Steven S. Zumdahl
`
`UNIVERSITY OF ILLINOIS
`
`Chemistry
`
`LEXINGTON, MASSACHUSETTS TORONTO
`
`Lupin Ex. 1035 (Page 1 of 190)
`
`

`

`To my parents and to Eunice, Whitney, and Leslie.
`
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`
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`
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`
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`
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`
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`
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`
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`
`Copyright © 1986 by D. C, Hcath and Company.
`
`All rights reserved. No part of this publication may be
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`
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`
`Printed in the United States of America.
`
`International Standard Book Number: 0-669-04529-2
`
`Library of Congress Catalog Card Number: 85-60981
`
`Lupin Ex. 1035 (Page 2 of 190)
`
`

`

`Molecules,
`Ions
`
`2.1 The Early HistoF/ of Chemistry
`2,9 Fundamental Chemical Laws
`2.3 Dalton’s Atomic Theo~/
`2.4 Early Experiments to
`Characterize the Atom
`The Electron
`Rad ioactivity
`The Nuclear Atom
`2,5 The Modern View of Atomic
`Str~cture~An Introduction
`2.6 Molecules and Ions
`2.? An Introduction to the Periodic
`Table
`2.8 Naming Compounds
`
`T o fully appreciate a science like chemistry, we must understand
`
`something of its origins. In this chapter we will see how chemistry
`developed, paying particular attention to the crucial experiments
`and fundamental laws that form the basis of modern chemistry. We
`will also consider chemical nomenclature and introduce some of the most
`important chemical concepts.
`
`The Early History of Chemistry
`
`PURPOSE
`
`To give a brief account of eafiy chemical discoveries.
`
`Chemistry has been important since ancient times. The processing o1 natural
`ores to produce metals for ornaments and weapons and the use of embNming fluids
`are two examples of chelrdcal phenomena that were utilized prior to 1000
`~l’hc Greeks were the first to try to explain why chemical changes occur. By
`about 400 me. they had proposed that all matter was composed of four fundamental
`substances: fire, earth, water, and air. The Greeks also considered the question of
`whether matter is continuous, and thus infinitely divisible into smaller pieces, or
`composed of small indivisible particles One st pporter f the latter p s tion was
`Democritus who used the term atomos (wh ch ater became atoms) to describe
`¯
`these ultlnmte particles. However, because the Greeks had no experiments to test
`their ideas, no definitive conclusion about the divisibility of matter was reached.
`
`31
`
`Lupin Ex. 1035 (Page 3 of 190)
`
`

`

`TheAIchemists, an oil on slate
`painting by Jan van der Stra~ in
`1570.
`
`The next 2000 years of chemical history were dominated by a pseudoscience
`called ~lehemy. Alchemists were often mystics and fakes who were obsessed with
`the idea of turning cheap metals into gold. However, this period also saw impontant
`discoveries: elements such as mercury, sulfur, and antimony were discovered, and
`alchemists learned how to prepare the mineral acids
`The foundations for modern chemistry were laid in the sixteenth century with
`the development of systematic metallurgy (extraction of metals from ores) by a
`German, Georg Bauer, and the medicinal application of minerals by a Swiss alche-
`mist called Paracelsus.
`The first person to perform truly quantitative physical expe(maents was Robert
`Boyle (1627-1691), who carefully measured the relationship between the pressure
`and volume of gases. When Boyle published his book, The Sceptical Chemist, in
`1661, the quantitative sciences of physics and chemistry were born. In addition to
`his results on the quantitative behavior of gases, Boyle’s other major contribution to
`chemistry consisted of his ideas about the chemical elements. Boyle held no precon-
`ceived notion about the number of elements. In his view a substance was an element
`unless it could be broken down into two or more simpler substances. As Boyle’s
`experimental definition of an element became generally accepted, the list of known
`elements began to grow, and the Greek system of four elements finally died. Al-
`though Boyle was an excellent scientist, he was not always right. For example, he
`clung to the alchemist’s views that metals were not true elements and that a way
`would eventually be found to change one metal to another.
`The phenomenon of combustion evoked intense interest in the seventeenth and
`eighteenth centuries: The German chemist Georg Stahl (1660-1734) suggested that
`a substance he called phlogiston flowed out of the burning material. Stahl postulated
`that a substance burning in a closed container eventually stopped burning, because
`the air in the container became saturated with phlogiston. Oxygen gas, discovered
`by Joseph Priestley (1733-1804), ,an English clergyman and scientist (Fig. 2. I),
`was found to support vigorous combustion and was thus supposed to be low in
`phlogiston. In fact, oxygen was originally called "dephlogisticated air."
`
`Figure 2.1
`
`Joseph Priestley was born in England on March
`13, 1733, and showed a great talent for
`science and languages from an eady age.
`Priestley performed many important scientific
`experiments, among them the discovery that
`the gas produced by the fermentation of grain
`0ater identified as carbon dioxide) could be
`dissolved in water to produce the pleasant
`drink called seltzer. AIs% as a result of meeting
`Benjamin Franklin in London in 1766, Priestley
`became interested in electricity and was the
`first to observe that graphite was an electrical
`conductor. However, Priestley’s greatest
`discovery occurred in 1774 when he isolated
`oxygen by heating mercuric oxide.
`Because of his nonconformist political
`views (he supported both the American and
`French revolutions), he was forced to leave
`England (a mob burned his house in
`Birmingham in 1791). He spent his last decade
`peacefully in the United States, and died in 1804
`
`32
`
`[] Chapter Two Atoms, Molecules, and Ions
`
`Lupin Ex. 1035 (Page 4 of 190)
`
`

`

`38cience
`sed with
`nportant
`red, and
`
`xry with
`~s) by a
`;s alchc-
`
`~ Robert
`
`~mist, in
`J.ition to
`ration to
`precon-
`element
`Boyle’s
`f known
`ied. A1-
`nple, he
`tt a way
`
`anth and
`sled that
`~stulated
`becausc
`;covered
`ig. 2.1),
`~ low in
`
`°2 Fundamental Chemical Laws
`PURPOSE
`
`To describe and illustrate the laws of conservation of mass, definite proportion,
`and multiple proportions.
`
`By the late eighteenth centm2¢, combustion had bccn studied extensivcly; the
`gascs carbon dioxide, nitrogen, hydrogen, and oxygen had been discovered; and the
`list of elements continued to grow. However, it was Antoine Lavoisier (1743-
`1794), a French chemist (Fig. 2.2), who finally explained the true nature of
`bustion, thus clearing the way for the tremendous progress that was made near the
`end of the eighteenth century. Lavoisier, like Boyle, rcgarded measurcment as the
`essential opcration of chemistry. His experiments, in which he carefully weighed
`the reactants and the products of various reactions, showed that mass was’ neither
`creawd nor destroyed. Lavoisier’s discovery of this law of conservation of nlass
`was the basis for the developmcnts in chemistry in the nineteenth century.
`Lavoisier’s quantitativc experiments showed that combustion involved oxygen
`(which Lavoisier named), not phlogiston. He also discovered that life was sup-
`ported by a process that also involved oxygen and was similar in many ways to
`combustion. In 1789 Law)isier published the first modern chemistry textbook, Ele-
`mentary Treatise on Chemistry, h~ which he presented a unified picture of the
`chemical knowlcdge assembled up to that timc. Unfortunately, in the same year the
`text was published, the French Revolution broke out. Lavoisier, who had been
`associated with collecting taxes for the government, was executed on tbe guillotine
`as an enemy of the people in 1794.
`After 1800 chemistry was dominated by scientists who, following Lavoisier’s
`lead, perlbrmed carcl)l weighing experiments to study the course of chemical reac-
`tions and to determh~e thc composition of various chentical compounds. One of
`these chemists, a Frenchman, Joscph Proust (1754 1826), showed that a given
`
`Figure 2.2
`
`Antoine Lavoisier was born in Paris on
`August 26, 1743. Although Lavoisier’s
`fdther wanted his son to follow him
`
`Lavoisier was fascinated by science.
`From the beginning of his scientific
`career, Lavoisier recognized the
`importance of accurate
`measurements. His careful weighings
`showed that mass was conserved in
`chemical reactions and that
`combustion involves reaction with
`oxygen. Also, he wrote the first
`modern chemistry textbook. It is not
`surprising that Lavoisier is often called
`fJ~e father of modern chemistry.
`To heip suppor~ his scientific
`work, Lavoisier invested in a private
`tax-collecting firm and married the
`daughter of one of the company
`executives. His connection to the tax
`coiled.ors proved fdtal, for radical
`French revolutionaries demanded his
`~xecution, which occurred on the
`guillotine on May 8, 1794.
`
`2.2 Fundamental Chemical Laws [] 33
`
`Lupin Ex. 1035 (Page 5 of 190)
`
`

`

`Figure 2.:3
`
`John Dalton (1766-1844), an
`Englishman, began teaching at a
`Quaker school wl,en he was 12. His
`tascination with science included an
`intense interest in meteorology (he
`kept careful daily weather records for
`46 years), which led to an interest in
`the gases of the air and their ultimate
`components, atoms. Dalton is best
`known tot his atomic theoly, in which
`he postulated that the fundamental
`differences among atoms are their
`masses He was the first to prepare a
`table of relative atomic weights.
`Dalton was a humble man with
`several apparent handicaps: he was
`poor; he was not articulate; he was
`not a skilled experimentalist; and he
`was color blind, a terrible probiem
`for a chemist in spite of these
`disadvantages, he helped to
`reVolutionize the science of
`chemistry.
`
`compound always contains exactly the same proportion q[ elements by weight. For
`example, Proust found that the substance copper carbonate is always 5.3 parts
`copper to 4 parts ox’vgen to 1 part carbon (by mass). The principle of the constant
`composition of compounds, orlgmall~ called Pr ust law, is now known as the law
`
`of definite proportion.
`Proust’s discovery stimulated .lohn Dalton (1766 1844), an English school-
`teacher (Fig. 2.3), to think about atonls. Dalton reasoned that if e]cnlents were
`composed of tiny individual particles, a given compound should alv~a.’rs contain the
`same combination of these atoms- This concept explained wh.’, the same relative
`
`masses of elements were always found in a given compound.
`But Dalton discovered ar~other prhaciple that convinced him even more of the
`existence of atoms. He noted, for example, that carbon and oxygen formed
`different compounds that contained different relative amounts of carbon and oxygen
`as shown by the following data:
`
`WeigN of oxygen that combines
`with 1 g of carbon
`
`Compound I
`Compound II
`
`1.33 g
`2.66 g
`
`Dalton noted that compound II contained twice as ranch oxygen per gram of carbon
`as conlpound I, a fact that could bc easily explained in tetras of atoms. Compognd I
`might be CO. and compound lI might be CO._. This principle, which was found to
`apply to compounds of other elements as well, beoame known as the law of inulti-
`pie proportions: when two elements.form a series (g" compoands, the ratios of the
`masses oj the second element that combine with 1 gram q~ the first element can
`always be reduced to ~mall whole nambers.
`}o make sure the significance of this observation is clear, in Sample Exercise
`2.1 we will consider data for a series of compounds consisting of nitrogen and
`
`oxygen.
`
`Sarnpl~ F~;ercis¢ ~. 1
`
`Tile following data were collected for several compounds of nitrogen and ox~geu:
`
`glass of n trogen thai combines
`with 1 g of oxygen
`
`Compound I
`Compound II
`Compound III
`
`1.750 g
`0.8750 g
`0.4375 g
`
`Show how these data illustrate tile law of multiple proportions.
`
`Solution
`For the law of multiple proportions to hold, the ratios of the masses of nitrogen
`combining v4th l gram of oxygen in each pair of compounds shonld be small whole
`numbers. We therefore compute the ratios as follows:
`
`34 % Chapter Two Atoms, Molecules, and Ions
`
`Lupin Ex. 1035 (Page 6 of 190)
`
`

`

`r. For
`parts
`lstant
`c law
`
`hool-
`were
`in the
`lative
`
`af the
`two
`~ygen
`
`:m’bon
`aund 1
`and to
`multi-
`of the
`
`aercise
`~n and
`
`xygen:
`
`I
`II
`
`II
`
`1II
`
`I
`
`III
`
`1.750
`0.875
`
`2
`1
`
`0.875
`
`0.4375
`
`2
`
`1.750 4
`
`0.4375 1
`
`These results support the law of multiple proportions.
`
`The significance of the data in Sample Exercise 2.1 is that compound I contains
`twice as much nitrogen (N) per grain of oxygen (O) as does compound II and that
`compound II contains twice as much nitrogen per grain of oxygen as does com-
`pound III. In terms of the numbers of atoms combining, these data can bc explained
`by any of the following formulas*:
`
`Compound 1
`Compound II
`Compound III
`
`N20
`NO
`NO2
`
`NO
`N402
`or NO_, or N202
`NO4
`N:O,I
`
`In tb.ct an infinite number of other possibilities exists. Dahon could not deduce
`absolute formulas from the available data on relative masses. However, the data on
`the composition of compounds in terms of the relative masses of the elements
`supported his hypothcsis that each element consisted of a certain type of atom and
`that compounds were fornmd from specific combinations of atoins.
`
`Dalton’s Atomic Theory
`
`PURPOSE
`
`td To dcscrlbe Dalton’s theo~ of atoms and show the significauce of Gay-Lus-
`sac’s experiments.
`
`In 1808 Dalton published A New System qfChemical Philosophy, in which he
`presented his theory of atoms:
`
`1, Each element is’ made up of tiny particle,v called atoms.
`2. The atoms of a given element are identical; the atoms of d!~f~rent elements are
`diff?’rent in some fi~tMamental way or ways.
`
`3. Chemical compounds are.~brmed when atoms combine with each other. A given
`compound always has the same relative numbers and types of atoms.
`4. Chemical reactions involve reorganization of the atoms changes in the way
`they are bound together. The atoms themselves are not changed in a chemical
`reaction.
`
`im)gen
`whole
`
`"Subscripts ale used to show the numbers of atoms present. The number i is understood and thus is not
`writtem The symbols for the eletnents axed the ,~,riting of chemical formulas will be illusn’ated turther in
`Sections 2.6 and 2.7.
`
`2,3 Dalton’s Atomic Theory I ] 35
`
`Lupin Ex. 1035 (Page 7 of 190)
`
`

`

`It is instructive to consider Dalton’s reasoning on the relative masses of the
`atoms of the varimls elements. In Dalton’s time water was known to bc composed of
`the elements hydrogen and oxygen, with 8 grams of oxygen present for every I
`gram of hydrogen. If the formula for water were OH, an oxygen atom would have to
`have 8 times the mass of a hydrogen atom. However, if the fom~ula for water were
`H,=O (two atoms of hydrogen for every oxygen atom), this would mean that each
`atom of oxygen is 16 times as heavy as each atom of hydrogen (since the ratio of the
`mass of one oxygen to that of ~wo hydrogens is 8 to 1). Because the formula for
`water was not then known, Dalton could not specify the relative masses of oxygen
`and hydrogen unambiguously, qo solve the problem, Dalton made a fundamental
`assumption: he decided that natm’e would be as simple as possible. This assumption
`led him to conclude that the formula for water should be OH. He thus assigned
`hydrogen a mass of I and oxygen a mass of 8.
`Using similar reasoning for other compounds, Dalton prepared the tirst table of
`atomic masses (called atomic weights by chenfists, since mass is often determined
`by comparison to a standard mass-- a process called weighing). Many of the mas~es
`were later proved to bc wrong because of Dalton’s incorrect assumptions about the
`formulas of certain compounds, but the constraction of a table of masses was an
`important step forward.
`Although not recognized as such for many years, the keys m determining abso-
`lute formulas for compounds were provided in the experimental work of the French
`chemist Joseph Gay-Lussac (1778 1850) and by ~he hypothesis of an Italian chem-
`ist named Amadeo Avogadm (1776-1856). In 1809 Gay-Lussac parti)rmed experi-
`ments in which he measured (under the same conditions of temperatm’e and pres-
`sure) the volumes of gases tha~ reacted with each other. For example, Gay-Lussac
`found that 2 volumes of hydrogen react with 1 volume of oxygen to form 2 volumes
`of gaseous water and that 1 volume of hydrogen reacts with I volume of chlorine to
`form 2 volumes ot hydrogen chloride. These results ,are represented schematically
`in Fig. 2.4
`In 1811 Avogadro interpreted these results by proposing that, at the same
`temperature and pressure, equal volumes oJ diJ]?~rent gases contain the same num-
`ber of particles. This assumption (c’,dled Avogadro’s hypothesis) makes sense if
`the distances between the particles in a gas are very great cmnpzu-ed to the sizes of
`the particles. Under these conditions the volume of a gas is determined by the
`number of molecules present, not by the size of the individual particles.
`
`Figure 2.4
`
`A representation of some of Gay-
`Lussac’s experimental results on
`combining gas volumes.
`
`30
`
`[] Chapter Two Atoms, Molecules, and Ions
`
`2 volumes hydrogen
`chloride
`
`Lupin Ex. 1035 (Page 8 of 190)
`
`

`

`)f the
`;cd (ff
`ery 1
`
`each
`:ffthc
`[a for
`:ygen
`~ental
`tption
`igned
`
`~le of
`nined
`:asses
`at the
`
`as an
`
`abso-
`rench
`hem-
`tperi-
`pres-
`ussac
`
`ine to
`ically
`
`ase if
`:es of
`y the
`
`If Aw)gadro’s hypothesis is correct, Gay-Lussac’s result,
`
`2 volumes of hydrogen react with 1 w)lume of oxygen
`~ 2 volumes of water vapor
`
`can be expressed as fbllows:
`
`2 molecules* of hydrogen react with 1 molecule of oxygen
`--~ 2 molecules of water
`
`These observations can best be explained by assunfing that gaseous hydrogen, oxy-
`gen, and chlorine are all composed oF diatomic (two-atom) molecules: Hz, 02, and
`C12, respectively. Gay-Lussac’s results can then be represented as shown in Fig.
`2.5. (Note that this reasoning suggests that the formula for water is H:O, not OH as
`Dalton believed.)
`
`Unfortunately, Avogadro’s interpretations were not accepted by most chemists
`and a half-century of ¢ordusion followed, in which many different assumptions
`were made about formulas and atomic masses.
`During the nineteenth century painstaking measurements were made of the
`masses of various elements that combined to form compounds. From these experi-
`ments a list of relative atomic masses could be detemfined. One of the chemists
`in’~olved in adding to this list was a Swede named J6ns Jakob Be~elins (1779-
`1848), who discovered the elements cerium, selenium, silicon, and thorium and
`developed the modem symbols for the elements used in writing the formulas of
`compounds.
`
`Early Experiments to
`Characterize the Atom
`
`PURPOSE
`To summarize the experiments that characterized the structure of the atom.
`
`Dalton’s atomic theory caused chemistry to become more systematic and more
`sensible. The concept of aioms was clearly a good idea, and scientists became very
`interested in the structure of the atom
`
`~A molecule is a collection of atoms (see Section 2.6 .
`
`Figure 2,5
`
`at the molecular level. The circles
`represent atoms in the molecules.
`
`2.4 Early Experiments to Characterize the Atom [] 37
`
`Lupin Ex. 1035 (Page 9 of 190)
`
`

`

`The Electron
`The first impmlant experiments thal led to an m~derstanding of the composition of
`the atom were done by the English physicist .l. J, Thomson (Fig. 2,6), who studied
`electrical discharges in partially evacuated tubes called cathode-ray tabes (l~ig. 2.7)
`durkng the period from 1898 to 1903. Thomson found that, when high voltage was
`applied to the tube, a "ra~’" he called a cathode ray (because it emanated from the
`negative clectrodc, or cafl~ode) was produced. Because tl~s ray was produced at the
`negative electrode and was repelled by the negative pole of an applied electric field
`(see Fig. 2.8), Thomson postulated that the ray ~as a stream of negatively charged
`particles, which he called electrons. By experiments in which he measured the
`deflection of the beain of elcctrons in a magnetic field, Thomson dctermined the
`
`charge4o-trglss ratio of an electron:
`
`~" = -1.76 x 10s C/g
`m
`
`where e represents the charge on the electron in coulombs and m represents the
`
`electrm~ mass in ~raiBs.
`One of Thomson’s prhnary goals in his cathode-ray tubc expcriments was to
`gain an understanding of the strncture of the atom. Hc reasoned that, since electrons
`~ould be produced frmn electrodes made of ~ arious types of mct~ds, all atoms must
`contain electrons. Since atoms were knowu to be electrically neutral, Thomson
`
`further assumed that atoms also must contain some positive charge. Thomson pos!u-
`lated* that ,’m atom consisted of a diffuse cloud of positive charge with the negatwe
`electrons embedded randomly h~ it. This model, shown in Fig. 2.9, is often called
`the plum pudding model because the electrons are like raisins dispet sed in a pudding
`(the positivc chargc cloud), as in plum pudding, a favorite Euglish dessert.
`In 1909 Robert Millikan (1868-1953), working at thc University of Chicago,
`perfomaed vcry clever experiments involving charged oil drops. These experiments
`allowed him to determine the magnitude of the electron chargc (see Fig. 2.10). Wifl~
`this value and the charge-to ~nass ratio determined by Thomson, Millikan was able
`to calculate the mass of the election as 9. l 1 × 10-=s gram.
`
`Sphericalcloud
`
`Fis~re 2.9
`One of the early models of the atom
`was Thomson’s plum pudding model,
`in which tile electrons were pictured
`as embedded in a positively charged
`spherical cioud, in tile same way as
`raisins are distributed in an old-
`fashioned plum pudding.
`
`J J. Thornson (1856 !940) was an
`English physicist at Cambridge
`University. He received the Nobel
`Prize in physics in 1906.
`
`FiB~ r~ 2.7
`Schematic of a cathode-ray tube, A
`stream of electrons passes between
`the electrodes. The fastomovlng
`e!ectrons excite the gas in the tube,
`causing a 8low between the
`electrodes.
`
`Appliea
`
`/
`
`(
`
`Figure 2.8
`
`Defledtion of cathode rays by an
`applied electric field.
`
`*~ltllou~4h] I Thomsonisgenerall~’givencreditfarthism°del’~heidea~asapparentlyflrstsaggested
`by the English mathematician at~d physicist William Thomson (better known as Lord Kelvin and not
`
`related to .I.J. Thomson).
`
`38
`
`[ ] Chapter Two Atoms, Molecules, and Ions
`
`Lupin Ex. 1035 (Page 10 of 190)
`
`

`

`on of
`udied
`¯ 2.7)
`
`~ was
`m the
`at the
`: field
`arged
`d the
`~d the
`
`ts the
`
`gative
`called
`dding
`
`icago,
`ments
`.With
`s able
`
`~ggested
`
`Figure 2.10
`
`A schematic representation of the
`apparatus Millikan used to determine
`the charge on the electron. The fall of
`charged oil droplets due to gravity
`can be halted by adjusting the
`voltage across the two plates. The
`voltage and the mass of an oil drop
`can then be used to calculate the
`charge on the oil drop. Millikaffs
`experiments showed that the charge
`on an oil drop is always a whole
`number multiple of the electron
`charge.
`
`(+)
`
`Radioactivity
`In the late nhaeteenth century, it was discovered that certain elements produce high-
`energy radiation. For example, in 1896 the French scientist Henri Becquerel found
`that a piece of a mineral containing uranium could produce its image on a photo-
`graphic plate in the absence of light. IIc attributed this phenomenon to a spontane-
`ous emission of radiation by the uranium, which he called radioactivity. Studies in
`the early twentieth century demonstrated tkree types of radioactive emission:
`gamma (3) rays, beta 03) particles, and alpha (a) particles. A y ray is high-energy
`"light"; a ~ particle is a high-speed electrou; and an c* particle has a 2+ charge,
`that is, a charge twice dmt of the electron and with the opposite sign. The mass of
`an c* particle is 7300 times that of the electron. More modes of radioactivity are now
`l~lown, and we will discuss them in Chapter 21. Ilere we will consider only a
`particles because they were used in some crucial early experiments.
`
`The Nuclear Atom
`In 1911 Eraest Rntherford (Fig. 2.11), who performed many of the pioneering
`experiments to explore radioactivity, carried out an experiment to test Thomson’s
`plum pudding model. The experiment involved directing c* particles at a thin sheet
`of metal hill, as illustrated in Fig. 2.12. Rutherford reasoned that, if Thomson’s
`model were accurate, the massive c~ patricles should crash through the thin foil like
`cannonballs through gauze, as shown in Fig. 2.13(a). He expected the ~ particles to
`travel through the lbil, with, at the most, very minor deflections h~ their paths. The
`results of the experiment were very different from those Ruthcrtord anticipated¯
`Although most of the ~ p,’~ticles passed straight through, many of the particles were
`deflected at large angles, as shown in Fig. 2.13(b), and some were reflected, never
`hitting the detector¯ This outcome was a great surprise to Rutherford. (He wrote that
`this result was comparable to shooting a howitzer at a piece of paper and having the
`shell reflected back.)
`Ruthertbrd knew from these results that the plum pudding model for the atom
`could not be correct. The large deflections of the c~ particles could only be caused by
`a center of concentrated positive charge, as illustrated h~ Fig. 2.13(b). Most of the c*
`particles pass directly through the foil because the atom is mostly open space. The
`deflected c* particles are those that had a "close encounter" with the positive center
`of the atom, and the few reflected r* particles are those that made a "direct hit" on
`the much more massive positive center.
`
`Fis~re 2.11
`
`Ernest Rutberford (1871-1937) was
`born on a farm in New Zealand. In
`1895 he placed second in a
`scholarship competition to at[end
`Cambridge University, but was
`awarded the scholarship when the
`winner decided to stay home and
`get married. As a scientist in England,
`Rutherford did much of the early
`work on characterizing radioactivity.
`He named the ~ and/~ particles and
`the T ray and coined the term half-
`/if~ to describe an important attribute
`of radioactive elements. His
`experiments on the behavior of ~
`particles striking thin metal foils led
`him to postulate the nuciear atom He
`also invented the name proton for
`the nucieus of the hydrogen atom.
`He received the Nobel Prize in
`chemistry in 1908
`
`2.4 Early Experiments to Characterize the Atom ~ 39
`
`Lupin Ex. 1035 (Page 11 of 190)
`
`

`

`Beam of
`alpha particles
`
`Some alpha
`
`(a)
`
`Thin metal foil
`
`Most parUcles pass
`straight through loll
`
`deflect scattcred
`alpha particles
`
`Figure 2.12
`
`Rutherford’s e~perime~ on ~-pa~cicle bombardment of metal foil.
`
`In Rutherford’s mind these results could only be explained in terms of a nu-
`clear atom--an atom with a dense center of positive charge (the nucleus) with
`electrons moving around at a distance that is large relative to the nuclear radius.
`
`The Modern View of Atomic
`2.5 Structure--An Introduction
`PURPOSE
`[] To describe features of subatomic particles.
`To explain the use of the symbol ~X to describe a given atom.
`
`In the years since Thomson and Rutherford, a great deal has been learned about
`atomic structure. Because much of this material will be covered in detail in later
`chapters, only an introduction will be given here. The simplest view of the atom is
`that it consists of a tiny nucleus (with a radius of about 10 is cm) and electrons that
`move about the nucleus at an average distance of about 10-s cm away from it (see
`Fig. 2.14).
`As we will see later, the chemistry of an atom mainly results from its electrons.
`For this reason chemists can be satisfied with a relatively crude nuclear model. The
`nucleus is assumed to contain protons, which have a positive charge equal in
`magnitude to the electron’s negative charge, and neutrons, ~vhich have the same
`mass as a proton but no charge. The,relative masses and charges of the electron,
`proton, and neutron are shown in Table 2.1.
`Two striking things about the nucleus are its small size compared to the overall
`size of the atom and its extremely high density. The tiny nucleus accounts for
`almost all of the atom’s mass. Its great density is dramatically demonstrated by the
`fact that a piece of nuclear material about the size of a pea would have a mass of 250
`million tons!
`
`I~igu re 2.13
`
`(a) The results of the metal foil
`experiment if Thomson’s model were
`correct. (b) Actual results.
`
`Figure 2.14
`
`A nuclear atom viewed in cross
`section.
`
`40
`
`[] Chapter Two Atoms, Molecules, and Ions
`
`Lupin Ex. 1035 (Page 12 of 190)
`
`

`

`The Mass and Charge of the Electron, Proton, and Neutron
`
`Particle
`
`Electron
`Protot~
`NeutroII
`
`Mass
`
`9.11 × 10 2~ g
`1.67 × 10-24 g
`1.67 × 10 24 g
`
`Charge*
`
`1 -
`1 +
`none
`
`*The magnitade of flae charge of the electron 0a~d the proton is 1.60 × 10 ~9 C.
`
`Table 2A
`
`An important question to consider at this point is "lfall atoms are composed of
`these same components, why do different atoms have different chemical proper-
`ties?" The answer to this question lies in the number and arrangement of the elec-
`trons. The electrons comprise most of the atomic volume and thus are the parts that
`"intermingle" when atoms combine to form molecules. Therefore, the number of
`electrons possessed by a given atom ~’eatly affects its ability to interact with other
`atoms. As a result, the atoms of different elements, which have different numbers of
`protons and electrons, show different chemical behavior.
`A sodium atom has 11 protons in its nucleus. Since atoms have no net charge,
`the number of electrons must equal the number of protons. Therefore, a sodium
`atom has 11 electrons moving around its nucleus. It is always true that a sodium
`atom has 11 protons and 11 electrons. However, each sodium atom also has neu-
`trons in its nucleus, and different types of sodium atoms exist that have different
`numbers of neutrons. For example, consider the sodium atoms represented in Fig.
`2.15. These two atoms are |sotopes, or atoms with tlw same number of protons but
`different numbers of neutrons. Note that the symbol for one particular type of
`sodium atom is written
`
`Mass number ~ ~Na "-- Element symbol
`
`Atomic Number/m
`
`where the atomic number Z (number of protons) is written as a subscript and the
`mass number A (the total number of protons and neutrons) is written as a super-
`script. (The particular atom represented here is called "sodium twenty-three." It
`has 11 electrons, 11 protons, and 12 neutrons.) Because the chemistry of an atom is
`due to its electrons, isotopes show almost identical chemical properties. In nature
`elemants are usually found as a mixture of isotopes.
`
`Figure 2.15
`
`Two isotopes of sodium. Both have
`eleven protons and electrons, but
`they differ in the number of neutrons
`in their nuclei.
`
`2.5 The Modern V’ew of Atomic Structure--An Introduction ~] 41
`
`Lupin Ex. 1035 (Page 13 of 190)
`
`with
`dius.
`
`tbout
`later
`)m is
`~ that
`
`(see
`
`rODS.
`The
`al in
`same
`
`tron,
`
`~erall
`s for
`y the
`f 25(I
`
`

`

`Write the symbol for the atom that has an atomic nmnber of 9 and a mass number of
`19. How many electrons and how many neutrons does this atom have?
`
`Solution
`The atomic namber 9 means the atom has 9 protons. This element is called fluorine,
`
`symbolized by F. Tlae atom is represented as
`
`and is called ’ fluonne mneteen.’ Since the atom has 9 protons, it mast also have 9
`electrons to achieve electrical neutrality. The mass number gives the total nmuber of
`protons and ncutrous, which means that this atom has 10 neutrons.
`
`.6 Molecules and Ions
`
`PURPOSE
`U~ To introduce basic ideas of bonding in molecules.
`To show various ways of representing molecules.
`
`From a chemist’s viewpoint the most interesting characteristic of an atom is its
`ability to combine with other atoms to form compounds. It was John Dalton who
`first recognized that chmnical compounds were collections of atoms, but he could
`not determine the structure of atoms or their means tor binding to eac