Introduction to Chemistry

580 Pages·1968·21.452 MB·English

by  Amos Turk (Auth.)

ACADEMIC PRESS INTERNATIONAL EDITION
I N T R O D U C T I ON
TO C H E M I S T RY
AMOS TURK • HERBERT MEISLICH
FRANK BRESCIA • JOHN ARENTS
Department of Chemistry, The City College
of the City University of New York
ACADEMIC PRESS NEW YORK and LONDON
ACADEMIC PRESS INTERNATIONAL EDITION
This edition not for sale in the United States of America
and Canada.
Copyright © 1968, by Academic Press Inc.
All rights reserved.
No part of this book may be reproduced in any form,
by photostat, microfilm, or any other means, without
written permission from the publishers.
ACADEMIC PRESS INC.
Ill Fifth Avenue, New York, New York 10003
United Kingdom Edition published by
ACADEMIC PRESS INC. (LONDON) LTD.
Berkeley Square House, London W1X6BA
Library of Congress Catalog Card Number: 68-14653
Printed in the United States of America
PREFACE
The pursuit of chemistry as a science and a technology is one of the
most productive of man's activities today. Concepts, techniques, tools,
and the instruction of the student are all undergoing very rapid
advances.
An introductory textbook of chemistry should catch the spirit of
vitality, of intellectual challenge and excitement, of professional tra­
dition and style, and of great utility for man's needs that pervades the
world of the chemist.
This text is designed to serve in the one-year course in general chem­
istry. It has been our objective to hold the interest of the student who
does not have vocational objectives in chemistry, and also to provide
the necessary preparation for students who do plan to continue their
studies in the advanced courses.
The approaches used in first-year chemistry texts have been classi­
fied* as the "high road" and the "low road" methods. The "high road"
begins with the study of atoms and proceeds from the properties of
these chemical microcosms to evolve subject matter like stoichiom-
etry, and molecular structure and transformations. The "low road"
proceeds from macroscopic beginnings—the densities of vapors, the
composition of materials, and the like—and thus lays the basis for our
atomistic concepts.
The approach of this text does not fit well into either category; it is,
perhaps, a "middle road." After an introduction to atomic structure,
the text proceeds to chemical periodicity, and to a classical considera­
tion of atomic weights. Thus, by the time that stoichiometry is intro­
duced, the student has been exposed to the study of the atom from
both macroscopic and microscopic viewpoints.
Having considered atoms and the arithmetic and energetics of their
combination into molecules (Chapters 1-6), the next phase (Chapters
7-9) considers the nature of the interactions among atoms—in other
words, the nature of chemical bonding. After this introduction to the
question of how atoms aggregate to form molecules, it seems reason­
able to proceed to the question of how molecules aggregate to form
materials. This phase is covered in Chapters 10-13, which deal with
the nature of intermolecular forces, and the study of the states of
matter. Chapters 14 to 19 are concerned with the statics and dynamics
* Leonard K. Nash, Stoichiometry. New York: Addison-Wesley Publishing Company, 1966.
ν
vi • PREFACE
of chemistry—equilibrium and kinetics. Included in these chapters
are ionic equilibrium, acids and bases, and a study of galvanic cells.
The final portions of the book, Chapters 20-26, apply the previously
developed principles to a descriptive study of chemistry—representa­
tive and transition elements, organic chemistry, metals, nuclear chem­
istry, polymers, and biochemistry. These descriptive sections do not
abandon the approach of the first two thirds of the text. There is no
question of whether "principles" or "facts" take precedence, for the
two do not separate easily. It is important to introduce the descriptive
matter of chemistry in such a way that the student learns the facts as
his understanding of principles is reinforced.
Many of the features of our more comprehensive text, Fundamentals
of Chemistry: A Modern Introduction, are retained here. There are
two sets of problems for almost every chapter. The first set can be
assigned in toto to cover the essential ideas of the chapter. The second
set, "additional problems," provides additional drill, extends the con­
tents of the chapter, or, in some cases, poses a greater challenge.
Answers are provided for about one-half of the problems requiring
computation.
One other matter. A chemistry textbook usually serves in part as a
reference for the experiments in the general chemistry laboratory. We
are very much concerned with safety, and have taken every reasonable
opportunity to instill the philosophy of safe practice into these pages,
even though this is not a laboratory book. The laboratory manual that
is an optional accompaniment of this text, Fundamentals of Chemistry:
Laboratory Studies, emphasizes safety along with critical selection of
experimental procedures and the use of techniques of quantitative
chemistry.
We express our gratitude to our wives, our students, and our col­
leagues for their many kinds of help. We also thank Mrs. Coleman
London for research assistance and Mrs. Evelyn Manacek who typed
most of the manuscript. We particularly wish to acknowledge our ap­
preciation to Professor Darrell Eyman of the University of Iowa, whose
contribution to the concept of the "middle road" did much to set the
approach of this book, and whose critical reading of the entire manu­
script provided helpful insights and suggestions.
Α. Τ., Η. M., F. B., J. A.
New York,
February, 1968
1 • INTRODUCTION
1.1 • GENERAL AND HISTORICAL
Chemistry deals with the properties and transformations of materials.
Materials are samples of matter; they include the secretion from a
pituitary gland, the varnish on a table, the odorant from a flower, cobra
venom, a flame, a helium atom, and a proton. Chemists observe how
materials change; they coordinate their observations into useful con­
cepts, and they predict conditions under which specific changes will
occur—often to produce materials never previously observed. Chem­
istry's origins are ancient. Metallurgy, leather tanning, fermentation,
and the manufacture of soap, glass, and pigments were all devised be­
fore man acquired the concepts that now underlie his understanding of
chemical change. Progress in chemical theory and technology during
the past two centuries, however, has been spectacular compared with
that in all previous history. The most valuable group of ideas held by
chemists has proved to be the atomic and molecular hypotheses.
What we now consider to be modern chemistry began with investi­
gations of gases under pressure and vacuum (a discipline then called
"pneumatics") in the seventeenth century by Evangelista Torricelli,
Blaise Pascal, Otto von Guericke, and Robert Boyle. These studies led
to improvements in laboratory techniques that helped another genera­
tion of scientists (notably Karl Wilhelm Scheele, Henry Cavendish,
Joseph Priestley, and Antoine Lavoisier), about a century later, to
formulate a quantitative basis for chemical changes, especially com­
bustion and other reactions involving oxygen. These advances, in turn,
set the stage for the chemical pioneers of the nineteenth century
(John Dalton, Amadeo Avogadro, Jons Jakob Berzelius, and Stanislao
ι
2 • INTRODUCTION
Cannizzaro) to interpret chemical changes in terms of atoms and mole
cules, and to devise rational systems of atomic and molecular weights.
The latter half of that century witnessed a very fruitful growth of sys
tematizing concepts—the periodic table, the structural theory of or
ganic chemistry, and stereochemistry (the geometry of molecules). In
1896 Henri Becquerel discovered radioactivity, thus initiating a new
chain of discoveries that led to a great refinement of our ideas about
the atom, and to new understanding of chemical processes. It is to
these ideas that we shall turn to provide the major conceptual frame
work for our study of chemistry.
Refer to Appendix! for a review of physical concepts, measure­
ment scales, and significant figures. This material is fundamental
to discussions throughout the book. Some of the problems at the
end of this chapter are based on material in the Appendix.
1.2 • DEFINITIONS OF SOME CHEMICAL TERMS
We shall define and discuss briefly some terms which, over a period
of many years, have become part of the language of chemists.
The properties of a material are its distinguishing characteristics.
Accidental or extensive properties depend on the amount of matter
present in a sample of the material; specific, or intensive, propertie,s on
the other hand, do not depend on the amount of matter in the sample.
Thus, the white color of a piece of chalk is a specific property; its
length is an accidental one.
A substance is any variety of matter of recognizably definite compo
sition and specific properties. The term is used in distinction to body,
or object, which refers to a particular item of matter. Thus, a chair
(object) is made of wood (substance). The composition of a substance
is its makeup of constituent substances, usually expressed in terms
of percent or fraction by weight.
Some substances have precisely fixed compositions which are asso
ciated with their properties; they are said to be pure substances. For
example, red iron rust can be obtained as a pure substance comprising
69.94% iron and 30.06% oxygen. Coal, on the other hand, is not a
pure substance; its carbon content ranges from 35 to 84%. Of course,
a pure substance may be contaminated by admixture of foreign matter.
The important point, however, is that the pure substance, when it is
recovered from such a mixture, retains its definite composition and
specific properties.
It is believed that the constant compositions associated with pure
substances are maintained by linkages among elementary units of
3-1.2 DEFINITIONS OF SOME CHEMICAL TERMS
matter; such linkages are called chemical bonds. A transformation
accompanied by the making or breaking of chemical bonds is called
a chemical change, or chemical reaction. Examples are combustion,
corrosion, photosynthesis, and digestion.
A physical change of a substance does not involve change of definite
composition or specific properties. Alterations in the dimensions of
objects, or in the states of aggregation of their constituents, are con­
sidered to be physical changes. Examples are fracture, deformation,
pulverizing, drawing (as of a metal wire), thermal expansion or con­
traction* melting, boiling, and freezing.
The types of behavior that a substance exhibits in chemical reactions
are called its chemical propertie;s other intensive characteristics of
a substance are called its physical propertie.s
Decomposition is a chemical reaction in which the constituent enti­
ties of a substance break down into simpler forms. Most of the many
substances known to man can undergo decompositions that involve
net energy changes up to about 2 χ 103 cal/g (released) or 3 χ 104
cal/g (absorbed) and which yield two or more decomposition products.
A relatively few substances (somewhat over one hundred) do not de­
compose at all within these ranges of energy change or, if they do,
give only one ultimate product (for example, ozone -> oxygen). Such
substances are considered to be the stuff of which all other substances
are made, and are called elements. The fundamental unit of the ele­
ment is the atom.
A nonelemental pure substance is called a compound substance or
a compound.
Electrically neutral individual particles of ordinary matter whose
atoms are linked together by chemical bonds are called molecules.
Individual unbonded atoms are also considered to be molecules,- for
example, the molecules in helium gas are individual atoms. There is
no absolute upper limit of size for molecules. Viruses, which reach
dimensions of hundreds of Angstrom units,* are sometimes called
"giant molecules." Molecules of substances that are gaseous in ordi­
nary terrestrial environments are usually less than 10 A; small mole­
cules like those of water and hydrogen chloride are around 2 to 4 A.
Electrically charged atoms or groups of atoms are called ions. Posi­
tive ions are cations, negative ones anions. The smallest ions are indi­
vidual charged atoms (for example, sodium ion, Na +, or fluoride ion,
F~). Ions may also be groups of relatively few atoms, such as sulfate
ion, S042_. Scientists concerned with air use the term "atmospheric
ions" to denote charged dust particles. Ions can be arranged in some
closely packed pattern, as in a crystal. In such cases, it is sometimes
conventional to designate as a molecule the smallest group of ions
* One Angstrom unit = 10~8 cm; see Appendix 1.
4 • INTRODUCTION
whose charges just neutralize each other—for example, a molecule of
calcium chloride consisting of one Ca2+ and two Ch ions. We shall,
however, use the word "molecule" to denote only individual entities
as described above.
1.3 • THE PURITY OF COMPOUNDS
A pure substance can be defined as one with a precisely fixed compo
sition that is associated with its properties. We usually think of it as
being entirely composed of like molecules. Blit the situation is really
not so simple. When a given constituent has been separated from a
mixture, and no further efforts succeed in separating that constituent
into additional components, we must assume that it is pure. As far as
laboratory operations are concerned, there is no way of showing that
the assumption is false except by developing a more effective method
to separate the "pure" substance into new components. In laboratory
practice, then, the criterion of purity of a substance is simply the inabil
ity of the experimenter to isolate or otherwise detect foreign material.
Purification involves physical separation, which occurs as a conse
quence of the differences in properties among the components to be
separated. For example, bricks and grains of sand differ from each
other in size; a sieve will retain the bricks, and allow the sand to fall
through under the influence of gravity. Materials may differ from each
other in many attributes other than particle size—electrical, magnetic,
and solubility properties are among those which frequently serve as
bases for purification procedures. The important thing to remember is
that the chemical investigator does not take the viewpoint that puri
fication ceases when the compound is pure, rather that the compound
must be called pure when purification ceases.
1.4 • CHEMICAL SYMBOLS, FORMULAS, AND EQUATIONS
Atoms or elements are denoted by symbols of one or two letters, like
H, U, W, Ba, At, and Zn. (See inside back cover for names.)
Compounds or molecules are represented by formulas that consist of
symbols and subscripts, sometimes with parentheses. The subscript
denotes the number of atoms represented by the symbols to which it
is attached. Thus (COOH) is a formula that represents a molecule of
2
oxalic acid or the substance oxalic acid. The molecule consists of 2
atoms each of carbon and hydrogen, and 4 atoms of oxygen. The sub
stance consists of matter that is an aggregate of such molecules. The
formula for phosphorus vapor is P ; this tells us that the molecules
4
consist of 4 atoms each. Sometimes formulas are written so as to give
other information in addition to the numbers or ratios of atoms; such
representations will be taken up in later chapters. In the nineteenth
century it was customary to use dots to separate portions of a formula
5-1.4 CHEMICAL SYMBOLS, FORMULAS, AND EQUATIONS
that correspond to formulas of simpler substances. For example, cal
cium sulfate, CaS04, was written as CaOS03. Today such usage per
sists in some formulas of substances, called hydrates, that produce
water when they decompose, like CuS04-5H20. We prefer to write
CuS04(H20)5. 4
Chemical transformations are represented by chemical equations,
which tell us what molecules or substances react and what ones are
produced, and in what molecular ratios. The equation for the burning
of methane in oxygen to produce carbon dioxide and water is
CH4 + 202 > C02 + 2H20
Each coefficient applies to the entire formula that follows it. Thus 2H 20
means 2(H20). This gives the following molecular ratios: reacting mate
rials, 2 molecules of oxygen to 1 of methane; products, 2 molecules of
water to 1 of carbon dioxide. The equation written above is balanced
because the same number and kinds of atoms, one of carbon and four
each of hydrogen and oxygen, appear on each side of the arrow. The
balancing of chemical equations is taken up in more detail in Chapters
5 and 15.
PROBLEMS
1. Chemical and physical change Identify the chemical and the physical
changes in the following sequences.
(a) Carrots for an ulcer patient are scraped, cooked, mashed, sieved, swal
lowed, digested, absorbed, and utilized metabolically for coenzyme
manufacture.
(b) Oxygen is liquefied, poured into rocket storage tanks, used as a fuel
component during flight, used to pressurize the capsule, inhaled by the
astronauts, and converted during respiration to C02 and H20.
(c) A lump of sugar is ground to a powder and then heated in air. It melts,
then darkens, and finally bursts into flame and burns completely.
2. Definitions of terms Discuss the inadequacies and limitations of the follow
ing definitions. Supply better ones, (a) A molecule is the smallest particle of
matter that retains the specific properties of the matter, (b) An element is a
substance that cannot be decomposed by chemical means, (c) An ion is any
electrically charged body, (d) A compound is a substance that contains more
than one element.
3. Purification Black gunpowder consists of sulfur, charcoal, and potassium
nitrate. Sulfur dissolves in the volatile solvent carbon bisulfide; potassium ni
trate dissolves in water; charcoal is insoluble in both. Outline a procedure for
separation and isolation of the components, and suggest methods to identify
them.
4. Purification For each of the following purification methods, point out what
physical separation is occurring, and what differences in the properties of the
substances involved make the separation possible.
6 • INTRODUCTION
(a) Scrap iron is recovered from junk with the aid of a magnet.
(b) Moving ionized material is deflected in a magnetic field; the lighter par­
ticles deflect to the greater extent. The separation is recorded on a photo­
graphic plate.
(c) A perfume chemist places a drop of an essential oil on a strip of blotting
paper, which he leaves open to evaporation. He sniffs the paper each
hour, until a constant, characteristic odor is noted.
5. Purity of matter In 1868 a sample of water, after treatment by distillation
and other methods, resisted further attempts at purification and was called
"pure." Techniques available in 1968 made it possible to separate the "pure"
sample into several components, including protium deuterium oxide and dideu-
terium oxide. Has the purity of the unchanged sample deteriorated because of
advance in technique? Was the sample pure in 1868? Is it possible for future
advances in methods to show that components separated today are "impure"?
6. Chemical symbols, formulas For the following formulas, give the number of
atoms of each element found in a molecule of the compound; give the total
number of atoms in one molecule of the compound: (NH 4)2HP04; AI(OH)-
(CH:iC02)2l K3Fe(C2OJ3(H20)3.
The following problems are related to material in Appendix 1.
7. Matter and energy Explain the fallacy in each of the following statements.
(a) Smoke pollution from power plants could be profitably controlled by
collecting all the stack emissions and chemically reconverting them
to fuel.
(b) A window air conditioner will also provide cooling if it is placed inside
the room with the windows and doors closed.
(c) Energy used to pulverize coal is lost when the coal burns.
8. Significant figures How many significant figures are there in each of the
following quantities? (a) 6.0342 g; (b) 2.00 χ 10"3 ml; (c) 1000 dollars; (d)
0.002°C; (e) 62.9%.
9. Units; significant figures It has been said that a number as large as 10100 is
never needed to express any magnitude that has physical significance. What is
the ratio between the apparent mass of a neutron star, estimated density 6 χ
1013 g/cm3 and diameter 1.6 χ 104 m, and that of the electron, approximate
rest mass 9.11 χ 10~28 g?
10. Density Copper pellets are poured into a measuring cylinder up to the
50-ml mark. The cylinder is then filled with water to the same mark. Data are
Weight of empty cylinder 104 g
Weight of cylinder + copper 371 g
Weight of cylinder + copper + water 391 g
Temperature of system 20°C
Density of water at 20°C 0.998 g/ml
Calculate the density of copper at 20°C.
11. Pressure The density of mercury at 0°C is 13.595 g/cm3. Prove that 1
torr = pressure exerted by a mass of 1.3595 g on an area of 1 cm2.
ANSWER 10. 8.9 g/ml.