of sulphur, oxide of
phosphorus, and oxide of iron. In general, then, _an oxide is a compound
of oxygen with another element_. A great many substances of this class
are known; in fact, the oxides of all the common elements have been
prepared, with the exception of those of fluorine and bromine. Some of
these are familiar compounds. Water, for example, is an oxide of
hydrogen, and lime an oxide of the metal calcium.
~Products of combustion.~ The particular oxides formed by the combustion
of any substance are called _products of combustion_ of that substance.
Thus oxide of sulphur is the product of the combustion of sulphur; oxide
of iron is the product of the combustion of iron. It is evident that the
products of the combustion of any substance must weigh more than the
original substance, the increase in weight corresponding to the amount
of oxygen taken up in the act of combustion. For example, when iron
burns the oxide of iron formed weighs more than the original iron.
In some cases the products of combustion are invisible gases, so that
the substance undergoing combustion is apparently destroyed. Thus, when
a candle burns it is consumed, and so far as the eye can judge nothing
is formed during combustion. That invisible gases are formed, however,
and that the weight of these is greater than the weight of the candle
may be shown by the following experiment.
[Illustration: Fig. 7]
A lamp chimney is filled with sticks of the compound known as
sodium hydroxide (caustic soda), and suspended from the beam of
the balance, as shown in Fig. 7. A piece of candle is placed on
the balance pan so that the wick comes just below the chimney,
and the balance is brought to a level by adding weights to the
other pan. The candle is then lighted. The products formed pass
up through the chimney and are absorbed by the sodium
hydroxide. Although the candle burns away, the pan upon which
it rests slowly sinks, showing that the combustion is attended
by an increase in weight.
~Combustion in air and in oxygen.~ Combustion in air and in
oxygen differs only in rapidity, the products formed being
exactly the same. That the process should take place less
rapidly in the former is readily understood, for the air is
only about one fifth oxygen, the remaining four fifths being
inert gases. Not only is less oxygen available, but much of the
heat is absorbed in raising the temperature of the inert gases
surrounding the substance undergoing combustion, and the
temperature reached in the combustion is therefore less.
~Phlogiston theory of combustion.~ The French chemist Lavoisier
(1743-1794), who gave to oxygen its name was the first to show
that combustion is due to union with oxygen. Previous to his
time combustion was supposed to be due to the presence of a
substance or principle called _phlogiston_. One substance was
thought to be more combustible than another because it
contained more phlogiston. Coal, for example, was thought to be
very rich in phlogiston. The ashes left after combustion would
not burn because all the phlogiston had escaped. If the
phlogiston could be restored in any way, the substance would
then become combustible again. Although this view seems absurd
to us in the light of our present knowledge, it formerly had
general acceptance. The discovery of oxygen led Lavoisier to
investigate the subject, and through his experiments he arrived
at the true explanation of combustion. The discovery of oxygen
together with the part it plays in combustion is generally
regarded as the most important discovery in the history of
chemistry. It marked the dawn of a new period in the growth of
the science.
~Combustion in the broad sense.~ According to the definition given above,
the presence of oxygen is necessary for combustion. The term is
sometimes used, however, in a broader sense to designate any chemical
change attended by the evolution of heat and light. Thus iron and
sulphur, or hydrogen and chlorine under certain conditions, will combine
so rapidly that light is evolved, and the action is called a combustion.
Whenever combustion takes place in the air, however, the process is one
of oxidation.
~Spontaneous combustion.~ The temperature reached in a given
chemical action, such as oxidation, depends upon the rate at
which the reaction takes place. This rate is usually increased
by raising the temperature of the substances taking part in the
action.
When a slow oxidation takes place under such conditions that
the heat generated is not lost by being conducted away, the
temperature of the substance undergoing oxidation is raised,
and this in turn hastens the rate of oxidation. The rise in
temperature may continue in this way until the kindling
temperature of the substance is reached, when combustion
begins. Combustion occurring in this way is called _spontaneous
combustion_.
Certain oils, such as the linseed oil used in paints, slowly
undergo oxidation at ordinary temperatures, and not
infrequently the origin of fires has been traced to the
spontaneous combustion of oily rags. The spontaneous combustion
of hay has been known to set barns on fire. Heaps of coal have
been found to be on fire when spontaneous combustion offered
the only possible explanation.
~Importance of oxygen.~ 1. Oxygen is essential to life. Among living
organisms only certain minute forms of plant life can exist without it.
In the process of respiration the air is taken into the lungs where a
certain amount of oxygen is absorbed by the blood. It is then carried to
all parts of the body, oxidizing the worn-out tissues and changing them
into substances which may readily be eliminated from the body. The heat
generated by this oxidation is the source of the heat of the body. The
small amount of oxygen which water dissolves from the air supports all
the varied forms of aquatic animals.
2. Oxygen is also essential to decay. The process of decay is really a
kind of oxidation, but it will only take place in the presence of
certain minute forms of life known as bacteria. Just how these assist in
the oxidation is not known. By this process the dead products of animal
and vegetable life which collect on the surface of the earth are slowly
oxidized and so converted into harmless substances. In this way oxygen
acts as a great purifying agent.
3. Oxygen is also used in the treatment of certain diseases in which the
patient is unable to inhale sufficient air to supply the necessary
amount of oxygen.
OZONE
~Preparation.~ When electric sparks are passed through oxygen or air a
small percentage of the oxygen is converted into a substance called
_ozone_, which differs greatly from oxygen in its properties. The same
change can also be brought about by certain chemical processes. Thus, if
some pieces of phosphorus are placed in a bottle and partially covered
with water, the presence of ozone may soon be detected in the air
contained in the bottle. The conversion of oxygen into ozone is attended
by a change in volume, 3 volumes of oxygen forming 2 volumes of ozone.
If the resulting ozone is heated to about 300°, the reverse change
takes place, the 2 volumes of ozone being changed back into 3 volumes of
oxygen. It is possible that traces of ozone exist in the atmosphere,
although its presence there has not been definitely proved, the tests
formerly used for its detection having been shown to be unreliable.
~Properties.~ As commonly prepared, ozone is mixed with a large excess of
oxygen. It is possible, however, to separate the ozone and thus obtain
it in pure form. The gas so obtained has the characteristic odor noticed
about electrical machines when in operation. By subjecting it to great
pressure and a low temperature, the gas condenses to a bluish liquid,
boiling at -119°. When unmixed with other gases ozone is very explosive,
changing back into oxygen with the liberation of heat. Its chemical
properties are similar to those of oxygen except that it is far more
active. Air or oxygen containing a small amount of ozone is now used in
place of oxygen in certain manufacturing processes.
~The difference between oxygen and ozone.~ Experiments show that in
changing oxygen into ozone no other kind of matter is either added to
the oxygen or withdrawn from it. The question arises then, How can we
account for the difference in their properties? It must be remembered
that in all changes we have to take into account _energy_ as well as
_matter_. By changing the amount of energy in a substance we change its
properties. That oxygen and ozone contain different amounts of energy
may be shown in a number of ways; for example, by the fact that the
conversion of ozone into oxygen is attended by the liberation of heat.
The passage of the electric sparks through oxygen has in some way
changed the energy content of the element and thus it has acquired new
properties. _Oxygen and ozone must, therefore, be regarded as identical
so far as the kind of matter of which they are composed is concerned.
Their different properties are due to their different energy contents._
~Allotropic states or forms of matter.~ Other elements besides oxygen may
exist in more than one form. These different forms of the same element
are called _allotropic states_ or _forms_ of the element. These forms
differ not only in physical properties but also in their energy
contents. Elements often exist in a variety of forms which look quite
different. These differences may be due to accidental causes, such as
the size or shape of the particles or the way in which the element was
prepared. Only such forms, however, as have different energy contents
are properly called allotropic forms.
MEASUREMENT OF GAS VOLUMES
~Standard conditions.~ It is a well-known fact that the volume occupied by
a definite weight of any gas can be altered by changing the temperature
of the gas or the pressure to which it is subjected. In measuring the
volume of gases it is therefore necessary, for the sake of accuracy, to
adopt some standard conditions of temperature and pressure. The
conditions agreed upon are (1) a temperature of 0°, and (2) a pressure
equal to the average pressure exerted by the atmosphere at the sea
level, that is, 1033.3 g. per square centimeter. These conditions of
temperature and pressure are known as the _standard conditions_, and
when the volume of a gas is given it is understood that the measurement
was made under these conditions, unless it is expressly stated
otherwise. For example, the weight of a liter of oxygen has been given
as 1.4285 g. This means that one liter of oxygen, measured at a
temperature of 0° and under a pressure of 1033.3 g. per square
centimeter, weighs 1.4285 g.
The conditions which prevail in the laboratory are never the standard
conditions. It becomes necessary, therefore, to find a way to calculate
the volume which a gas will occupy under standard conditions from the
volume which it occupies under any other conditions. This may be done in
accordance with the following laws.
~Law of Charles.~ This law expresses the effect which a change in the
temperature of a gas has upon its volume. It may be stated as follows:
_For every degree the temperature of a gas rises above zero the volume
of the gas is increased by 1/273 of the volume which it occupies at
zero; likewise for every degree the temperature of the gas falls below
zero the volume of the gas is decreased by 1/273 of the volume which it
occupies at zero, provided in both cases that the pressure to which the
gas is subjected remains constant._
If V represents the volume of gas at 0°, then the volume at 1° will be
V + 1/273 V; at 2° it will be V + 2/273 V; or, in general, the
volume v, at the temperature t, will be expressed by the formula
(1) v = V + t/273 V,
or (2) v = V(1 + (t/273)).
Since 1/273 = 0.00366, the formula may be written
(3) v = V(1 + 0.00366t).
Since the value of V (volume under standard conditions) is the one
usually sought, it is convenient to transpose the equation to the
following form:
(4) V = v/(1 + 0.00366t).
The following problem will serve as an illustration of the application
of this equation.
The volume of a gas at 20° is 750 cc.; find the volume it will occupy at
0°, the pressure remaining constant.
In this case, v = 750 cc. and t = 20. By substituting these values,
equation (4) becomes
V = 750/(1 + 0.00366 × 20) = 698.9 cc.
~Law of Boyle.~ This law expresses the relation between the volume
occupied by a gas and the pressure to which it is subjected. It may be
stated as follows: _The volume of a gas is inversely proportional to the
pressure under which it is measured, provided the temperature of the gas
remains constant._
If V represents the volume when subjected to a pressure P and v
represents its volume when the pressure is changed to p, then, in
accordance with the above law, V : v :: p : P, or VP = vp.
In other words, for a given weight of a gas the product of the numbers
representing its volume and the pressure to which it is subjected is a
constant.
Since the pressure of the atmosphere at any point is indicated by the
barometric reading, it is convenient in the solution of the problems to
substitute the latter for the pressure measured in grams per square
centimeter. The average reading of the barometer at the sea level is 760
mm., which corresponds to a pressure of 1033.3 g. per square centimeter.
The following problem will serve as an illustration of the application
of Boyle's law.
A gas occupies a volume of 500 cc. in a laboratory where the barometric
reading is 740 mm. What volume would it occupy if the atmospheric
pressure changed so that the reading became 750 mm.?
Substituting the values in the equation VP = vp, we have 500 × 740 =
v × 750, or v = 493.3 cc.
~Variations in the volume of a gas due to changes both in temperature and
pressure.~ Inasmuch as corrections must be made as a rule for both
temperature and pressure, it is convenient to combine the equations
given above for the corrections for each, so that the two corrections
may be made in one operation. The following equation is thus obtained:
(5) V_{s} = vp/(760(1 + 0.00366t)),
in which V_{s} represents the volume of a gas under standard
conditions and v, p, and t the volume, pressure, and temperature
respectively at which the gas was actually measured.
The following problem will serve to illustrate the application of this
equation.
A gas having a temperature of 20° occupies a volume of 500 cc. when
subjected to a pressure indicated by a barometric reading of 740 mm.
What volume would this gas occupy under standard conditions?
In this problem v = 500, p = 740, and t = 20. Substituting these
values in the above equation, we get
V_{s} = (500 × 740)/(760 (1 + 0.00366 × 20)) = 453.6 cc.
[Illustration: Fig. 8]
~Variations in the volume of a gas due to the pressure of aqueous vapor.~
In many cases gases are collected over water, as explained under the
preparation of oxygen. In such cases there is present in the gas a
certain amount of water vapor. This vapor exerts a definite pressure,
which acts in opposition to the atmospheric pressure and which therefore
must be subtracted from the latter in determining the effective pressure
upon the gas. Thus, suppose we wish to determine the pressure to which
the gas in tube A (Fig. 8) is subjected. The tube is raised or lowered
until the level of the water inside and outside the tube is the same.
The atmosphere presses down upon the surface of the water (as indicated
by the arrows), thus forcing the water upward within the tube with a
pressure equal to the atmospheric pressure. The full force of this
upward pressure, however, is not spent in compressing the gas within the
tube, for since it is collected over water it contains a certain amount
of water vapor. This water vapor exerts a pressure (as indicated by the
arrow within the tube) in opposition to the upward pressure. It is
plain, therefore, that the effective pressure upon the gas is equal to
the atmospheric pressure less the pressure exerted by the aqueous vapor.
The pressure exerted by the aqueous vapor increases with the
temperature. The figures representing the extent of this pressure (often
called the _tension of aqueous vapor_) are given in the Appendix. They
express the pressure or tension in millimeters of mercury, just as the
atmospheric pressure is expressed in millimeters of mercury.
Representing the pressure of the aqueous vapor by a, formula (5)
becomes
(6) V_{s} = v(p - a)/(760(1 + 0.00366t)).
The following problem will serve to illustrate the method of applying
the correction for the pressure of the aqueous vapor.
The volume of a gas measured over water in a laboratory where the
temperature is 20° and the barometric reading is 740 mm. is 500 cc. What
volume would this occupy under standard conditions?
The pressure exerted by the aqueous vapor at 20° (see table in Appendix)
is equal to the pressure exerted by a column of mercury 17.4 mm. in
height. Substituting the values of v, t, p, and a in formula
(6), we have
(6) V_{s} = 500(740 - 17.4)/(760(1 + 0.00366 × 20)) = 442.9 cc.
~Adjustment of tubes before reading gas volumes.~ In measuring the volumes
of gases collected in graduated tubes or other receivers, over a liquid
as illustrated in Fig. 8, the reading should be taken after raising or
lowering the tube containing the gas until the level of the liquid
inside and outside the tube is the same; for it is only under these
conditions that the upward pressure within the tube is the same as the
atmospheric pressure.
EXERCISES
1. What is the meaning of the following words? phlogiston, ozone,
phosphorus. (Consult dictionary.)
2. Can combustion take place without the emission of light?
3. Is the evolution of light always produced by combustion?
4. (a) What weight of oxygen can be obtained from 100 g. of water?
(b) What volume would this occupy under standard conditions?
5. (a) What weight of oxygen can be obtained from 500g. of mercuric
oxide? (b) What volume would this occupy under standard conditions?
6. What weight of each of the following compounds is necessary to
prepare 50 l. of oxygen? (a) water; (b) mercuric oxide; (c)
potassium chlorate.
7. Reduce the following volumes to 0°, the pressure remaining constant:
(a) 150 cc. at 10°; (b) 840 cc. at 273°.
8. A certain volume of gas is measured when the temperature is 20°. At
what temperature will its volume be doubled?
9. Reduce the following volumes to standard conditions of pressure, the
temperature remaining constant: (a) 200 cc. at 740 mm.; (b) 500 l.
at 380 mm.
10. What is the weight of 1 l. of oxygen when the pressure is 750 mm.
and the temperature 0°?
11. Reduce the following volumes to standard conditions of temperature
and pressure: (a) 340 cc. at 12° and 753 mm; (b) 500 cc. at 15° and
740 mm.
12. What weight of potassium chlorate is necessary to prepare 250 l. of
oxygen at 20° and 750 mm.?
13. Assuming the cost of potassium chlorate and mercuric oxide to be
respectively $0.50 and $1.50 per kilogram, calculate the cost of
materials necessary for the preparation of 50 l. of oxygen from each of
the above compounds.
14. 100 g. of potassium chlorate and 25 g. of manganese dioxide were
heated in the preparation of oxygen. What products were left in the
flask, and how much of each was present?
CHAPTER III
HYDROGEN
~Historical.~ The element hydrogen was first clearly recognized as a
distinct substance by the English investigator Cavendish, who in 1766
obtained it in a pure state, and showed it to be different from the
other inflammable airs or gases which had long been known. Lavoisier
gave it the name hydrogen, signifying water former, since it had been
found to be a constituent of water.
~Occurrence.~ In the free state hydrogen is found in the atmosphere, but
only in traces. In the combined state it is widely distributed, being a
constituent of water as well as of all living organisms, and the
products derived from them, such as starch and sugar. About 10% of the
human body is hydrogen. Combined with carbon, it forms the substances
which constitute petroleum and natural gas.
It is an interesting fact that while hydrogen in the free state
occurs only in traces on the earth, it occurs in enormous
quantities in the gaseous matter surrounding the sun and
certain other stars.
~Preparation from water.~ Hydrogen can be prepared from water by several
methods, the most important of which are the following.
1. _By the electric current._ As has been indicated in the preparation
of oxygen, water is easily separated into its constituents, hydrogen and
oxygen, by passing an electric current through it under certain
conditions.
2. _By the action of certain metals._ When brought into contact with
certain metals under appropriate conditions, water gives up a portion
or the whole of its hydrogen, its place being taken by the metal. In the
case of a few of the metals this change occurs at ordinary temperatures.
Thus, if a bit of sodium is thrown on water, an action is seen to take
place at once, sufficient heat being generated to melt the sodium, which
runs about on the surface of the water. The change which takes place
consists in the displacement of one half of the hydrogen of the water by
the sodium, and may be represented as follows:
_ _ _ _
| hydrogen | | sodium |
sodium + | hydrogen |(water) = | hydrogen |(sodium hydroxide) + hydrogen
|_oxygen _| |_oxygen _|
The sodium hydroxide formed is a white solid which remains dissolved in
the undecomposed water, and may be obtained by evaporating the solution
to dryness. The hydrogen is evolved as a gas and may be collected by
suitable apparatus.
Other metals, such as magnesium and iron, decompose water rapidly, but
only at higher temperatures. When steam is passed over hot iron, for
example, the iron combines with the oxygen of the steam, thus displacing
the hydrogen. Experiments show that the change may be represented as
follows:
_ _
| hydrogen | _ _ _ _
iron + | hydrogen |(water) = | iron |(iron oxide) + | hydrogen |
|_oxygen _| |_oxygen _| |_hydrogen_|
The iron oxide formed is a reddish-black compound, identical with that
obtained by the combustion of iron in oxygen.
~Directions for preparing hydrogen by the action of steam on
iron.~ The apparatus used in the preparation of hydrogen from
iron and steam is shown in Fig. 9. A porcelain or iron tube
B, about 50 cm. in length and 2 cm. or 3 cm. in diameter, is
partially filled with fine iron wire or tacks and connected as
shown in the figure. The tube B is heated, slowly at first,
until the iron is red-hot. Steam is then conducted through the
tube by boiling the water in the flask A. The hot iron
combines with the oxygen in the steam, setting free the
hydrogen, which is collected over water. The gas which first
passes over is mixed with the air previously contained in the
flask and tube, and is allowed to escape, _since a mixture of
hydrogen with oxygen or air explodes violently when brought in
contact with a flame_. It is evident that the flask A must be
disconnected from the tube before the heat is withdrawn.
That the gas obtained is different from air and oxygen may be
shown by holding a bottle of it mouth downward and bringing a
lighted splint into it. The hydrogen is ignited and burns with
an almost colorless flame.
[Illustration Fig. 9]
~Preparation from acids~ (_usual laboratory method_). While hydrogen can
be prepared from water, either by the action of the electric current or
by the action of certain metals, these methods are not economical and
are therefore but little used. In the laboratory hydrogen is generally
prepared from compounds known as acids, all of which contain hydrogen.
When acids are brought in contact with certain metals, the metals
dissolve and set free the hydrogen of the acid. Although this reaction
is a quite general one, it has been found most convenient in preparing
hydrogen by this method to use either zinc or iron as the metal and
either hydrochloric or sulphuric acid as the acid. Hydrochloric acid is
a compound consisting of 2.77% hydrogen and 97.23% chlorine, while
sulphuric acid consists of 2.05% hydrogen, 32.70% sulphur, and 65.25%
oxygen.
The changes which take place in the preparation of hydrogen from zinc
and sulphuric acid (diluted with water) may be represented as follows:
_ _ _ _
| hydrogen |(sulphuric | zinc |(zinc
zinc + | sulphur | acid) = | sulphur | sulphate) + hydrogen
|_oxygen _| |_oxygen _|
In other words, the zinc has taken the place of the hydrogen in
sulphuric acid. The resulting compound contains zinc, sulphur, and
oxygen, and is known as zinc sulphate. This remains dissolved in the
water present in the acid. It may be obtained in the form of a white
solid by evaporating the liquid left after the metal has passed into
solution.
When zinc and hydrochloric acid are used the following changes take
place:
_ _ _ _
| hydrogen |(hydrochloric | zinc |(zinc
zinc + |_chlorine_| acid) = |_chlorine_| chloride) + hydrogen
When iron is used the changes which take place are exactly similar to
those just given for zinc.
[Illustration Fig. 10.]
~Directions for preparing hydrogen from acids.~ The preparation
of hydrogen from acids is carried out in the laboratory as
follows: The metal is placed in a flask or wide-mouthed bottle
A (Fig. 10) and the acid is added slowly through the funnel
tube B. The metal dissolves in the acid, while the hydrogen
which is liberated escapes through the exit tube C and is
collected over water. It is evident that the hydrogen which
passes over first is mixed with the air from the bottle A.
Hence care must be taken not to bring a flame near the exit
tube, since, as has been stated previously, such a mixture
explodes with great violence when brought in contact with a
flame.
~Precautions.~ Both sulphuric acid and zinc, if impure, are
likely to contain small amounts of arsenic. Such materials
should not be used in preparing hydrogen, since the arsenic
present combines with a portion of the hydrogen to form a very
poisonous gas known as arsine. On the other hand, chemically
pure sulphuric acid, i.e. sulphuric acid that is entirely free
from impurities, will not act upon chemically pure zinc. The
reaction may be started, however, by the addition of a few
drops of a solution of copper sulphate or platinum
tetrachloride.
~Physical properties.~ Hydrogen is similar to oxygen in that it is a
colorless, tasteless, odorless gas. It is characterized by its extreme
lightness, being the lightest of all known substances. One liter of the
gas weighs only 0.08984 g. On comparing this weight with that of an
equal volume of oxygen, viz., 1.4285 g., the latter is found to be 15.88
times as heavy as hydrogen. Similarly, air is found to be 14.38 times as
heavy as hydrogen. Soap bubbles blown with hydrogen rapidly rise in the
air. On account of its lightness it is possible to pour it upward from
one bottle into another. Thus, if the bottle A (Fig. 11) is filled
with hydrogen, placed mouth downward by the side of bottle _B_, filled
with air, and is then gradually inverted under B as indicated in the
figure, the hydrogen will flow upward into bottle _B_, displacing the
air. Its presence in bottle B may then be shown by bringing a lighted
splint to the mouth of the bottle, when the hydrogen will be ignited by
the flame. It is evident, from this experiment, that in order to retain
the gas in an open bottle the bottle must be placed mouth downward.
[Illustration Fig. 11]
Hydrogen is far more difficult to liquefy than any other gas, with the
exception of helium, a rare element recently found to exist in the
atmosphere. The English scientist Dewar, however, in 1898 succeeded not
only in obtaining hydrogen in liquid state but also as a solid. Liquid
hydrogen is colorless and has a density of only 0.07. Its boiling point
under atmospheric pressure is -252°. Under diminished pressure the
temperature has been reduced to -262°. The solubility of hydrogen in
water is very slight, being still less than that of oxygen.
Pure hydrogen produces no injurious results when inhaled. Of course one
could not live in an atmosphere of the gas, since oxygen is essential to
respiration.
~Chemical properties.~ At ordinary temperatures hydrogen is not an active
element. A mixture of hydrogen and chlorine, however, will combine with
explosive violence at ordinary temperature if exposed to the sunlight.
The union can be brought about also by heating. The product formed in
either case is hydrochloric acid. Under suitable conditions hydrogen
combines with nitrogen to form ammonia, and with sulphur to form the
foul-smelling gas, hydrogen sulphide. The affinity of hydrogen for
oxygen is so great that a mixture of hydrogen and oxygen or hydrogen
and air explodes with great violence when heated to the kindling
temperature (about 612°). Nevertheless under proper conditions hydrogen
may be made to burn quietly in either oxygen or air. The resulting
hydrogen flame is almost colorless and is very hot. The combustion of
the hydrogen is, of course, due to its union with oxygen. The product of
the combustion is therefore a compound of hydrogen and oxygen. That this
compound is water may be shown easily by experiment.
[Illustration Fig. 12]
~Directions for burning hydrogen in air.~ The combustion of
hydrogen in air may be carried out safely as follows: The
hydrogen is generated in the bottle A (Fig. 12), is dried by
conducting it through the tube X, filled with some substance
(generally calcium chloride) which has a great attraction for
moisture, and escapes through the tube T, the end of which is
drawn out to a jet. The hydrogen first liberated mixes with the
air contained in the generator. If a flame is brought near the
jet before this mixture has all escaped, a violent and very
dangerous explosion results, since the entire apparatus is
filled with the explosive mixture. On the other hand, if the
flame is not applied until all the air has been expelled, the
hydrogen is ignited and burns quietly, since only the small
amount of it which escapes from the jet can come in contact
with the oxygen of the air at any one time. By holding a cold,
dry bell jar or bottle over the flame, in the manner shown in
the figure, the steam formed by the combustion of the hydrogen
is condensed, the water collecting in drops on the sides of the
jar.
~Precautions.~ In order to avoid danger it is absolutely necessary to
prove that the hydrogen is free from air before igniting it. This can be
done by testing small amounts of the escaping gas. A convenient and safe
method of doing this is to fill a test tube with the gas by inverting it
over the jet. The hydrogen, on account of its lightness, collects in the
tube, displacing the air. After holding it over the jet for a few
moments in order that it may be filled with the gas, the tube is gently
brought, mouth downward, to the flame of a burner placed not nearer than
an arm's length from the jet. If the hydrogen is mixed with air a slight
explosion occurs, but if pure it burns quietly in the tube. The
operation is repeated until the gas burns quietly, when the tube is
quickly brought back over the jet for an instant, whereby the escaping
hydrogen is ignited by the flame in the tube.
[Illustration. Fig. 13]
~A mixture of hydrogen and oxygen is explosive.~ That a mixture of
hydrogen and air is explosive may be shown safely as follows: A cork
through which passes a short glass tube about 1 cm. in diameter is
fitted air-tight into the tubule of a bell jar of 2 l. or 3 l. capacity.
(A thick glass bottle with bottom removed may be used.) The tube is
closed with a small rubber stopper and the bell jar filled with
hydrogen, the gas being collected over water. When entirely filled with
the gas the jar is removed from the water and supported by blocks of
wood in order to leave the bottom of the jar open, as shown in Fig. 13.
The stopper is now removed from the tube in the cork, and the hydrogen,
which on account of its lightness escapes from the tube, is at once
lighted. As the hydrogen escapes, the air flows in at the bottom of the
jar and mixes with the remaining portion of the hydrogen, so that a
mixture of the two soon forms, and a loud explosion results. The
explosion is not dangerous, since the bottom of the jar is open, thus
leaving room for the expansion of the hot gas.
Since air is only one fifth oxygen, the remainder being inert gases, it
may readily be inferred that a mixture of hydrogen with pure oxygen
would be far more explosive than a mixture of hydrogen with air. Such
mixtures should not be made except in small quantities and by
experienced workers.
~Hydrogen does not support combustion.~ While hydrogen is readily
combustible, it is not a supporter of combustion. In other words,
substances will not burn in it. This may be shown by bringing a lighted
candle supported by a stiff wire into a bottle or cylinder of the pure
gas, as shown in Fig. 14. The hydrogen is ignited by the flame of the
candle and burns at the mouth of the bottle, where it comes in contact
with the oxygen in the air. When the candle is thrust up into the gas,
its flame is extinguished on account of the absence of oxygen. If slowly
withdrawn, the candle is relighted as it passes through the layer of
burning hydrogen.
[Illustration: Fig. 14]
[Illustration: Fig. 15]
~Reduction.~ On account of its great affinity for oxygen, hydrogen has the
power of abstracting it from many of its compounds. Thus, if a stream of
hydrogen, dried by passing through the tube B (Fig. 15), filled with
calcium chloride, is conducted through the tube C containing some
copper oxide, heated to a moderate temperature, the hydrogen abstracts
the oxygen from the copper oxide. The change may be represented as
follows:
hydrogen + {copper} {hydrogen}
{oxygen}(copper oxide) = {oxygen }(water) + copper
The water formed collects in the cold portions of the tube C near its
end. In this experiment the copper oxide is said to undergo reduction.
_Reduction may therefore be defined as the process of withdrawing oxygen
from a compound._
~Relation of reduction to oxidation.~ At the same time that the copper
oxide is reduced it is clear that the hydrogen is oxidized, for it
combines with the oxygen given up by the copper oxide. The two processes
are therefore very closely related, and it usually happens that when one
substance is oxidized some other substance is reduced. That substance
which gives up its oxygen is called an _oxidizing agent_, while the
substance which unites with the oxygen is called a _reducing agent_.
~The oxyhydrogen blowpipe.~ This is a form of apparatus used for burning
hydrogen in pure oxygen. As has been previously stated, the flame
produced by the combustion of hydrogen in the air is very hot. It is
evident that if pure oxygen is substituted for air, the temperature
reached will be much higher, since there are no inert gases to absorb
the heat. The oxyhydrogen blowpipe, used to effect this combination,
consists of a small tube placed within a larger one, as shown in Fig.
16.
[Illustration: Fig. 16]
The hydrogen, stored under pressure, generally in steel cylinders, is
first passed through the outer tube and ignited at the open end of the
tube. The oxygen from a similar cylinder is then conducted through the
inner tube, and mixes with the hydrogen at the end of the tube. In order
to produce the maximum heat, the hydrogen and oxygen must be admitted to
the blowpipe in the exact proportion in which they combine, viz., 2
volumes of hydrogen to 1 of oxygen, or by weight, 1 part of hydrogen to
7.94 parts of oxygen. The intensity of the heat may be shown by bringing
into the flame pieces of metal such as iron wire or zinc. These burn
with great brilliancy. Even platinum, having a melting point of 1779°,
may be melted by the heat of the flame.
While the oxyhydrogen flame is intensely hot, it is almost non-luminous.
If directed against some infusible substance like ordinary lime (calcium
oxide), the heat is so intense that the lime becomes incandescent and
glows with a brilliant light. This is sometimes used as a source of
light, under the name of _Drummond_ or _lime light_.
[Illustration: Fig. 17]
~The blast lamp.~ A similar form of apparatus is commonly used in the
laboratory as a source of heat under the name _blast lamp_ (Fig. 17).
This differs from the oxyhydrogen blowpipe only in the size of the
tubes. In place of the hydrogen and oxygen the more accessible coal gas
and air are respectively used. The former is composed largely of a
mixture of free hydrogen and gaseous compounds of carbon and hydrogen.
While the temperature of the flame is not so high as that of the
oxyhydrogen blowpipe, it nevertheless suffices for most chemical
operations carried out in the laboratory.
~Uses of hydrogen.~ On account of its cost, hydrogen is but little used
for commercial purposes. It is sometimes used as a material for the
inflation of balloons, but usually the much cheaper coal gas is
substituted for it. Even hot air is often used when the duration of
ascension is very short. It has been used also as a source of heat and
light in the oxyhydrogen blowpipe. Where the electric current is
available, however, this form of apparatus has been displaced almost
entirely by the electric light and electric furnace, which are much more
economical and more powerful sources of light and heat.
EXERCISES
1. Will a definite weight of iron decompose an unlimited weight of
steam?
2. Why is oxygen passed through the inner tube of the oxyhydrogen
blowpipe rather than the outer?
3. In Fig. 14, will the flame remain at the mouth of the tube?
4. From Fig. 15, suggest a way for determining experimentally the
quantity of water formed in the reaction.
5. Distinguish clearly between the following terms: oxidation,
reduction, combustion, and kindling temperature.
6. Is oxidation always accompanied by reduction?
7. What is the source of heat in the lime light? What is the exact use
of lime in this instrument?
8. In Fig. 12, why is it necessary to dry the hydrogen by means of the
calcium chloride in the tube X?
9. At what pressure would the weight of 1 l. of hydrogen be equal to
that of oxygen under standard conditions?
10. (a) What weight of hydrogen can be obtained from 150 g. of
sulphuric acid? (b) What volume would this occupy under standard
conditions? (c) The density of sulphuric acid is 1.84. What volume
would the 150 g. of the acid occupy?
11. How many liters of hydrogen can be obtained from 50 cc. of sulphuric
acid having a density of 1.84?
12. Suppose you wish to fill five liter bottles with hydrogen, the gas
to be collected over water in your laboratory, how many cubic
centimeters of sulphuric acid would be required?
CHAPTER IV
COMPOUNDS OF HYDROGEN AND OXYGEN; WATER AND HYDROGEN DIOXIDE
WATER
~Historical.~ Water was long regarded as an element. In 1781 Cavendish
showed that it is formed by the union of hydrogen and oxygen. Being a
believer in the phlogiston theory, however, he failed to interpret his
results correctly. A few years later Lavoisier repeated Cavendish's
experiments and showed that water must be regarded as a compound of
hydrogen and oxygen.
~General methods employed for the determination of the composition of a
compound.~ The composition of a compound may be determined by either of
two general processes these are known as _analysis_ and _synthesis_.
1. _Analysis_ is the process of decomposing a compound into its
constituents and determining what these constituents are. The analysis
is _qualitative_ when it results in merely determining what elements
compose the compound; it is _quantitative_ when the exact percentage of
each constituent is determined. Qualitative analysis must therefore
precede quantitative analysis, for it must be known what elements, are
in a compound before a method can be devised for determining exactly how
much of each is present.
2. _Synthesis_ is the process of forming a compound from its constituent
parts. It is therefore the reverse of analysis. Like analysis, it may be
either qualitative or quantitative.
~Application of these methods to the determination of the composition of
water.~ The determination of the composition of water is a matter of
great interest not only because of the importance of the compound but
also because the methods employed illustrate the general methods of
analysis and synthesis.
~Methods based on analysis.~ The methods based on analysis may be either
qualitative or quantitative in character.
[Illustration: Fig. 18]
1. _Qualitative analysis._ As was stated in the study of oxygen, water
may be separated into its component parts by means of the electric
current. The form of apparatus ordinarily used for effecting this
analysis is shown in Fig. 18. A platinum wire, to the end of which is
attached a small piece of platinum foil (about 15 mm. by 25 mm.), is
fused through each of the tubes B and D, as shown in the figure. The
stopcocks at the ends of these tubes are opened and water, to which has
been added about one tenth of its volume of sulphuric acid, is poured
into the tube A until the side tubes B and D are completely
filled. The stopcocks are then closed. The platinum wires extending into
the tubes B and D are now connected with the wires leading from two
or three dichromate cells joined in series. The pieces of platinum foil
within the tubes thus become the electrodes, and the current flows from
one to the other through the acidulated water. As soon as the current
passes, bubbles of gas rise from each of the electrodes and collect in
the upper part of the tubes. The gas rising from the negative electrode
is found to be hydrogen, while that from the positive electrode is
oxygen. It will be seen that the volume of the hydrogen is approximately
double that of the oxygen. Oxygen is more soluble in water than
hydrogen, and a very little of it is also lost by being converted into
ozone and other substances. It has been found that when the necessary
corrections are made for the error due to these facts, the volume of the
hydrogen is exactly double that of the oxygen.
Fig. 19 illustrates a simpler form of apparatus, which may be used in
place of that shown in Fig. 18. A glass or porcelain dish is partially
filled with water to which has been added the proper amount of acid. Two
tubes filled with the same liquid are inverted over the electrodes. The
gases resulting from the decomposition of the water collect in the
tubes.
[Illustration: Fig. 19]
2. _Quantitative analysis._ The analysis just described is purely
qualitative and simply shows that water contains hydrogen and oxygen. It
does not prove the absence of other elements; indeed it does not prove
that the hydrogen and oxygen are present in the proportion in which they
are liberated by the electric current. The method may be made
quantitative, however, by weighing the water decomposed and also the
hydrogen and oxygen obtained in its decomposition. If the combined
weights of the hydrogen and oxygen exactly equal the weight of the water
decomposed, then it would be proved that the water consists of hydrogen
and oxygen in the proportion in which they are liberated by the electric
current. This experiment is difficult to carry out, however, so that the
more accurate methods based on synthesis are used.
~Methods based on synthesis.~ Two steps are necessary to ascertain the
exact composition of water by synthesis: (1) to show by qualitative
synthesis that water is formed by the union of oxygen with hydrogen; (2)
to determine by quantitative synthesis in what proportion the two
elements unite to form water. The fact that water is formed by the
combination of oxygen with hydrogen was proved in the preceding chapter.
The quantitative synthesis may be made as follows:
[Illustration: Fig. 20]
The combination of the two gases is brought about in a tube called a
eudiometer. This is a graduated tube about 60 cm. long and 2 cm. wide,
closed at one end (Fig. 20). Near the closed end two platinum wires are
fused through the glass, the ends of the wires within the tube being
separated by a space of 2 mm or 3 mm. The tube is entirely filled with
mercury and inverted in a vessel of the same liquid. Pure hydrogen is
passed into the tube until it is about one fourth filled. The volume of
the gas is then read off on the scale and reduced to standard
conditions. Approximately an equal volume of pure oxygen is then
introduced and the volume again read off and reduced to standard
conditions. This gives the total volume of the two gases. From this the
volume of the oxygen introduced may be determined by subtracting from
it the volume of the hydrogen. The combination of the two gases is now
brought about by connecting the two platinum wires with an induction
coil and passing a spark from one wire to the other. Immediately a
slight explosion occurs. The mercury in the tube is at first depressed
because of the expansion of the gases due to the heat generated, but at
once rebounds, taking the place of the gases which have combined to form
water. The volume of the water in the liquid state is so small that it
may be disregarded in the calculations. In order that the temperature of
the residual gas and the mercury may become uniform, the apparatus is
allowed to stand for a few minutes. The volume of the gas is then read
off and reduced to standard conditions, so that it may be compared with
the volumes of the hydrogen and oxygen originally taken. The residual
gas is then tested in order to ascertain whether it is hydrogen or
oxygen, experiments having proved that it is never a mixture of the two.
From the information thus obtained the composition of the water may be
calculated. Thus, suppose the readings were as follows:
Volume of hydrogen taken 20.3 cc.
Volume of hydrogen and oxygen 38.7
Volume of oxygen 18.4
Volume of gas left after combination has taken place (oxygen) 8.3
The 20.3 cc. of hydrogen have combined with 18.4 cc. minus 8.3 cc. (or
10.1 cc.) of oxygen; or approximately 2 volumes of hydrogen have
combined with 1 of oxygen. Since oxygen is 15.88 times as heavy as
hydrogen, the proportion by weight in which the two gases combine is 1
part of hydrogen to 7.94 of oxygen.
~Precaution.~ If the two gases are introduced into the eudiometer in the
exact proportions in which they combine, after the combination has taken
place the liquid will rise and completely fill the tube. Under these
conditions, however, the tube is very likely to be broken by the sudden
upward rush of the liquid. Hence in performing the experiment care is
taken to introduce an excess of one of the gases.
~A more convenient form of eudiometer.~ A form of eudiometer (Fig. 21)
different from that shown on page 43 is sometimes used to avoid the
calculations necessary in reducing the volumes of the gases to the same
conditions of temperature and pressure in order to make comparisons.
With this apparatus it is possible to take the readings of the volumes
under the same conditions of temperature and pressure, and thus compare
them directly. The apparatus (Fig. 21) is filled with mercury and the
gases introduced into the tube A. The experiment is carried out as in
the preceding one, except that before taking the reading of the gas
volumes, mercury is either added to the tube B or withdrawn from it by
means of the stopcock C, until it stands at exactly the same height in
both tubes. The gas inclosed in tube A is then under atmospheric
pressure; and since but a few minutes are required for performing the
experiment, the conditions of temperature and pressure may be regarded
as constant. Hence the volumes of the hydrogen and oxygen and of the
residual gas may be read off from the tube and directly compared.
[Illustration: Fig. 21]
~Method used by Berzelius and Dumas.~ The method used by these
investigators enables us to determine directly the proportion by weight
in which the hydrogen and oxygen combine. Fig. 22 illustrates the
apparatus used in making this determination. B is a glass tube
containing copper oxide. C and D are glass tubes filled with calcium
chloride, a substance which has great affinity for water. The tubes B
and C, including their contents, are carefully weighed, and the
apparatus connected as shown in the figure. A slow current of pure
hydrogen is then passed through A, and that part of the tube B which
contains copper oxide is carefully heated. The hydrogen combines with
the oxygen present in the copper oxide to form water, which is absorbed
by the calcium chloride in tube C. The calcium chloride in tube D
prevents any moisture entering tube C from the air. The operation is
continued until an appreciable amount of water has been formed. The
tubes B and C are then weighed once more. The loss of weight in the
tube B will exactly equal the weight of oxygen taken up from the
copper oxide in the formation of the water. The gain in weight in the
tube C will exactly equal the weight of the water formed. The
difference in these weights will of course equal the weight of the
hydrogen present in the water formed.
[Illustration: Fig. 22]
~Dumas' results.~ The above method for the determination of the
composition of water was first used by Berzelius in 1820. The work was
repeated in 1843 by Dumas, the average of whose results is as follows:
Weight of water formed 236.36 g.
Oxygen given up by the copper oxide 210.04
------
Weight of hydrogen present in water 26.32
According to this experiment the ratio of hydrogen to oxygen in water is
therefore 26.32 to 210.04, or as l to 7.98
~Morley's results.~ The American chemist Morley has recently determined
the composition of water, extreme precautions being taken to use pure
materials and to eliminate all sources of error. The hydrogen and oxygen
which combined, as well as the water formed, were all accurately
weighed. According to Morley's results, 1 part of hydrogen by weight
combines with 7.94 parts of oxygen to form water.
~Comparison of results obtained.~ From the above discussions it is easy to
see that it is by experiment alone that the composition of a compound
can be determined. Different methods may lead to slightly different
results. The more accurate the method chosen and the greater the skill
with which the experiment is carried out, the more accurate will be the
results. It is generally conceded by chemists that the results obtained
by Morley in reference to the composition of water are the most accurate
ones. In accordance with these results, then, _water must be regarded as
a compound containing hydrogen and oxygen in the proportion of 1 part by
weight of hydrogen to 7.94 parts by weight of oxygen_.
~Relation between the volume of aqueous vapor and the volumes of the
hydrogen and oxygen which combine to form it.~ When the quantitative
synthesis of water is carried out in the eudiometer as described above,
the water vapor formed by the union of the hydrogen and oxygen at once
condenses. The volume of the resulting liquid is so small that it may be
disregarded in making the calculations. If, however, the experiment is
carried out at a temperature of 100° or above, the water-vapor formed is
not condensed and it thus becomes possible to compare the volume of the
vapor with the volumes of hydrogen and oxygen which combined to form it.
This can be accomplished by surrounding the arm A of the eudiometer
(Fig. 23) with the tube B through which is passed the vapor obtained
by boiling some liquid which has a boiling point above 100°. In this way
it has been proved that 2 volumes of hydrogen and 1 volume of oxygen
combine to form exactly 2 volumes of water vapor, the volumes all being
measured under the same conditions of temperature and pressure. It will
be noted that the relation between these volumes may be expressed by
whole numbers. The significance of this very important fact will be
discussed in a subsequent chapter.
[Illustration: Fig. 23]
~Occurrence of water.~ Water not only covers about three fourths of the
surface of the earth, and is present in the atmosphere in the form of
moisture, but it is also a common constituent of the soil and rocks and
of almost every form of animal and vegetable organism. The human body is
nearly 70% water. This is derived not only from the water which we drink
but also from the food which we eat, most of which contains a large
percentage of water. Thus potatoes contain about 78% of water, milk 85%,
beef over 50%, apples 84%, tomatoes 94%.
~Impurities in water.~ Chemically pure water contains only hydrogen and
oxygen. Such a water never occurs in nature, however, for being a good
solvent, it takes up certain substances from the rocks and soil with
which it comes in contact. When such waters are evaporated these
substances are deposited in the form of a residue. Even rain water,
which is the purest form occurring in nature, contains dust particles
and gases dissolved from the atmosphere. The foreign matter in water is
of two kinds, namely, _mineral_, such as common salt and limestone, and
_organic_, that is the products of animal and vegetable life.
~Mineral matter in water.~ The amount and nature of the mineral
matter present in different waters vary greatly, depending on
the character of the rocks and soil with which the waters come
in contact. The more common of the substances present are
common salt and compounds of calcium, magnesium, and iron. One
liter of the average river water contains about 175 mg. of
mineral matter. Water from deep wells naturally contains more
mineral matter than river water, generally two or three times
as much, while sea water contains as much as 35,000 mg. to the
liter.
~Effect of impurities on health.~ The mineral matter in water does not,
save in very exceptional cases, render the water injurious to the human
system. In fact the presence of a certain amount of such matter is
advantageous, supplying the mineral constituents necessary for the
formation of the solid tissues of the body. The presence of organic
matter, on the other hand, must always be regarded with suspicion. This
organic matter may consist not only of the products of animal and
vegetable life but also of certain microscopic forms of living organisms
which are likely to accompany such products. Contagious diseases are
known to be due to the presence in the body of minute living organisms
or germs. Each disease is caused by its own particular kind of germ.
Through sewage these germs may find their way from persons afflicted
with disease into the water supply, and it is principally through the
drinking water that certain of these diseases, especially typhoid fever,
are spread. It becomes of great importance, therefore, to be able to
detect such matter when present in drinking water as well as to devise
methods whereby it can be removed or at least rendered harmless.
~Analysis of water.~ The mineral analysis of a water is, as the
name suggests, simply the determination of the mineral matter
present. Sanitary analysis, on the other hand, is the
determination of the organic matter present. The physical
properties of a water give no conclusive evidence as to its
purity, since a water may be unfit for drinking purposes and
yet be perfectly clear and odorless. Neither can any reliance
be placed on the simple methods often given for testing the
purity of water. Only the trained chemist can carry out such
methods of analysis as can be relied upon.
[Illustration: Fig. 24]
~Purification of water.~ Three general methods are used for the
purification of water, namely, _distillation_, _filtration_, and
_boiling_.
1. _Distillation._ The most effective way of purifying natural waters is
by the process of distillation. This consists in boiling the water and
condensing the steam. Fig. 24 illustrates the process of distillation,
as commonly conducted in the laboratory. Ordinary water is poured into
the flask A and boiled. The steam is conducted through the condenser
B, which consists essentially of a narrow glass tube sealed within a
larger one, the space between the two being filled with cold water,
which is admitted at C and escapes at D. The inner tube is thus kept
cool and the steam in passing through it is condensed. The water formed
by the condensation of the steam collects in the receiver E and is
known as _distilled_ water. Such water is practically pure, since the
impurities are nonvolatile and remain in the flask A.
~Commercial distillation.~ In preparing distilled water on a
large scale, the steam is generated in a boiler or other metal
container and condensed by passing it through a pipe made of
metal, generally tin. This pipe is wound into a spiral and is
surrounded by a current of cold water. Distilled water is used
by the chemist in almost all of his work. It is also used in
the manufacture of artificial ice and for drinking water.
~Fractional distillation.~ In preparing distilled water, it is
evident that if the natural water contains some substance which
is volatile its vapor will pass over and be condensed with the
steam, so that the distillate will not be pure water. Even such
mixtures, however, may generally be separated by repeated
distillation. Thus, if a mixture of water (boiling point 100°)
and alcohol (boiling point 78°) is distilled, the alcohol,
having the lower boiling point, tends to distill first,
followed by the water. The separation of the two is not
perfect, however, but may be made nearly so by repeated
distillations. The process of separating a mixture of volatile
substances by distillation is known as _fractional
distillation_.
2. _Filtration._ The process of distillation practically removes all
nonvolatile foreign matter, mineral as well as organic. In purifying
water for drinking purposes, however, it is only necessary to eliminate
the latter or to render it harmless. This is ordinarily done either by
filtration or boiling. In filtration the water is passed through some
medium which will retain the organic matter. Ordinary charcoal is a
porous substance and will condense within its pores the organic matter
in water if brought in contact with it. It is therefore well adapted to
the construction of filters. Such filters to be effective must be kept
clean, since it is evident that the charcoal is useless after its pores
are filled. A more effective type of filter is the Chamberlain-Pasteur
filter. In this the water is forced through a porous cylindrical cup,
the pores being so minute as to strain out the organic matter.
~City filtration beds.~ For purifying the water supply of cities,
large filtration beds are prepared from sand and gravel, and
the water is allowed to filter through these. Some of the
impurities are strained out by the filter, while others are
decomposed by the action of certain kinds of bacteria present
in the sand. Fig. 25 shows a cross section of a portion of the
filter used in purifying the water supply of Philadelphia. The
water filters through the sand and gravel and passes into the
porous pipe A, from which it is pumped into the city mains.
The filters are covered to prevent the water from freezing in
cold weather.
[Illustration: Fig. 25]
3. _Boiling._ A simpler and equally efficient method for purifying water
for drinking purposes consists in boiling the water. It is the germs in
water that render it dangerous to health. These germs are living forms
of matter. If the water is boiled, the germs are killed and the water
rendered safe. While these germs are destroyed by heat, cold has little
effect upon them. Thus Dewar, in working with liquid hydrogen, exposed
some of these minute forms of life to the temperature of boiling
hydrogen (-252°) without killing them.
~Self-purification of water.~ It has long been known that water
contaminated with organic matter tends to purify itself when exposed to
the air. This is due to the fact that the water takes up a small amount
of oxygen from the air, which gradually oxidizes the organic matter
present in the water. While water is undoubtedly purified in this way,
the method cannot be relied upon to purify a contaminated water so as to
render it safe for drinking purposes.
~Physical properties.~ Pure water is an odorless and tasteless liquid,
colorless in thin layers, but having a bluish tinge when observed
through a considerable thickness. It solidifies at 0° and boils at 100°
under the normal pressure of one atmosphere. If the pressure is
increased, the boiling point is raised. When water is cooled it steadily
contracts until the temperature of 4° is reached: it then expands. Water
is remarkable for its ability to dissolve other substances, and is the
best solvent known. Solutions of solids in water are more frequently
employed in chemical work than are the solid substances, for chemical
action between substances goes on more readily when they are in solution
than it does when they are in the solid state.
~Chemical properties.~ Water is a very stable substance, or, in other
words, it does not undergo decomposition readily. To decompose it into
its elements by heat alone requires a very high temperature; at 2500°,
for example, only about 5% of the entire amount is decomposed. Though
very stable towards heat, water can be decomposed in other ways, as by
the action of the electrical current or by certain metals.
~Heat of formation and heat of decomposition are equal.~ The fact
that a very high temperature is necessary to decompose water
into hydrogen and oxygen is in accord with the fact that a
great deal of heat is evolved by the union of hydrogen and
oxygen; for it has been proved that the heat necessary to
decompose a compound into its elements (heat of decomposition)
is equal to the heat evolved in the formation of a compound
from its elements (heat of formation).
~Water of crystallization.~ When a solid is dissolved in water and the
resulting solution is allowed to evaporate, the solid separates out,
often in the form of crystals. It has been found that the crystals of
many compounds, although perfectly dry, give up a definite amount of
water when heated, the substance at the same time losing its crystalline
form. Such water is called _water of crystallization_. This varies in
amount with different compounds, but is perfectly definite for the same
compound. Thus, if a perfectly dry crystal of copper sulphate is
strongly heated in a tube, water is evolved and condenses on the sides
of the tube, the crystal crumbling to a light powder. The weight of the
water evolved is always equal to exactly 36.07% of the weight of copper
sulphate crystals heated. The water must therefore be in chemical
combination with the substance composing the crystal; for if simply
mixed with it or adhering to it, not only would the substance appear
moist but the amount present would undoubtedly vary. The combination,
however, must be a very weak one, since the water is often expelled by
even a gentle heat. Indeed, in some cases the water is given up on
simple exposure to air. Such compounds are said to be _efflorescent_.
Thus a crystal of sodium sulphate (Glauber's salt) on exposure to air
crumbles to a fine powder, owing to the escape of its water of
crystallization. Other substances have just the opposite property: they
absorb moisture when exposed to the air. For example, if a bit of dry
calcium chloride is placed in moist air, in the course of a few hours it
will have absorbed sufficient moisture to dissolve it. Such substances
are said to be _deliquescent_. A deliquescent body serves as a good
drying or _desiccating_ agent. We have already employed calcium chloride
as an agent for absorbing the moisture from hydrogen. Many substances,
as for example quartz, form crystals which contain no water of
crystallization.
~Mechanically inclosed water.~ Water of crystallization must be
carefully distinguished from water which is mechanically
inclosed in a crystal and which can be removed by powdering the
crystal and drying. Thus, when crystals of common salt are
heated, the water inclosed in the crystal is changed into steam
and bursts the crystal with a crackling sound. Such crystals
are said to _decrepitate_. That this water is not combined is
proved by the fact that the amount present varies and that it
has all the properties of water.
~Uses of water.~ The importance of water in its relation to life and
commerce is too well known to require comment. Its importance to the
chemist has also been pointed out. It remains to call attention to the
fact that it is used as a standard in many physical measurements. Thus
0° and 100° on the centigrade scale are respectively the freezing and
the boiling points of water under normal pressure. The weight of 1 cc.
of water at its point of greatest density is the unit of weight in the
metric system, namely, the gram. It is also taken as the unit for the
determination of the density of liquids and solids as well as for the
measurement of amounts of heat.
HYDROGEN DIOXIDE
~Composition.~ As has been shown, 1 part by weight of hydrogen combines
with 7.94 parts by weight of oxygen to form water. It is possible,
however, to obtain a second compound of hydrogen and oxygen differing
from water in composition in that 1 part by weight of hydrogen is
combined with 2 × 7.94, or 15.88 parts, of oxygen. This compound is
called _hydrogen dioxide_ or _hydrogen peroxide_, the prefixes _di-_ and
_per-_ signifying that it contains more oxygen than hydrogen oxide,
which is the chemical name for water.
~Preparation.~ Hydrogen dioxide cannot be prepared cheaply by the direct
union of hydrogen and oxygen, and indirect methods must therefore be
used. It is commonly prepared by the action of a solution of sulphuric
acid on barium dioxide. The change which takes place may be indicated as
follows:
sulphuric acid + barium dioxide = barium sulphate + hydrogen dioxide
-------------- -------------- --------------- ----------------
hydrogen barium barium hydrogen
sulphur oxygen sulphur oxygen
oxygen oxygen
In other words, the barium and hydrogen in the two compounds exchange
places. By this method a dilute solution of the dioxide in water is
obtained. It is possible to separate the dioxide from the water by
fractional distillation. This is attended with great difficulties,
however, since the pure dioxide is explosive. The distillation is
carried on under diminished pressure so as to lower the boiling points
as much as possible; otherwise the high temperature would decompose the
dioxide.
~Properties.~ Pure hydrogen dioxide is a colorless sirupy liquid having a
density of 1.49. Its most characteristic property is the ease with which
it decomposes into water and oxygen. One part by weight of hydrogen is
capable of holding firmly only 7.94 parts of oxygen. The additional 7.94
parts of oxygen present in hydrogen dioxide are therefore easily
evolved, the compound breaking down into water and oxygen. This
decomposition is attended by the generation of considerable heat. In
dilute solution hydrogen dioxide is fairly stable, although such a
solution should be kept in a dark, cool place, since both heat and light
aid in the decomposition of the dioxide.
~Uses.~ Solutions of hydrogen dioxide are used largely as oxidizing
agents. The solution sold by druggists contains 3% of the dioxide and is
used in medicine as an antiseptic. Its use as an antiseptic depends upon
its oxidizing properties.
EXERCISES
1. Why does the chemist use distilled water in making solutions, rather
than filtered water?
2. How could you determine the total amount of solid matter dissolved in
a sample of water?
3. How could you determine whether a given sample of water is distilled
water?
4. How could the presence of air dissolved in water be detected?
5. How could the amount of water in a food such as bread or potato be
determined?
6. Would ice frozen from impure water necessarily be free from disease
germs?
7. Suppose that the maximum density of water were at 0° in place of 4°;
what effect would this have on the formation of ice on bodies of water?
8. Is it possible for a substance to contain both mechanically inclosed
water and water of crystallization?
9. If steam is heated to 2000° and again cooled, has any chemical change
taken place in the steam?
10. Why is cold water passed into C instead of D (Fig. 24)?
11. Mention at least two advantages that a metal condenser has over a
glass condenser.
12. Draw a diagram of the apparatus used in your laboratory for
supplying distilled water.
13. 20 cc. of hydrogen and 7 cc. of oxygen are placed in a eudiometer
and the mixture exploded. (a) How many cubic centimeters of aqueous
vapor are formed? (b) What gas and how much of it remains in excess?
14. (a) What weight of water can be formed by the combustion of 100 L
of hydrogen, measured under standard conditions? (b)What volume of
oxygen would be required in (a)? (c)What weight of potassium
chlorate is necessary to prepare this amount of oxygen?
15. What weight of oxygen is present in 1 kg. of the ordinary hydrogen
dioxide solution? In the decomposition of this weight of the dioxide
into water and oxygen, what volume of oxygen (measured under standard
conditions) is evolved?
CHAPTER V
THE ATOMIC THEORY
~Three fundamental laws of matter.~ Before we can gain any very definite
idea in regard to the structure of matter, and the way in which
different kinds of substances act chemically upon each other, it is
necessary to have clearly in view three fundamental laws of matter.
These laws have been established by experiment, and any conception which
may be formed concerning matter must therefore be in harmony with them.
The laws are as follows:
~Law of conservation of matter.~ This law has already been touched upon in
the introductory chapter, and needs no further discussion. It will be
recalled that it may be stated thus: _Matter can neither be created nor
destroyed, though it can be changed from one form into another._
~Law of definite composition.~ In the earlier days of chemistry there was
much discussion as to whether the composition of a given compound is
always precisely the same or whether it is subject to some variation.
Two Frenchmen, Berthollet and Proust, were the leaders in this
discussion, and a great deal of most useful experimenting was done to
decide the question. Their experiments, as well as all succeeding ones,
have shown that the composition of a pure chemical compound is always
exactly the same. Water obtained by melting pure ice, condensing steam,
burning hydrogen in oxygen, has always 11.18% hydrogen and 88.82% oxygen
in it. Red oxide of mercury, from whatever source it is obtained,
contains 92.6% mercury and 7.4% oxygen. This truth is known as _the law
of definite composition_, and may be stated thus: _The composition of a
chemical compound never varies._
~Law of multiple proportion.~ It has already been noted, however, that
hydrogen and oxygen combine in two different ratios to form water and
hydrogen dioxide respectively. It will be observed that this fact does
not contradict the law of definite composition, for entirely different
substances are formed. These compounds differ from each other in
composition, but the composition of each one is always constant. This
ability of two elements to unite in more than one ratio is very
frequently observed. Carbon and oxygen combine in two different ratios;
nitrogen and oxygen combine to form as many as five distinct compounds,
each with its own precise composition.
In the first decade of the last century John Dalton, an English
school-teacher and philosopher, endeavored to find some rule which holds
between the ratios in which two given substances combine. His studies
brought to light a very simple relation, which the following examples
will make clear. In water the hydrogen and oxygen are combined in the
ratio of 1 part by weight of hydrogen to 7.94 parts by weight of oxygen.
In hydrogen dioxide the 1 part by weight of hydrogen is combined with
15.88 parts by weight of oxygen. The ratio between the amounts of oxygen
which combine with the same amount of hydrogen to form water and
hydrogen dioxide respectively is therefore 7.94: 15.88, or 1: 2.
[Illustration: JOHN DALTON (English) (1766-1844)
Developed the atomic theory; made many studies on the properties and the
composition of gases. His book entitled "A New System of Chemical
Philosophy" had a large influence on the development of chemistry]
Similarly, the element iron combines with oxygen to form two oxides, one
of which is black and the other red. By analysis it has been shown that
the former contains 1 part by weight of iron combined with 0.286 parts
by weight of oxygen, while the latter contains 1 part by weight of iron
combined with 0.429 parts by weight of oxygen. Here again we find that
the amounts of oxygen which combine with the same fixed amount of iron
to form the two compounds are in the ratio of small whole numbers, viz.,
2:3.
Many other examples of this simple relation might be given, since it has
been found to hold true in all cases where more than one compound is,
formed from the same elements. Dalton's law of multiple proportion
states these facts as follows: _When any two elements,_ A _and_ B,
_combine to form more than one compound, the amounts of_ B _which unite
with any fixed amount of_ A _bear the ratio of small whole numbers to
each other_.
~Hypothesis necessary to explain the laws of matter.~ These three
generalizations are called _laws_, because they express in concise
language truths which are found by careful experiment to hold good in
all cases. They do not offer any explanation of the facts, but merely
state them. The human mind, however, does not rest content with the mere
bare facts, but seeks ever to learn the explanation of the facts. A
suggestion which is offered to explain such a set of facts is called an
_hypothesis_. The suggestion which Dalton offered to explain the three
laws of matter, called the _atomic hypothesis_, was prompted by his view
of the constitution of matter, and it involves three distinct
assumptions in regard to the nature of matter and chemical action.
Dalton could not prove these assumptions to be true, but he saw that if
they were true the laws of matter become very easy to understand.
~Dalton's atomic hypothesis.~ The three assumptions which Dalton made in
regard to the nature of matter, and which together constitute the atomic
hypothesis, are these:
1. All elements are made up of minute, independent particles which
Dalton designated as _atoms_.
2. All atoms of the same element have equal masses; those of different
elements have different masses; in any change to which an atom is
subjected its mass does not change.
3. When two or more elements unite to form a compound, the action
consists in the union of a definite small number of atoms of each
element to form a small particle of the compound. The smallest particles
of a given compound are therefore exactly alike in the number and kinds
of atoms which they contain, and larger masses of the substances are
simply aggregations of these least particles.
~Molecules and atoms.~ Dalton applied the name atom not only to the minute
particles of the elements but also to the least particles of compounds.
Later Avogadro, an Italian scientist, pointed out the fact that the two
are different, since the smallest particle of an element is a unit,
while that of a compound must have at least two units in it. He
suggested the name _molecule_ for the least particle of a compound which
can exist, retaining the name _atom_ for the smallest particle of an
element. In accordance with this distinction, we may define the atom and
the molecule as follows: _An atom is the smallest particle of an element
which can exist. A molecule is the smallest particle of a compound which
can exist._ It will be shown in a subsequent chapter that sometimes two
or more atoms of the same element unite with each other to form
molecules of the element. While the term atom, therefore, is applicable
only to elements, the term molecule is applicable both to elements and
compounds.
~The atomic hypothesis and the laws of matter.~ Supposing the atomic
hypothesis to be true, let us now see if it is in harmony with the laws
of matter.
1. _The atomic hypothesis and the law of conservation of matter._ It is
evident that if the atoms never change their masses in any change which
they undergo, the total quantity of matter can never change and the law
of conservation of matter must follow.
2. _The atomic hypothesis and the law of definite composition._
According to the third supposition, when iron combines with sulphur the
union is between definite numbers of the two kinds of atoms. In the
simplest case one atom of the one element combines with one atom of the
other. If the sulphur and the iron atoms never change their respective
masses when they unite to form a molecule of iron sulphide, all iron
sulphide molecules will have equal amounts of iron in them and also of
sulphur. Consequently any mass made up of iron sulphide molecules will
have the same fraction of iron by weight as do the individual iron
sulphide molecules. Iron sulphide, from whatever source, will have the
same composition, which is in accordance with the law of definite
composition.
3. _The atomic hypothesis and the law of multiple proportion._ But this
simplest case may not always be the only one. Under other conditions one
atom of iron might combine with two of sulphur to form a molecule of a
second compound. In such a case the one atom of iron would be in
combination with twice the mass of sulphur that is in the first
compound, since the sulphur atoms all have equal masses. What is true
for one molecule will be true for any number of them; consequently when
such quantities of these two compounds are selected as are found to
contain the same amount of iron, the one will contain twice as much
sulphur as the other.
The combination between the atoms may of course take place in other
simple ratios. For example, two atoms of one element might combine with
three or with five of the other. In all such cases it is clear that the
law of multiple proportion must hold true. For on selecting such numbers
of the two kinds of molecules as have the same number of the one kind of
atoms, the numbers of the other kind of atoms will stand in some simple
ratio to each other, and their weights will therefore stand in the same
simple ratio.
~Testing the hypothesis.~ Efforts have been made to find compounds which
do not conform to these laws, but all such attempts have resulted in
failure. If such compounds should be found, the laws would be no longer
true, and the hypothesis of Dalton would cease to possess value. When an
hypothesis has been tested in every way in which experiment can test it,
and is still found to be in harmony with the facts in the case, it is
termed a _theory_. We now speak of the atomic theory rather than of the
atomic hypothesis.
~Value of a theory.~ The value of a theory is twofold. It aids in the
clear understanding of the laws of nature because it gives an
intelligent idea as to why these laws should be in operation.
A theory also leads to discoveries. It usually happens that in testing a
theory much valuable work is done, and many new facts are discovered.
Almost any theory in explaining given laws will involve a number of
consequences apart from the laws it seeks to explain. Experiment will
soon show whether these facts are as the theory predicts they will be.
Thus Dalton's atomic theory predicted many properties of gases which
experiment has since verified.
~Atomic weights.~ It would be of great advantage in the study of chemistry
if we could determine the weights of the different kinds of atoms. It is
evident that this cannot be done directly. They are so small that they
cannot be seen even with a most powerful microscope. It is calculated
that it would take 200,000,000 hydrogen atoms placed side by side to
make a row one centimeter long. No balance can weigh such minute
objects. It is possible, however, to determine their relative
weights,--that is, how much heavier one is than another. _These relative
weights of the atoms are spoken of as the atomic weights of the
elements._
If elements were able to combine in only one way,--one atom of one with
one atom of another,--the problem of determining the atomic weights
would be very simple. We should merely have to take some one convenient
element as a standard, and find by experiment how much of each other
element would combine with a fixed weight of it. The ratios thus found
would be the same ratios as those between the atoms of the elements, and
thus we should have their relative atomic weights. The law of multiple
proportion calls attention to the fact that the atoms combine in other
ratios than 1: 1, and there is no direct way of telling which one, if
any, of the several compounds in a given case is the one consisting of a
single atom of each element.
If some way were to be found of telling how much heavier the entire
molecule of a compound is than the atom chosen as a standard,--that is,
of determining the molecular weights of compounds,--the problem could be
solved, though its solution would not be an entirely simple matter.
There are ways of determining the molecular weights of compounds, and
there are other experiments which throw light directly upon the relative
weights of the atoms. These methods cannot be described until the facts
upon which they rest have been studied. It will be sufficient for the
present to assume that these methods are trustworthy.
~Standard for atomic weights.~ Since the atomic weights are merely
relative to some one element chosen as a standard, it is evident that
any one of the elements may serve as this standard and that any
convenient value may be assigned to its atom. At one time oxygen was
taken as this standard, with the value 100, and the atomic weights of
the other elements were expressed in terms of this standard. It would
seem more rational to take the element of smallest atomic weight as the
standard and give it unit value; accordingly hydrogen was taken as the
standard with an atomic weight of 1. Very recently, however, this unit
has been replaced by oxygen, with an atomic weight of 16.
~Why oxygen is chosen as the standard for atomic weights.~ In the
determination of the atomic weight of an element it is necessary to find
the weight of the element which combines with a definite weight of
another element, preferably the element chosen as the standard. Since
oxygen combines with the elements far more readily than does hydrogen to
form definite compounds, it is far better adapted for the standard
element, and has accordingly replaced hydrogen as the standard. Any
definite value might be given to the weight of the oxygen atom. In
assigning a value to it, however, it is convenient to choose a whole
number, and as small a number as possible without making the atomic
weight of any other element less than unity. For these reasons the
number 16 has been chosen as the atomic weight of oxygen. This makes
the atomic weight of hydrogen equal to 1.008, so that there is but
little difference between taking oxygen as 16 and hydrogen as 1 for the
unit.
The atomic weights of the elements are given in the Appendix.
EXERCISES
1. Two compounds were found to have the following compositions: (a)
oxygen = 69.53%, nitrogen = 30.47%; (b) oxygen = 53.27%, nitrogen =
46.73%. Show that the law of multiple proportion holds in this case.
2. Two compounds were found to have the following compositions: (a)
oxygen = 43.64%, phosphorus = 56.36%; (b) oxygen = 56.35%, phosphorus
= 43.65%. Show that the law of multiple proportion holds in this case.
3. Why did Dalton assume that all the atoms of a given element have the
same weight?
CHAPTER VI
CHEMICAL EQUATIONS AND CALCULATIONS
~Formulas.~ Since the molecule of any chemical compound consists of a
definite number of atoms, and this number never changes without
destroying the identity of the compound, it is very convenient to
represent the composition of a compound by indicating the composition of
its molecules. This can be done very easily by using the symbols of the
atoms to indicate the number and the kind of the atoms which constitute
the molecule. HgO will in this way represent mercuric oxide, a molecule
of which has been found to contain 1 atom each of mercury and oxygen.
H_{2}O will represent water, the molecules of which consist of 1 atom of
oxygen and 2 of hydrogen, the subscript figure indicating the number of
the atoms of the element whose symbol precedes it. H_{2}SO_{4} will
stand for sulphuric acid, the molecules of which contain 2 atoms of
hydrogen, 1 of sulphur, and 4 of oxygen. The combination of symbols
which represents the molecule of a substance is called its _formula_.
~Equations.~ When a given substance undergoes a chemical change it is
possible to represent this change by the use of such symbols and
formulas. In a former chapter it was shown that mercuric oxide
decomposes when heated to form mercury and oxygen. This may be expressed
very briefly in the form of the equation
(1) HgO = Hg + O.
When water is electrolyzed two new substances, hydrogen and oxygen, are
formed from it. This statement in the form of an equation is
(2) H_{2}O = 2H + O.
The coefficient before the symbol for hydrogen indicates that a single
molecule of water yields two atoms of hydrogen on decomposition.
In like manner the combination of sulphur with iron is expressed by the
equation
(3) Fe + S = FeS.
The decomposition of potassium chlorate by heat takes place as
represented by the equation
(4) KClO_{3} = KCl + 3O.
~Reading of equations.~ Since equations are simply a kind of shorthand way
of indicating chemical changes which occur under certain conditions, in
reading an equation the full statement for which it stands should be
given. Equation (1) should be read, "Mercuric oxide when heated gives
mercury and oxygen"; equation (2) is equivalent to the statement, "When
electrolyzed, water produces hydrogen and oxygen"; equation (3), "When
heated together iron and sulphur unite to form iron sulphide"; equation
(4), "Potassium chlorate when heated yields potassium chloride and
oxygen."
~Knowledge required for writing equations.~ In order to write such
equations correctly, a considerable amount of exact knowledge is
required. Thus, in equation (1) the fact that red oxide of mercury has
the composition represented by the formula HgO, that it is decomposed by
heat, that in this decomposition mercury and oxygen are formed and no
other products,--all these facts must be ascertained by exact experiment
before the equation can be written. An equation expressing these facts
will then have much value.
Having obtained an equation describing the conduct of mercuric oxide on
being heated, it will not do to assume that other oxides will behave in
like manner. Iron oxide (FeO) resembles mercuric oxide in many respects,
but it undergoes no change at all when heated. Manganese dioxide, the
black substance used in the preparation of oxygen, has the formula
MnO_{2}. When this substance is heated oxygen is set free, but the metal
manganese is not liberated; instead, a different oxide of manganese
containing less oxygen is produced. The equation representing the
reaction is
3MnO_{2} = Mn_{3}O_{4} + 2O.
~Classes of reactions.~ When a chemical change takes place in a substance
the substance is said to undergo a reaction. Although a great many
different reactions will be met in the study of chemistry, they may all
be grouped under the following heads.
1. _Addition._ This is the simplest kind of chemical action. It consists
in the union of two or more substances to produce a new substance. The
combination of iron with sulphur is an example:
Fe + S = FeS.
2. _Decomposition._ This is the reverse of addition, the substance
undergoing reaction being parted into its constituents. The
decomposition of mercuric oxide is an example: HgO = Hg + O.
3. _Substitution._ It is sometimes possible for an element in the free
state to act upon a compound in such a way that it takes the place of
one of the elements of the compound, liberating it in turn. In the study
of the element hydrogen it was pointed out that hydrogen is most
conveniently prepared by the action of sulphuric or hydrochloric acid
upon zinc. When sulphuric acid is used a substance called zinc sulphate,
having the composition represented by the formula ZnSO_{4}, is formed
together with hydrogen. The equation is
Zn + H_{2}SO_{4} = ZnSO_{4} + 2H.
When hydrochloric acid is used zinc chloride and hydrogen are the
products of reaction:
Zn + 2HCl = ZnCl_{2} + 2H.
When iron is used in place of zinc the equation is
Fe + H_{2}SO_{4} = FeSO_{4} + 2H.
These reactions are quite similar, as is apparent from an examination of
the equations. In each case 1 atom of the metal replaces 2 atoms of
hydrogen in the acid, and the hydrogen escapes as a gas. When an element
in the free state, such as the zinc in the equations just given, takes
the place of some one element in a compound, setting it free from
chemical combination, the act is called _substitution_.
Other reactions illustrating substitution are the action of sodium on
water,
Na + H_{2}O = NaOH + H;
and the action of heated iron upon water,
3Fe + 4H_{2}O = Fe_{3}O_{4} + 8H.
4. _Double decomposition._ When barium dioxide (BaO_{2}) is treated with
sulphuric acid two compounds are formed, namely, hydrogen dioxide
(H_{2}O_{2}) and barium sulphate (BaSO_{4}). The equation is
BaO_{2} + H_{2}SO_{4} = BaSO_{4} + H_{2}O_{2}.
In this reaction it will be seen that the two elements barium and
hydrogen simply exchange places. Such a reaction is called a _double
decomposition_. We shall meet with many examples of this kind of
chemical reactions.
~Chemical equations are quantitative.~ The use of symbols and formulas in
expressing chemical changes has another great advantage. Thus, according
to the equation
H_{2}O = 2H + O,
1 molecule of water is decomposed into 2 atoms of hydrogen and 1 atom of
oxygen. But, as we have seen, the relative weights of the atoms are
known, that of hydrogen being 1.008, while that of oxygen is 16. The
molecule of water, being composed of 2 atoms of hydrogen and 1 atom of
oxygen, must therefore weigh relatively 2.016 + 16, or 18.016. The
amount of hydrogen in this molecule must be 2.016/18.016, or 11.18% of
the whole, while the amount of oxygen must be 16/18.018, or 88.82% of
the whole. Now, since any definite quantity of water is simply the sum
of a great many molecules of water, it is plain that the fractions
representing the relative amounts of hydrogen and oxygen present in a
molecule must likewise express the relative amounts of hydrogen and
oxygen present in any quantity of water. Thus, for example, in 20 g. of
water there are 2.016/18.016 × 20, or 2.238 g. of hydrogen, and
16/18.016 × 20, or 17.762 g. of oxygen. These results in reference to
the composition of water of course agree exactly with the facts obtained
by the experiments described in the chapter on water, for it is because
of those experiments that the values 1.008 and 16 are given to hydrogen
and oxygen respectively.
It is often easier to make calculations of this kind in the form of a
proportion rather than by fractions. Since the molecule of water and
the two atoms of hydrogen which it contains have the ratio by weight of
18.016: 2.016, any mass of water has the same ratio between its total
weight and the weight of the hydrogen in it. Hence, to find the number
of grams (x) of hydrogen in 20 g. of water, we have the proportion
18.016 : 2.016 :: 20 g. : x (grams of hydrogen).
Solving for x, we get 2.238 for the number of grams of hydrogen.
Similarly, to find the amount (x) of oxygen present in the 20 g. of
water, we have the proportion
18.016 : 16 :: 20 : x
from which we find that x = 17.762 g.
Again, suppose we wish to find what weight of oxygen can be obtained
from 15 g. of mercuric oxide. The equation representing the
decomposition of mercuric oxide is
HgO = Hg + O.
The relative weights of the mercury and oxygen atoms are respectively
200 and 16. The relative weight of the mercuric oxide molecule must
therefore be the sum of these, or 216. The molecule of mercuric oxide
and the atom of oxygen which it contains have the ratio 216: 16. This
same ratio must therefore hold between the weight of any given quantity
of mercuric oxide and that of the oxygen which it contains. Hence, to
find the weight of oxygen in 15 g. of mercuric oxide, we have the
proportion
216 : 16 :: 15 : x (grams of oxygen).
On the other hand, suppose we wish to prepare, say, 20 g. of oxygen. The
problem is to find out what weight of mercuric oxide will yield 20 g. of
oxygen. The following proportion evidently holds
216 : 16 :: x (grams of mercuric oxide) : 20;
from which we get x = 270.
In the preparation of hydrogen by the action of sulphuric acid upon
zinc, according to the equation,
Zn + H_{2}SO_{4} = ZnSO_{4} + 2 H,
suppose that 50 g. of zinc are available; let it be required to
calculate the weight of hydrogen which can be obtained. It will be seen
that 1 atom of zinc will liberate 2 atoms of hydrogen. The ratio by
weight of a zinc to an hydrogen atom is 65.4: 1.008; of 1 zinc atom to 2
hydrogen atoms, 65.4: 2.016. Zinc and hydrogen will be related in this
reaction in this same ratio, however many atoms of zinc are concerned.
Consequently in the proportion
65.4 : 2.016 :: 50 : x,
x will be the weight of hydrogen set free by 50 g. of zinc. The weight
of zinc sulphate produced at the same time can be found from the
proportion
65.4 : 161.46 :: 50 : x;
where 161.46 is the molecular weight of the zinc sulphate, and x the
weight of zinc sulphate formed. In like manner, the weight of sulphuric
acid used up can be calculated from the proportion
65.4 : 98.076 :: 50 : x.
These simple calculations are possible because the symbols and formulas
in the equations represent the relative weights of the substances
concerned in a chemical reaction. When once the relative weights of the
atoms have been determined, and it has been agreed to allow the symbols
to stand for these relative weights, an equation or formula making use
of the symbols becomes a statement of a definite numerical fact, and
calculations can be based on it.
~Chemical equations not algebraic.~ Although chemical equations are
quantitative, it must be clearly understood that they are not algebraic.
A glance at the equations
7 + 4 = 11, 8 + 5 = 9 + 4
will show at once that they are true. The equations
HgO = Hg + O, FeO = Fe + O
are equally true in an algebraic sense, but experiment shows that only
the first is true chemically, for iron oxide (FeO) cannot be directly
decomposed into iron and oxygen. Only such equations as have been found
by careful experiment to express a real chemical transformation, true
both for the kinds of substances as well as for the weights, have any
value.
_Chemical formulas and equations, therefore, are a concise way of
representing qualitatively and quantitatively facts which have been
found by experiment to be true in reference to the composition of
substances and the changes which they undergo._
~Formulas representing water of crystallization.~ An examination of
substances containing water of crystallization has shown that in every
case the water is present in such proportion by weight as can readily be
represented by a formula. For example, copper sulphate (CuSO_{4}) and
water combine in the ratio of 1 molecule of the sulphate to 5 of water;
calcium sulphate (CaSO_{4}) and water combine in the ratio 1: 2 to form
gypsum. These facts are expressed by writing the formulas for the two
substances with a period between them. Thus the formula for crystallized
copper sulphate is CuSO_{4}·5H_{2}O; that of gypsum is CaSO_{4}·2H_{2}O.
~Heat of reaction.~ Attention has frequently been directed to the fact
that chemical changes are usually accompanied by heat changes. In
general it has been found that in every chemical action heat is either
absorbed or given off. By adopting a suitable unit for the measurement
of heat, the heat change during a chemical reaction can be expressed in
the equation for the reaction.
Heat cannot be measured by the use of a thermometer alone, since the
thermometer measures the intensity of heat, not its quantity. The
easiest way to measure a quantity of heat is to note how warm it will
make a definite amount of a given substance chosen as a standard. Water
has been chosen as the standard, and the unit of heat is called a
_calorie. A calorie is defined as the amount of heat required to raise
the temperature of one gram of water one degree._
By means of this unit it is easy to indicate the heat changes in a given
chemical reaction. The equation
2H + O = H_{2}O + 68,300 cal.
means that when 2.016 g. of hydrogen combine with 16 g. of oxygen,
18.016 g. of water are formed and 68,300 cal. are set free.
C + 2S = CS_{2} - 19,000 cal.
means that an expenditure of 19,000 cal. is required to cause 12 g. of
carbon to unite with 64.12 g. of sulphur to form 76.12 g. of carbon
disulphide. In these equations it will be noted that the symbols stand
for as many grams of the substance as there are units in the weights of
the atoms represented by the symbols. This is always understood to be
the case in equations where the heat of reaction is given.
~Conditions of a chemical action are not indicated by equations.~
Equations do not tell the conditions under which a reaction will take
place. The equation
HgO = Hg + O
does not tell us that it is necessary to keep the mercuric oxide at a
high temperature in order that the decomposition may go on. The equation
Zn + 2HCl = ZnCl_{2} + 2H
in no way indicates the fact that the hydrochloric acid must be
dissolved in water before it will act upon the zinc. From the equation
H + Cl = HCl
it would not be suspected that the two gases hydrogen and chlorine will
unite instantly in the sunlight, but will stand mixed in the dark a long
time without change. It will therefore be necessary to pay much
attention to the details of the conditions under which a given reaction
occurs, as well as to the expression of the reaction in the form of an
equation.
EXERCISES
1. Calculate the percentage composition of the following substances:
(a) mercuric oxide; (b) potassium chlorate; (c) hydrochloric acid;
(d) sulphuric acid. Compare the results obtained with the compositions
as given in Chapters II and III.
2. Determine the percentage of copper, sulphur, oxygen, and water in
copper sulphate crystals. What weight of water can be obtained from 150
g. of this substance?
3. What weight of zinc can be dissolved in 10 g. of sulphuric acid? How
much zinc sulphate will be formed?
4. How many liters of hydrogen measured under standard conditions can be
obtained from the action of 8 g. of iron on 10 g. of sulphuric acid? How
much iron sulphate (FeSO_{4}) will be formed?
5. 10 g. of zinc were used in the preparation of hydrogen; what weight
of iron will be required to prepare an equal volume?
6. How many grams of barium dioxide will be required to prepare 1 kg. of
common hydrogen dioxide solution? What weight of barium sulphate will be
formed at the same time?
7. What weight of the compound Mn_{3}O_{4} will be formed by strongly
heating 25 g. of manganese dioxide? What volume of oxygen will be given
off at the same time, measured under standard conditions?
8. (a) What is the weight of 100 l. of hydrogen measured in a
laboratory in which the temperature is 20° and pressure 750 mm.? (b)
What weight of sulphuric acid is necessary to prepare this amount of
hydrogen? (c) The density of sulphuric acid is 1.84. Express the acid
required in (b) in cubic centimeters.
9. What weight of potassium chlorate is necessary to furnish sufficient
oxygen to fill four 200 cc. bottles in your laboratory (the gas to be
collected over water)?
CHAPTER VII
NITROGEN AND THE RARE ELEMENTS: ARGON, HELIUM, NEON, KRYPTON, XENON
~Historical.~ Nitrogen was discovered by the English chemist Rutherford in
1772. A little later Scheele showed it to be a constituent of air, and
Lavoisier gave it the name _azote_, signifying that it would not support
life. The name _nitrogen_ was afterwards given it because of its
presence in saltpeter or niter. The term azote and symbol Az are still
retained by the French chemists.
~Occurrence.~ Air is composed principally of oxygen and nitrogen in the
free state, about 78 parts by volume out of every 100 parts being
nitrogen. Nitrogen also occurs in nature in the form of potassium
nitrate (KNO_{3})--commonly called saltpeter or niter--as well as in
sodium nitrate (NaNO_{3}). Nitrogen is also an essential constituent of
all living organisms; for example, the human body contains about 2.4% of
nitrogen.
~Preparation from air.~ Nitrogen can be prepared from air by the action of
some substance which will combine with the oxygen, leaving the nitrogen
free. Such a substance must be chosen, however, as will combine with the
oxygen to form a product which is not a gas, and which can be readily
separated from the nitrogen. The substances most commonly used for this
purpose are phosphorus and copper.
1. _By the action of phosphorus._ The method used for the preparation of
nitrogen by the action of phosphorus is as follows:
The phosphorus is placed in a little porcelain dish, supported on a cork
and floated on water (Fig. 26). It is then ignited by contact with a hot
wire, and immediately a bell jar or bottle is brought over it so as to
confine a portion of the air. The phosphorus combines with the oxygen to
form an oxide of phosphorus, known as phosphorus pentoxide. This is a
white solid which floats about in the bell jar, but in a short time it
is all absorbed by the water, leaving the nitrogen. The withdrawal of
the oxygen is indicated by the rising of the water in the bell jar.
[Illustration: Fig. 26]
2. _By the action of copper._ The oxygen present in the air may also be
removed by passing air slowly through a heated tube containing copper.
The copper combines with the oxygen to form copper oxide, which is a
solid. The nitrogen passes on and may be collected over water.
~Nitrogen obtained from air is not pure.~ Inasmuch as air, in
addition to oxygen and nitrogen, contains small amounts of
other gases, and since the phosphorus as well as the copper
removes only the oxygen, it is evident that the nitrogen
obtained by these methods is never quite pure. About 1% of the
product is composed of other gases, from which it is very
difficult to separate the nitrogen. The impure nitrogen so
obtained may, however, be used for a study of most of the
properties of nitrogen, since these are not materially affected
by the presence of the other gases.
~Preparation from compounds of nitrogen.~ Pure nitrogen may be obtained
from certain compounds of the element. Thus, if heat is applied to the
compound ammonium nitrite (NH_{4}NO_{2}), the change represented in the
following equation takes place:
NH_{4}NO_{2} = 2H_{2}O + 2N.
~Physical properties.~ Nitrogen is similar to oxygen and hydrogen in that
it is a colorless, odorless, and tasteless gas. One liter of nitrogen
weighs 1.2501 g. It is almost insoluble in water. It can be obtained in
the form of a colorless liquid having a boiling point of -195° at
ordinary pressure. At -214° it solidifies.
~Chemical properties.~ Nitrogen is characterized by its inertness. It is
neither combustible nor a supporter of combustion. At ordinary
temperatures it will not combine directly with any of the elements
except under rare conditions. At higher temperatures it combines with
magnesium, lithium, titanium, and a number of other elements. The
compounds formed are called _nitrides_, just as compounds of an element
with oxygen are called _oxides_. When it is mixed with oxygen and
subjected to the action of electric sparks, the two gases slowly combine
forming oxides of nitrogen. A mixture of nitrogen and hydrogen when
treated similarly forms ammonia, a gaseous compound of nitrogen and
hydrogen. Since we are constantly inhaling nitrogen, it is evident that
it is not poisonous. Nevertheless life would be impossible in an
atmosphere of pure nitrogen on account of the exclusion of the necessary
oxygen.
~Argon, helium, neon, krypton, xenon.~ These are all rare
elements occurring in the air in very small quantities. Argon,
discovered in 1894, was the first one obtained. Lord Rayleigh,
an English scientist, while engaged in determining the exact
weights of various gases, observed that the nitrogen obtained
from the air is slightly heavier than pure nitrogen obtained
from its compounds. After repeating his experiments many times,
always with the same results, Rayleigh finally concluded that
the nitrogen which he had obtained from the air was not pure,
but was mixed with a small amount of some unknown gas, the
density of which is greater than that of nitrogen. Acting on
this assumption, Rayleigh, together with the English chemist
Ramsay, attempted to separate the nitrogen from the unknown
gas. Knowing that nitrogen would combine with magnesium, they
passed the nitrogen obtained from the air and freed from all
known substances through tubes containing magnesium heated to
the necessary temperature. After repeating this operation, they
finally succeeded in obtaining from the atmospheric nitrogen a
small volume of gas which would not combine with magnesium and
hence could not be nitrogen. This proved to be a new element,
to which they gave the name _argon_. As predicted, this new
element was found to be heavier than nitrogen, its density as
compared with hydrogen as a standard being approximately 20,
that of nitrogen being only 14. About 1% of the atmospheric
nitrogen proved to be argon. The new element is characterized
by having no affinity for other elements. Even under the most
favorable conditions it has not been made to combine with any
other element. On this account it was given the name argon,
signifying lazy or idle. Like nitrogen, it is colorless,
odorless, and tasteless. It has been liquefied and solidified.
Its boiling point is -187°.
Helium was first found in the gases expelled from certain
minerals by heating. Through the agency of the spectroscope it
had been known to exist in the sun long before its presence on
the earth had been demonstrated,--a fact suggested by the name
helium, signifying the sun. Its existence in traces in the
atmosphere has also been proven. It was first liquefied by
Onnes in July, 1908. Its boiling point, namely -269°, is the
lowest temperature yet reached.
The remaining elements of this group--neon, krypton, and
xenon--have been obtained from liquid air. When liquid air is
allowed to boil, the constituents which are the most difficult
to liquefy, and which therefore have the lowest boiling points,
vaporize first, followed by the others in the order of their
boiling points. It is possible in this way to make at least a
partial separation of the air into its constituents, and Ramsay
thus succeeded in obtaining from liquid air not only the known
constituents, including argon and helium, but also the new
elements, neon, krypton, and xenon. These elements, as well as
helium, all proved to be similar to argon in that they are
without chemical activity, apparently forming no compounds
whatever. The percentages present in the air are very small.
The names, neon, krypton, xenon, signify respectively, new,
hidden, stranger.
EXERCISES
1. How could you distinguish between oxygen, hydrogen, and nitrogen?
2. Calculate the relative weights of nitrogen and oxygen; of nitrogen
and hydrogen.
3. In the preparation of nitrogen from the air, how would hydrogen do as
a substance for the removal of the oxygen?
4. What weight of nitrogen can be obtained from 10 l. of air measured
under the conditions of temperature and pressure which prevail in your
laboratory?
5. How many grams of ammonium nitrite are necessary in the preparation
of 20 l. of nitrogen measured over water under the conditions of
temperature and pressure which prevail in your laboratory?
6. If 10 l. of air, measured under standard conditions, is passed over
100 g. of hot copper, how much will the copper gain in weight?
[Illustration: WILLIAM RAMSAY (Scotch) (1855-)
Has made many studies in the physical properties of substances;
discovered helium; together with Lord Rayleigh and others he discovered
argon, krypton, xenon, and neon; has contributed largely to the
knowledge of radio-active substances, showing that radium gradually
gives rise to helium; professor at University College, London]
CHAPTER VIII
THE ATMOSPHERE
~Atmosphere and air.~ The term _atmosphere_ is applied to the gaseous
envelope surrounding the earth. The term _air_ is generally applied to a
limited portion of this envelope, although the two words are often used
interchangeably. Many references have already been made to the
composition and properties of the atmosphere. These statements must now
be collected and discussed somewhat more in detail.
~Air formerly regarded as an element.~ Like water, air was at first
regarded as elementary in character. Near the close of the eighteenth
century Scheele, Priestley, and Lavoisier showed by their experiments
that it is a mixture of at least two gases,--those which we now call
oxygen and nitrogen. By burning substances in an inclosed volume of air
and noting the contraction in volume due to the removal of the oxygen,
they were able to determine with some accuracy the relative volumes of
oxygen and nitrogen present in the air.
~The constituents of the atmosphere.~ The constituents of the atmosphere
may be divided into two general groups: those which are essential to
life and those which are not essential.
1. _Constituents essential to life._ In addition to oxygen and nitrogen
at least two other substances, namely, carbon dioxide and water vapor,
must be present in the atmosphere in order that life may exist. The
former of these is a gaseous compound of carbon and oxygen having the
formula CO_{2}. Its properties will be discussed in detail in the
chapter on the compounds of carbon. Its presence in the air may be shown
by causing the air to bubble through a solution of calcium hydroxide
(Ca(OH)_{2}), commonly called lime water. The carbon dioxide combines
with the calcium hydroxide in accordance with the following equation:
Ca(OH)_{2} + CO_{2} = CaCO_{3} + H_{2}O.
The resulting calcium carbonate (CaCO_{3}) is insoluble in water and
separates in the form of a white powder, which causes the solution to
appear milky.
The presence of water vapor is readily shown by its condensation on cold
objects as well as by the fact that a bit of calcium chloride when
exposed to the air becomes moist, and may even dissolve in the water
absorbed from the air.
2. _Constituents not essential to life._ In addition to the essential
constituents, the air contains small percentages of various other gases,
the presence of which so far as is known is not essential to life. This
list includes the rare elements, argon, helium, neon, krypton, and
xenon; also hydrogen, ammonia, hydrogen dioxide, and probably ozone.
Certain minute forms of life (germs) are also present, the decay of
organic matter being due to their presence.
~Function of each of the essential constituents.~ (1) The oxygen
directly supports life through respiration. (2) The nitrogen,
on account of its inactivity, serves to dilute the oxygen, and
while contrary to the older views, it is possible that life
might continue to exist in the absence of the atmospheric
nitrogen, yet the conditions of life would be entirely changed.
Moreover, nitrogen is an essential constituent of all animal
and plant life. It was formerly supposed that neither animals
nor plants could assimilate the free nitrogen, but it has been
shown recently that the plants of at least one natural order,
the Leguminosæ, to which belong the beans, peas, and clover,
have the power of directly assimilating the free nitrogen from
the atmosphere. This is accomplished through the agency of
groups of bacteria, which form colonies in little tubercles on
the roots of the plants. These bacteria probably assist in the
absorption of nitrogen by changing the free nitrogen into
compounds which can be assimilated by the plant. Fig. 27 shows
the tubercles on the roots of a variety of bean. (3) The
presence of water vapor in the air is necessary to prevent
excessive evaporation from both plants and animals. (4) Carbon
dioxide is an essential plant food.
[Illustration: Fig. 27]
~The quantitative analysis of air.~ A number of different methods have
been devised for the determination of the percentages of the
constituents present in the atmosphere. Among these are the following.
1. _Determination of oxygen._ (1) The oxygen is withdrawn from a
measured volume of air inclosed in a tube, by means of phosphorus.
To make the determination, a graduated tube is filled with
water and inverted in a vessel of water. Air is introduced into
the tube until it is partially filled with the gas. The volume
of the inclosed air is carefully noted and reduced to standard
conditions. A small piece of phosphorus is attached to a wire
and brought within the tube as shown in Fig. 28. After a few
hours the oxygen in the inclosed air will have combined with
the phosphorus, the water rising to take its place. The
phosphorus is removed and the volume is again noted and reduced
to standard conditions. The contraction in the volume of the
air is equal to the volume of oxygen absorbed.
[Illustration: Fig. 28]
(2) The oxygen may also be estimated by passing a measured volume of air
through a tube containing copper heated to a high temperature. The
oxygen in the air combines with the copper to form copper oxide (CuO).
Hence the increase in the weight of the copper equals the weight of the
oxygen in the volume of air taken.
(3) A more accurate method is the following. A eudiometer tube is filled
with mercury and inverted in a vessel of the same liquid. A convenient
amount of air is then introduced into the tube and its volume accurately
noted. There is then introduced more than sufficient hydrogen to combine
with the oxygen present in the inclosed air, and the volume is again
accurately noted. The mixture is then exploded by an electric spark, and
the volume is once more taken. By subtracting this volume from the total
volume of the air and hydrogen there is obtained the contraction in
volume due to the union of the oxygen and hydrogen. The volume occupied
by the water formed by the union of the two gases is so small that it
may be disregarded in the calculation. Since oxygen and hydrogen combine
in the ratio 1: 2 by volume, it is evident that the contraction in
volume due to the combination is equal to the volume occupied by the
oxygen in the air contained in the tube, plus twice this volume of
hydrogen. In other words, one third of the total contraction is equal to
the volume occupied by the oxygen in the inclosed air. The following
example will make this clear:
Volume of air in tube 50.0 cc.
Volume after introducing hydrogen 80.0
Volume after combination of oxygen and hydrogen 48.5
Contraction in volume due to combination (80 cc.-48.5 cc.) 31.5
Volume of oxygen in 50 cc. of air (1/3 of 31.5) 10.5
All these methods agree in showing that 100 volumes of dry air contain
approximately 21 volumes of oxygen.
2. _Determination of nitrogen._ If the gas left after the removal of
oxygen from a portion of air is passed over heated magnesium, the
nitrogen is withdrawn, argon and the other rare elements being left. It
may thus be shown that of the 79 volumes of gas left after the removal
of the oxygen from 100 volumes of air, approximately 78 are nitrogen and
0.93 argon. The other elements are present in such small quantities that
they may be neglected.
3. _Determination of carbon dioxide._ The percentage of carbon dioxide
in any given volume of air may be determined by passing the air over
calcium hydroxide or some other compound which will combine with the
carbon dioxide. The increase in the weight of the hydroxide equals the
weight of the carbon dioxide absorbed. The amount present in the open
normal air is from 3 to 4 parts by volume in 10,000 volumes of air, or
about 0.04%.
4. _Determination of water vapor._ The water vapor present in a given
volume of air may be determined by passing the air over calcium chloride
(or some other compound which has a strong affinity for water), and
noting the increase in the weight of the chloride. The amount present
varies not only with the locality, but there is a wide variation from
day to day in the same locality because of the winds and changes in
temperature.
~Processes affecting the composition of the air.~ The most important of
these processes are the following.
1. _Respiration._ In the process of respiration some of the oxygen in
the inhaled air is absorbed by the blood and carried to all parts of the
body, where it combines with the carbon of the worn-out tissues. The
products of oxidation are carried back to the lungs and exhaled in the
form of carbon dioxide. The amount exhaled by an adult averages about 20
l. per hour. Hence in a poorly ventilated room occupied by a number of
people the amount of carbon dioxide rapidly increases. While this gas is
not poisonous unless present in large amounts, nevertheless air
containing more than 15 parts in 10,000 is not fit for respiration.
2. _Combustion._ All of the ordinary forms of fuel contain large
percentages of carbon. On burning, this carbon combines with oxygen in
the air, forming carbon dioxide. Combustion and respiration, therefore,
tend to diminish the amount of oxygen in the air and to increase the
amount of carbon dioxide.
3. _Action of plants._ Plants have the power, when in the sunlight, of
absorbing carbon dioxide from the air, retaining the carbon and
returning at least a portion of the oxygen to the air. It will be
observed that these changes are just the opposite of those brought about
by the processes of respiration and combustion.
~Poisonous effect of exhaled air.~ The differences in the
percentages of oxygen, carbon dioxide, and moisture present in
inhaled air and exhaled air are shown in the following
analyses.
INHALED AIR EXHALED AIR
Oxygen 21.00% 16.00%
Carbon dioxide 0.04 4.38
Moisture variable saturated
The foul odor of respired air is due to the presence of a
certain amount of organic matter. It is possible that this
organic matter rather than the carbon dioxide is responsible
for the injurious effects which follow the respiration of
impure air. The extent of such organic impurities present may
be judged, however, by the amount of carbon dioxide present,
since the two are exhaled together.
~The cycle of carbon in nature.~ Under the influence of sunlight,
the carbon dioxide absorbed from the air by plants reacts with
water and small amounts of other substances absorbed from the
soil to form complex compounds of carbon which constitute the
essential part of the plant tissue. This reaction is attended
by the evolution of oxygen, which is restored to the air. The
compounds resulting from these changes are much richer in their
energy content than are the substances from which they are
formed; hence a certain amount of energy must have been
absorbed in their formation. The source of this energy is the
sun's rays.
If the plant is burned, the changes which took place in the
formation of the compounds present are largely reversed. The
carbon and hydrogen present combine with oxygen taken from the
air to form carbon dioxide and water, while the energy absorbed
from the sun's rays is liberated in the form of energy of heat.
If, on the other hand, the plant is used as food, the compounds
present are used in building up the tissues of the body. When
this tissue breaks down, the changes which it undergoes are
very similar to those which take place when the plant is
burned. The carbon and hydrogen combine with the inhaled oxygen
to form carbon dioxide and water, which are exhaled. The energy
possessed by the complex substances is liberated partly in the
form of energy of heat, which maintains the heat of the body,
and partly in the various forms of muscular energy. The carbon
originally absorbed from the air by the plant in the form of
carbon dioxide is thus restored to the air and is ready to
repeat the cycle of changes.
~The composition of the air is constant.~ Notwithstanding the changes
constantly taking place which tend to alter the composition of the air,
the results of a great many analyses of air collected in the open fields
show that the percentages of oxygen and nitrogen as well as of carbon
dioxide are very nearly constant. Indeed, so constant are the
percentages of oxygen and nitrogen that the question has arisen, whether
these two elements are not combined in the air, forming a definite
chemical compound. That the two are not combined but are simply mixed
together can be shown in a number of ways, among which are the
following.
1. When air dissolves in water it has been found that the ratio of
oxygen to nitrogen in the dissolved air is no longer 21: 78, but more
nearly 35: 65. If it were a chemical compound, the ratio of oxygen to
nitrogen would not be changed by solution in water.
2. A chemical compound in the form of a liquid has a definite boiling
point. Water, for example, boils at 100°. Moreover the steam which is
thus formed has the same composition as the water. The boiling point of
liquid air, on the other hand, gradually rises as the liquid boils, the
nitrogen escaping first followed by the oxygen. If the two were
combined, they would pass off together in the ratio in which they are
found in the air.
~Why the air has a constant composition.~ If air is a mixture and changes
are constantly taking place which tend to modify its composition, how,
then, do we account for the constancy of composition which the analyses
reveal? This is explained by several facts. (1) The changes which are
caused by the processes of combustion and respiration, on the one hand,
and the action of plants, on the other, tend to equalize each other. (2)
The winds keep the air in constant motion and so prevent local changes.
(3) The volume of the air is so vast and the changes which occur are so
small compared with the total amount of air that they cannot be readily
detected. (4) Finally it must be noted that only air collected in the
open fields shows this constancy in composition. The air in a poorly
ventilated room occupied by a number of people rapidly changes in
composition.
~The properties of the air.~ Inasmuch as air is composed principally of a
mixture of oxygen and nitrogen, which elements have already been
discussed, its properties may be inferred largely from those of the two
gases. One liter weighs 1.2923 g. It is thus 14.38 times as heavy as
hydrogen. At the sea level it exerts an average pressure sufficient to
sustain a column of mercury 760 mm. in height. This is taken as the
standard pressure in determining the volumes of gases as well as the
boiling points of liquids. Water may be made to boil at any temperature
between 0° and considerably above 100° by simply varying the pressure.
It is only when the pressure upon it is equal to the normal pressure of
the atmosphere at the sea level, as indicated by a barometric reading of
760 mm., that it boils at 100°.
~Preparation of liquid air.~ Attention has been called to the fact that
both oxygen and nitrogen can be obtained in the liquid state by strongly
cooling the gases and applying great pressure to them. Since air is
largely a mixture of these two gases, it can be liquefied by the same
methods.
The methods for liquefying air have been simplified greatly in
that the low temperature required is obtained by allowing a
portion of the compressed air to expand. The expansion of a gas
is always attended by the absorption of heat. In liquefying air
the apparatus is so constructed that the heat absorbed is
withdrawn from air already under great pressure. This process
is continued until the temperature is lowered to the point of
liquefaction.
[Illustration: Fig. 29]
~The Dewar bulb.~ It is not possible to preserve air in the liquid state
in a closed vessel, on account of the enormous pressure exerted by it in
its tendency to pass into the gaseous state. It may however be preserved
for some hours or even days before it will completely evaporate, by
simply placing it in an open vessel surrounded by a nonconducting
material. The most efficient vessel for this purpose is the _Dewar bulb_
shown in Fig. 29. The air is withdrawn from the space between the two
walls, thus making it nonconducting.
~Properties and uses of liquid air.~ When first prepared, liquid air is
cloudy because of the presence of particles of solid carbon dioxide.
These may be filtered off, leaving a liquid of slightly bluish color. It
begins to boil at about -190°, the nitrogen passing off first, gradually
followed by the oxygen, the last portions being nearly pure oxygen. To a
certain extent oxygen is now prepared in this way for commercial
purposes.
The extremely low temperature of liquid air may be inferred from the
fact that mercury when cooled by it is frozen to a mass so hard that it
may be used for driving nails.
Liquid air is used in the preparation of oxygen and as a cooling agent
in the study of the properties of matter at low temperatures. It has
thus been found that elements at extremely low temperatures largely lose
their chemical activity.
EXERCISES
1. When oxygen and nitrogen are mixed in the proportion in which they
exist in the atmosphere, heat is neither evolved nor absorbed by the
process. What important point does this suggest?
2. What essential constituent of the air is found in larger amount in
manufacturing districts than in the open country?
3. Can you suggest any reason why the growth of clover in a field
improves the soil?
4. Why are the inner walls of a Dewar bulb sometimes coated with a film
of silver?
5. To what is the blue color of liquid air due? Does this color increase
in intensity on standing?
6. When ice is placed in a vessel containing liquid air, the latter
boils violently. Explain.
7. Taking the volumes of the oxygen and nitrogen in 100 volumes of air
as 21 and 78 respectively, calculate the percentages of these elements
present by weight.
8. Would combustion be more intense in liquid air than in the gaseous
substance?
9. A tube containing calcium chloride was found to weigh 30.1293 g. A
volume of air which weighed 15.2134 g. was passed through, after which
the weight of the tube was found to be 30.3405 g. What was the
percentage amount of moisture present in the air?
10. 10 l. of air measured at 20° and 740 mm. passed through lime water
caused the precipitation of 0.0102 g. of CaCO_{3}. Find the number of
volumes of carbon dioxide in 10,000 volumes of the air.
CHAPTER IX
SOLUTIONS
~Definitions.~ When a substance disappears in a liquid in such a way as to
thoroughly mix with it and to be lost to sight as an individual body,
the resulting liquid is called a _solution_. The liquid in which the
substance dissolves is called the _solvent_, while the dissolved
substance is called the _solute_.
~Classes of solutions.~ Matter in any one of its physical states may
dissolve in a liquid, so that we may have solutions of gases, of
liquids, and of solids. Solutions of liquids in liquids are not often
mentioned in the following pages, but the other two classes will become
very familiar in the course of our study, and deserve special attention.
SOLUTION OF GASES IN LIQUIDS
[Illustration: Fig. 30]
It has already been stated that oxygen, hydrogen, and nitrogen are
slightly soluble in water. Accurate study has led to the conclusion that
all gases are soluble to some extent not only in water but in many other
liquids. The amount of a gas which will dissolve in a liquid depends
upon a number of conditions, and these can best be understood by
supposing a vessel B (Fig. 30), to be filled with the gas and inverted
over the liquid. Under these circumstances the gas cannot escape or
become mixed with another gas.
~Circumstances affecting the solubility of gases.~ A number of
circumstances affect the solubility of a gas in a liquid.
1. _Nature of the gas._ Other conditions being equal, each gas has its
own peculiar solubility, just as it has its own special taste or odor.
The solubility of gases varies between wide limits, as will be seen from
the following table, but as a rule a given volume of a liquid will not
dissolve more than two or three times its own volume of a gas.
_Solubility of Gases in Water_
1 l. of water at 760 mm. pressure and at 0° will dissolve:
Ammonia 1148.00 l.
Hydrochloric acid 503.00
Sulphur dioxide 79.79
Carbon dioxide 1.80
Oxygen 41.14 cc.
Hydrogen 21.15
Nitrogen 20.03
In the case of very soluble gases, such as the first three in the table,
it is probable that chemical combination between the liquid and the gas
takes place.
2. _Nature of the liquid._ The character of the liquid has much
influence upon the solubility of a gas. Water, alcohol, and ether have
each its own peculiar solvent power. From the solubility of a gas in
water, no prediction can be made as to its solubility in other liquids.
3. _Influence of pressure._ It has been found that the weight of gas
which dissolves in a given case is proportional to the pressure exerted
upon the gas. If the pressure is doubled, the weight of gas going into
solution is doubled; if the pressure is diminished to one half of its
original value, half of the dissolved gas will escape. Under high
pressure, large quantities of gas can be dissolved in a liquid, and when
the pressure is removed the gas escapes, causing the liquid to foam or
_effervesce_.
4. _Influence of temperature._ In general, the lower the temperature of
the liquid, the larger the quantity of gas which it can dissolve. 1000
volumes of water at 0° will dissolve 41.14 volumes of oxygen; at 50°,
18.37 volumes; at 100° none at all. While most gases can be expelled
from a liquid by boiling the solution, some cannot. For example, it is
not possible to expel hydrochloric acid gas completely from its solution
by boiling.
SOLUTION OF SOLIDS IN LIQUIDS
This is the most familiar class of solutions, since in the laboratory
substances are much more frequently used in the form of solutions than
in the solid state.
~Circumstances affecting the solubility of a solid.~ The solubility of a
solid in a liquid depends upon several factors.
1. _Nature of the solid._ Other conditions being the same, solids vary
greatly in their solubility in liquids. This is illustrated in the
following table:
_Table of Solubility of Solids at 18°_
100 cc. of water will dissolve:
Calcium chloride 71.0 g.
Sodium chloride 35.9
Potassium nitrate 29.1
Copper sulphate 21.4
Calcium sulphate 0.207
No solids are absolutely insoluble, but the amount dissolved may be so
small as to be of no significance for most purposes. Thus barium
sulphate, one of the most insoluble of common substances, dissolves in
water to the extent of 1 part in 400,000.
2. _Nature of the solvent._ Liquids vary much in their power to dissolve
solids. Some are said to be good solvents, since they dissolve a great
variety of substances and considerable quantities of them. Others have
small solvent power, dissolving few substances, and those to a slight
extent only. Broadly speaking, water is the most general solvent, and
alcohol is perhaps second in solvent power.
3. _Temperature._ The weight of a solid which a given liquid can
dissolve varies with the temperature. Usually it increases rapidly as
the temperature rises, so that the boiling liquid dissolves several
times the weight which the cold liquid will dissolve. In some instances,
as in the case of common salt dissolved in water, the temperature has
little influence upon the solubility, and a few solids are more soluble
in cold water than in hot. The following examples will serve as
illustrations:
_Table of Solubility at 0° and at 100°_
100 cc. of water will dissolve:
At 0° At 100°
Calcium chloride 49.6 g. 155.0 g.
Sodium chloride 35.7 39.8
Potassium nitrate 13.3 247.0
Copper sulphate 15.5 73.5
Calcium sulphate 0.205 0.217
Calcium hydroxide 0.173 0.079
~Saturated solutions.~ A liquid will not dissolve an unlimited quantity of
a solid. On adding the solid to the liquid in small portions at a time,
it will be found that a point is reached at which the liquid will not
dissolve more of the solid at that temperature. The solid and the
solution remain in contact with each other unchanged. This condition may
be described by saying that they are in equilibrium with each other. A
solution is said to be _saturated_ when it remains unchanged in
concentration in contact with some of the solid. The weight of the solid
which will completely saturate a definite volume of a liquid at a given
temperature is called the _solubility_ of the substance at that
temperature.
~Supersaturated solutions.~ When a solution, saturated at a given
temperature, is allowed to cool it sometimes happens that no solid
crystallizes out. This is very likely to occur when the vessel used is
perfectly smooth and the solution is not disturbed in any way. Such a
solution is said to be _supersaturated_. That this condition is unstable
can be shown by adding a crystal of the solid to the solution. All of
the solid in excess of the quantity required to saturate the solution at
this temperature will at once crystallize out, leaving the solution
saturated. Supersaturation may also be overcome in many cases by
vigorously shaking or stirring the solution.
~General physical properties of solutions.~ A few general statements may
be made in reference to the physical properties of solutions.
1. _Distribution of the solid in the liquid._ A solid, when dissolved,
tends to distribute itself uniformly through the liquid, so that every
part of the solution has the same concentration. The process goes on
very slowly unless hastened by stirring or shaking the solution. Thus,
if a few crystals of a highly colored substance such as copper sulphate
are placed in the bottom of a tall vessel full of water, it will take
weeks for the solution to become uniformly colored.
2. _Boiling points of solutions._ The boiling point of a liquid is
raised by the presence of a substance dissolved in it. In general the
extent to which the boiling point of a solvent is raised by a given
substance is proportional to the concentration of the solution, that
is, to the weight of the substance dissolved in a definite weight of the
solvent.
3. _Freezing points of solutions._ A solution freezes at a lower
temperature than the pure solvent. The lowering of the freezing point
obeys the same law which holds for the raising of the boiling point: the
extent of lowering is proportional to the weight of dissolved substance,
that is, to the concentration of the solution.
~Electrolysis of solutions.~ Pure water does not appreciably conduct the
electric current. If, however, certain substances such as common salt
are dissolved in the water, the resulting solutions are found to be
conductors of electricity. Such solutions are called _electrolytes_.
When the current passes through an electrolyte some chemical change
always takes place. This change is called _electrolysis_.
[Illustration: Fig. 31]
The general method used in the electrolysis of a solution is illustrated
in Fig. 31. The vessel D contains the electrolyte. Two plates or rods,
A and B, made of suitable material, are connected with the wires
from a battery (or dynamo) and dipped into the electrolyte, as shown in
the figure. These plates or rods are called _electrodes_. The electrode
connected with the zinc plate of the battery is the negative electrode
or _cathode_, while that connected with the carbon plate is the positive
electrode or _anode_.
~Theory of electrolytic dissociation.~ The facts which have just been
described in connection with solutions, together with many others, have
led chemists to adopt a theory of solutions called _the theory of
electrolytic dissociation_. The main assumptions in this theory are the
following.
1. _Formation of ions._ Many compounds when dissolved in water undergo
an important change. A portion of their molecules fall apart, or
_dissociate_, into two or more parts, called _ions_. Thus sodium nitrate
(NaNO_{3}) dissociates into the ions Na and NO_{3}; sodium chloride,
into the ions Na and Cl. These ions are free to move about in the
solution independently of each other like independent molecules, and for
this reason were given the name ion, which signifies a wanderer.
2. _The electrical charge of ions._ Each ion carries a heavy electrical
charge, and in this respect differs from an atom or molecule. It is
evident that the sodium in the form of an ion must differ in some
important way from ordinary sodium, for sodium ions, formed from sodium
nitrate, give no visible evidence of their presence in water, whereas
metallic sodium at once decomposes the water. The electrical charge,
therefore, greatly modifies the usual chemical properties of the
element.
3. _The positive charges equal the negative charges._ The ions formed by
the dissociation of any molecule are of two kinds. One kind is charged
with positive electricity and the other with negative electricity;
moreover the sum of all the positive charges is always equal to the sum
of all the negative charges. The solution as a whole is therefore
electrically neutral. If we represent dissociation by the usual chemical
equations, with the electrical charges indicated by + and - signs
following the symbols, the dissociation of sodium chloride molecules is
represented thus:
NaCl --> Na^{+}, Cl^{-}.
The positive charge on each sodium ion exactly equals the negative
charge on each chlorine ion. Sodium sulphate dissociates, as shown in
the equation
Na_{2}SO_{4} --> 2Na^{+}, SO_{4}^{--}.
Here the positive charge on the two sodium ions equals the double
negative charge on the SO_{4} ion.
4. _Not all compounds dissociate._ Only those compounds dissociate whose
solutions form electrolytes. Thus salt dissociates when dissolved in
water, the resulting solution being an electrolyte. Sugar, on the other
hand, does not dissociate and its solution is not a conductor of the
electric current.
5. _Extent of dissociation differs in different liquids._ While
compounds most readily undergo dissociation in water, yet dissociation
often occurs to a limited extent when solution takes place in liquids
other than water. In the discussion of solutions it will be understood
that the solvent is water unless otherwise noted.
~The theory of electrolytic dissociation and the properties of solutions.~
In order to be of value, this theory must give a reasonable explanation
of the properties of solutions. Let us now see if the theory is in
harmony with certain of these properties.
~The theory of electrolytic dissociation and the boiling and freezing
points of solutions.~ We have seen that the boiling point of a solution
of a substance is raised in proportion to the concentration of the
dissolved substance. This is but another way of saying that the change
in the boiling point of the solution is proportional to the number of
molecules of the dissolved substance present in the solution.
It has been found, however, that in the case of electrolytes the boiling
point is raised more than it should be to conform to this law. If the
solute dissociates into ions, the reason for this becomes clear. Each
ion has the same effect on the boiling point as a molecule, and since
their number is greater than the number of molecules from which they
were formed, the effect on the boiling point is abnormally great.
In a similar way, the theory furnishes an explanation of the abnormal
lowering of the freezing point of electrolytes.
~The theory of electrolytic dissociation and electrolysis.~ The changes
taking place during electrolysis harmonize very completely with the
theory of dissociation. This will become clear from a study of the
following examples.
[Illustration: Fig. 32]
1. _Electrolysis of sodium chloride._ Fig. 32 represents a vessel in
which the electrolyte is a solution of sodium chloride (NaCl). According
to the dissociation theory the molecules of sodium chloride dissociate
into the ions Na^{+} and Cl^{-}. The Na^{+} ions are attracted to the
cathode owing to its large negative charge. On coming into contact with
the cathode, the Na^{+} ions give up their positive charge and are then
ordinary sodium atoms. They immediately decompose the water according to
the equation
Na + H_{2}O = NaOH + H,
and hydrogen is evolved about the cathode.
The chlorine ions on being discharged at the anode in similar manner may
either be given off as chlorine gas, or may attack the water, as
represented in the equation
2Cl + H_{2}O = 2HCl + O.
2. _Electrolysis of water._ The reason for the addition of sulphuric
acid to water in the preparation of oxygen and hydrogen by electrolysis
will now be clear. Water itself is not an electrolyte to an appreciable
extent; that is, it does not form enough ions to carry a current.
Sulphuric acid dissolved in water is an electrolyte, and dissociates
into the ions 2 H^{+} and SO_{4}^{--}. In the process of electrolysis of
the solution, the hydrogen ions travel to the cathode, and on being
discharged escape as hydrogen gas. The SO_{4} ions, when discharged at
the anode, act upon water, setting free oxygen and once more forming
sulphuric acid:
SO_{4} + H_{2}O = H_{2}SO_{4} + O.
The sulphuric acid can again dissociate and the process repeat itself as
long as any water is left. Hence the hydrogen and oxygen set free in the
electrolysis of water really come directly from the acid but indirectly
from the water.
3. _Electrolysis of sodium sulphate._ In a similar way, sodium sulphate
(Na_{2}SO_{4}), when in solution, gives the ions 2 Na^{+} and
SO_{4}^{--}. On being discharged, the sodium atoms decompose water about
the cathode, as in the case of sodium chloride, while the SO_{4} ions
when discharged at the anode decompose the water, as represented in the
equation
SO_{4} + H_{2}O = H_{2}SO_{4} + O
[Illustration: Fig. 33]
That new substances are formed at the cathode and anode may be shown in
the following way. A U-tube, such as is represented in Fig. 33, is
partially filled with a solution of sodium sulphate, and the liquid in
one arm is colored with red litmus, that in the other with blue litmus.
An electrode placed in the red solution is made to serve as cathode,
while one in the blue solution is made the anode. On allowing the
current to pass, the blue solution turns red, while the red solution
turns blue. These are exactly the changes which would take place if
sodium hydroxide and sulphuric acid were to be set free at the
electrodes, as required by the theory.
~The properties of electrolytes depend upon the ions present.~ When a
substance capable of dissociating into ions is dissolved in water, the
properties of the solution will depend upon two factors: (1) the ions
formed from the substance; (2) the undissociated molecules. Since the
ions are usually more active chemically than the molecules, most of the
chemical properties of an electrolyte are due to the ions rather than to
the molecules.
The solutions of any two substances which give the same ion will have
certain properties in common. Thus all solutions containing the copper
ion (Cu^{++}) are blue, unless the color is modified by the presence of
ions or molecules having some other color.
EXERCISES
1. Distinguish clearly between the following terms: electrolysis,
electrolyte, electrolytic dissociation, ions, solute, solvent, solution,
saturated solution, and supersaturated solution.
2. Why does the water from some natural springs effervesce?
3. (a) Why does not the water of the ocean freeze? (b) Why will ice
and salt produce a lower temperature than ice alone?
4. Why does shaking or stirring make a solid dissolve more rapidly in a
liquid?
5. By experiment it was found that a certain volume of water was
saturated at 100° with 114 g. of potassium nitrate. On cooling to 0° a
portion of the substance crystallized. (a) How many grams of the
substance remained in solution? (b) What was the strength of the
solution at 18°? (c) How much water had been used in the experiment?
6. (a) 10 g. of common salt were dissolved in water and the solution
evaporated to dryness; what weight of solid was left? (b) 10 g. of
zinc were dissolved in hydrochloric acid and the solution evaporated to
dryness; what weight of solid was left?
7. Account for the fact that sugar sometimes deposits from molasses,
even when no evaporation has taken place.
8. (a) From the standpoint of the theory of electrolytic dissociation,
write the simple equation for a dilute solution of copper sulphate
(CuSO_{4}); this solution is blue. (b) In the same manner, write one
for sodium sulphate; this solution is colorless. (c) How would you
account for the color of the copper sulphate solution?
9. (a) As in the preceding exercise, write a simple equation for a
dilute solution of copper chloride (CuCl_{2}); this solution is blue.
(b) In the same manner, write one for sodium chloride; this solution
is colorless. To what is the blue color due?
10. What component is present in concentrated sulphuric acid that is
almost wanting in very dilute sulphuric acid?
11. Why will vegetables cook faster when boiled in strong salt water
than when boiled in pure water?
12. How do you explain the foaming of soda water?
CHAPTER X
ACIDS, BASES, AND SALTS; NEUTRALIZATION
~Acids, bases, and salts.~ The three classes of compounds known
respectively as acids, bases, and salts include the great majority of
the compounds with which we shall have to deal. It is important,
therefore, for us to consider each of these classes in a systematic way.
The individual members belonging to each class will be discussed in
detail in the appropriate places, but a few representatives of each
class will be described in this chapter with special reference to the
common properties in accordance with which they are classified.
~The familiar acids.~ _Hydrochloric acid_ is a gas composed of hydrogen and
chlorine, and has the formula HCl. The substance is very soluble in
water, and it is this solution which is usually called hydrochloric
acid. _Nitric acid_ is a liquid composed of hydrogen, nitrogen, and
oxygen, having the formula HNO_{3}. As sold commercially it is mixed
with about 32% of water. _Sulphuric acid_, whose composition is
represented by the formula H_{2}SO_{4}, is an oily liquid nearly twice
as heavy as water, and is commonly called _oil of vitriol_.
~Characteristics of acids.~ (1) All acids contain hydrogen. (2) When
dissolved in water the molecules of the acid dissociate into two kinds
of ions. One of these is always hydrogen and is the cation (+), while
the other consists of the remainder of the molecule and is the anion
(-). (3) The solution tastes sour. (4) It has the power to change the
color of certain substances called _indicators_. Thus blue litmus is
changed to red, and yellow methyl orange is changed to red. Since all
acids produce hydrogen cations, while the anions of each are different,
the properties which all acids have in common when in solution, such as
taste and action on indicators, must be attributed to the hydrogen ions.
DEFINITION: _An acid is a substance which produces hydrogen ions when
dissolved in water or other dissociating liquids._
~Undissociated acids.~ When acids are perfectly free from water, or are
dissolved in liquids like benzene which do not have the power of
dissociating them into ions, they should have no real acid properties.
This is found to be the case. Under these circumstances they do not
affect the color of indicators or have any of the properties
characteristic of acids.
The familiar bases. The bases most used in the laboratory are sodium
hydroxide (NaOH), potassium hydroxide (KOH), and calcium hydroxide
(Ca(OH)_{2}). These are white solids, soluble in water, the latter
sparingly so. Some bases are very difficultly soluble in water. The very
soluble ones with most pronounced basic properties are sometimes called
the _alkalis_.
~Characteristics of bases.~ (1) All bases contain hydrogen and oxygen. (2)
When dissolved in water the molecules of the base dissociate into two
kinds of ions. One of these is always composed of oxygen and hydrogen
and is the anion. It has the formula OH and is called the _hydroxyl
ion_. The remainder of the molecule, which usually consists of a single
atom, is the cation. (3) The solution of a base has a soapy feel and a
brackish taste. (4) It reverses the color change produced in indicators
by acids, turning red litmus blue, and red methyl orange yellow. Since
all bases produce hydroxyl anions, while the cations of each are
different, the properties which all bases have in common when in
solution must be due to the hydroxyl ions.
DEFINITION: _A base is a substance which produces hydroxyl ions when
dissolved in water or other dissociating liquids._
~Undissociated bases.~ Bases, in the absence of water or when dissolved in
liquids which do not dissociate them, should have none of the properties
characteristic of this class of substances. This has been found to be
the case. For example, they have no effect upon indicators under these
circumstances.
~Neutralization.~ When an acid and a base are brought together in solution
in proper proportion, the characteristic properties of each disappear.
The solution tastes neither sour nor brackish; it has no effect upon
indicators. There can therefore be neither hydrogen nor hydroxyl ions
present in the solution. A study of reactions of this kind has shown
that the hydrogen ions of the acid combine with the hydroxyl ions of the
base to form molecules of water, water being a substance which is not
appreciably dissociated into ions. This action of an acid on a base is
called _neutralization_. The following equations express the
neutralization of the three acids by three bases, water being formed in
each case.
Na^{+}, OH^{-} + H^{+}, Cl^{-} = Na^{+}, Cl^{-} + H_{2}O.
K^{+}, OH^{-} + H^{+}, NO_{3}^{-} = K^{+}, NO_{3}^{-} + H_{2}O.
Ca^{++}, (OH)_{2}^{--} + H_{2}^{++}, SO_{4}^{--}
= Ca^{++}, SO_{4}^{--} + 2H_{2}O.
DEFINITION: _Neutralization consists in the union of the hydrogen ion of
an acid with the hydroxyl ion of a base to form water._
~Salts.~ It will be noticed that in neutralization the anion of the acid
and the cation of the base are not changed. If, however, the water is
expelled by evaporation, these two ions slowly unite, and when the water
becomes saturated with the substance so produced, it separates in the
form of a solid called a _salt_.
DEFINITION: _A salt is a substance formed by the union of the anion of
an acid with the cation of a base._
~Characteristics of salts.~ (1) From the definition of a salt it will be
seen that there is no element or group of elements which characterize
salts. (2) Salts as a class have no peculiar taste. (3) In the absence
of all other substances they are without action on indicators. (4) When
dissolved in water they form two kinds of ions.
~Heat of neutralization.~ If neutralization is due to the union
of hydrogen ions with hydroxyl ions, and nothing more, it
follows that when a given weight of water is formed in
neutralization, the heat set free should always be the same, no
matter from what acid and base the two kinds of ions have been
supplied. Careful experiments have shown that this is the case,
provided no other reactions take place at the same time. When
18g. of water are formed in neutralization, 13,700 cal. of heat
are set free. This is represented in the equations
Na^{+}, OH^{-} + H^{+}, Cl^{-}
= Na^{+}, Cl^{-} + H_{2}O + 13,700 cal.
K^{+}, OH^{-} + H^{+}, NO_{3}^{-}
= K^{+}, NO_{3}^{-} + H_{2}O + 13,700 cal.
Ca^{++}, (OH)_{2}^{--} + H_{2}^{++}, SO_{4}^{--}
= Ca^{++}, SO_{4}^{--} + 2H_{2}O + 2 × 13,700 cal.
~Neutralization a quantitative act.~ Since neutralization is a definite
chemical act, each acid will require a perfectly definite weight of each
base for its neutralization. For example, a given weight of sulphuric
acid will always require a definite weight of sodium hydroxide, in
accordance with the equation
H_{2}, SO_{4} + 2Na, OH = Na_{2}, SO_{4} + 2H_{2}O.
~Determination of the ratio in neutralization.~ The quantities of
acid and base required in neutralization may be determined in
the following way. Dilute solutions of the two substances are
prepared, the sulphuric acid being placed in one of the
burettes (Fig. 34) and the sodium hydroxide in the other. The
levels of the two liquids are then brought to the zero marks of
the burettes by means of the stopcocks. A measured volume of
the acid is drawn off into a beaker, a few drops of litmus
solution added, and the sodium hydroxide is run in drop by drop
until the red litmus just turns blue. The volume of the sodium
hydroxide consumed is then noted. If the concentrations of the
two solutions are known, it is easy to calculate what weight of
sodium hydroxide is required to neutralize a given weight of
sulphuric acid. By evaporating the neutralized solution to
dryness, the weight of the sodium sulphate formed can be
determined directly. Experiment shows that the weights are
always in accordance with the equation in the preceding
paragraph.
[Illustration: Fig. 34]
~Extent of dissociation.~ The question will naturally arise, When an acid,
base, or salt dissolves in water, do all the molecules dissociate into
ions, or only a part of them? The experiments by which this question can
be answered cannot be described here. It has been found, however, that
only a fraction of the molecules dissociate. The percentage which will
dissociate in a given case depends upon several conditions, the chief of
which are: (1) The concentration of the solution. In concentrated
solutions only a very small percentage of dissociation occurs. As the
solution is diluted the percentage increases, and in very dilute
solutions it may be very large, though it is never complete in any
ordinary solution. (2) The nature of the dissolved compound. At equal
concentrations substances differ much among themselves in the percentage
of dissociation. The great majority of salts are about equally
dissociated. Acids and bases, on the contrary, show great differences.
Some are freely dissociated, while others are dissociated to but a
slight extent.
~Strength of acids and bases.~ Since acid and basic properties are due to
hydrogen and hydroxyl ions respectively, the acid or base which will
produce the greatest percentage of these ions at a given concentration
must be regarded as the strongest representative of its class. The acids
and bases described in the foregoing paragraphs are all quite strong. In
10% solutions they are dissociated to about 50%, and this is also
approximately the extent to which most salts are dissociated at this
same concentration.
~Partial neutralization.~ 1. _Basic salts._ The chemical action
between an acid and a base is not always as complete as has
been represented in the foregoing paragraphs. For example, if
the base magnesium hydroxide (Mg(OH)_{2}) and hydrochloric acid
(HCl) are brought together in the ratio of an equal number of
molecules of each, there will be only half enough hydrogen ions
for the hydroxyl ions present.
Mg, (OH)_{2} + H, Cl = Mg, OH, Cl + H_{2}O.
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