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C H A P T E R
Atomic
Structure
QUESTIONS YOU MAY HAVE ASKED YOURSELF
1.
2.
3.
4.
If we can’t see atoms, how can we discuss even smaller particles?
What is radioactivity? Is it contagious?
What causes the colors of fireworks?
Most lightbulbs produce white light. Why does a neon bulb give off colored
light?
5. I see the word quantum used; what does it mean?
6. What is “periodic” about the periodic table?
Images of the Invisible
Atoms are exceedingly tiny particles, much too small to see. Yet in this chapter, we discuss
their inner structure. If atoms are so small that we cannot see them, even with an optical
microscope, how can we possibly know anything about their structures?
It turns out that we really know quite a lot about the structure of atoms. Although scientists
have never examined the interior of an atom directly, they have been able to obtain a great deal
of indirect information. By designing some clever experiments and exercising their powers of deduction, they have been able to construct an amazingly detailed model of what an atom’s interior
must be like.
Atoms are much too small to be seen with an ordinary light microscope. However, since
1970 scientists have been able to obtain images of individual atoms. In order to see even rough
images of atoms, they must use special kinds of instruments such as the scanning tunneling microscope (STM). In this way they obtain pictures such as the one shown on page 40. We can
see outlines of atoms in such photographs, and we can tell quite a bit about how they are
arranged, but these pictures tell us nothing about the inner structure of the atoms. Why do we
care about the structure of particles as tiny as atoms? It is the arrangement of various parts of
atoms that determines the properties of different kinds of matter. Only by understanding atomic
structure can we learn how atoms combine to make millions of different substances. With such
knowledge, we can modify and synthesize materials to meet our needs more precisely. Knowledge of atomic structure is even essential to our health. Many medical diagnoses are based on
chemical analyses that have been developed from our understanding of atomic structure.
Perhaps of greater interest to you is the fact that your understanding of chemistry (as well as
much of biology and other sciences) depends, at least in part, on your knowledge of atomic
structure. Let’s start our study of atomic structure by going back to the time of John Dalton.
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3.1
Electricity source
(⫺)
(⫹)
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⫹
⫹
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Electricity and the Atom
John Dalton, who set forth his atomic theory in 1803, regarded the atom as hard and
indivisible. However, it wasn’t long before evidence accumulated to show that matter has electrical properties. Indeed, the electrolytic decomposition of water by
Nicholson and Carlisle in 1800 (Section 2.3) had already indicated this. Electricity
played an important role in unraveling the structure of the atom.
Static electricity has been known since ancient times, but the notion of continuous electric current developed in the nineteenth century. In 1800 Alessandro Volta
invented an electrochemical cell much like a modern battery. If the poles of a cell are
connected by a wire, current flows through the wire. The current is sustained by
chemical reactions inside the cell. Volta’s invention soon was applied in many areas
of science and everyday life.
⫹
Electrolysis
Anode
Cathode
왖 Figure 3.1 An electrolysis apparatus. The electricity source (for example, a battery) directs electrons
through wires from the anode to the
cathode. Cations 1+2 are attracted to
the cathode 1-2, and anions 1 -2 are
attracted to the anode 1+2. This migration of ions is the flow of electricity
through the solution.
왖 In addition to Michael Faraday’s work in
A N S W E R
electrochemistry, he also devised methods
for liquefying gases, and invented the first
electrical transformer.
1. If we can’t see atoms,
how can we discuss even
smaller particles?
The existence of electrons
and other subatomic particles was deduced from the
behavior of electricity. The
existence and behavior of
those subatomic particles
make atoms behave the way
they do.
Soon after Volta’s invention, Humphry Davy (1778–1829), a British chemist, built a
powerful battery that he used to pass electricity through molten (melted) salts.
Davy quickly discovered several new elements. In 1807 he liberated highly reactive
potassium metal from molten potassium hydroxide. Shortly thereafter he produced
sodium metal by passing electricity through molten sodium hydroxide. Within a
year, Davy had also produced magnesium, strontium, barium, and calcium metals
for the first time. The science of electrochemistry was born. (We will study electrochemistry in more detail in Chapter 8.)
Davy’s protégé Michael Faraday (1791–1867) greatly extended this new science.
Faraday defined many of the terms we still use today, such as electrolysis, the splitting of compounds by electricity (Figure 3.1). Lacking in his own formal education,
he consulted English classical scholar William Whewell (1794–1866), who suggested the name electrolyte for a compound that conducts electricity when melted or
dissolved in water. In the electrolysis apparatus, electrodes, carbon rods or metal
strips inserted into a molten compound or solution, carry the electric current. The
electrode that bears a positive charge is the anode, and the negatively charged electrode is the cathode. The entities that carry the electric current through a melted
compound or solution are called ions. An ion is an atom or a group of atoms bonded together that has an electric charge. An ion with a negative charge is an anion; it
travels toward the anode. A positively charged ion is a cation; it moves toward the
cathode.
Faraday’s work established that atoms are electrical in nature, but further details of atomic structure had to wait several decades for more powerful sources of
electrical voltage and better vacuum pumps.
Cathode Ray Tubes
In 1875 William Crookes (1832–1919), an English chemist, was able to construct a
low-pressure gas discharge tube that would allow electricity to pass through it. His
experiment is shown in Figure 3.2. Metal electrodes are sealed in the tube. It is connected to a vacuum pump, and most of the air is removed. A beam of current is seen
as a green fluorescence, observed when the beam strikes a screen coated with zinc
sulfide. This beam, which seems to leave the cathode and travel to the anode, is
called a cathode ray.
Thomson’s Experiment: Mass-to-Charge Ratio
Considerable speculation arose as to the nature of cathode rays and many experiments were undertaken. Were these rays actually beams of particles, or did they
consist of a wavelike form of energy much like visible light? The answer came (as
scientific answers should) from an experiment performed by the English physicist
Joseph John Thomson in 1897. Thomson showed that cathode rays were deflected
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3.1 Electricity and the Atom
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Negative plate
(⫺)
(⫹)
High
voltage
Cathode
Slit
Screen coated
with zinc
sulfide
(⫺)
Anode
(⫹)
Positive plate
(a)
(b)
왖 Figure 3.2 Thomson’s apparatus, showing deflection of cathode rays (a beam of electrons). Cathode rays are themselves invisible but are observed through the green fluorescence produced when they strike a zinc sulfide–coated screen.
The diagram (a) shows deflection of the beam in an electric field. The photograph (b) shows the deflection in a magnetic
field. The magnetic field is created by the magnet to the right of and slightly behind the screen. Cathode rays travel in
straight lines unless some kind of external field is applied.
in an electric field (see again Figure 3.2). The beam was attracted to the positive
plate and repelled by the negative plate. Thomson therefore concluded that cathode
rays consisted of negatively charged particles. His experiments also showed that
the particles were the same regardless of the materials from which the electrodes
were made or the type of gas in the tube. He concluded that these negative particles
are part of all kinds of atoms. Thomson named these negatively charged units
electrons. Cathode rays, then, are beams of electrons emanating from the cathode of
a gas discharge tube.
Cathode rays are deflected in magnetic fields as well as in electric fields. The
greater the charge on a particle, the more it is deflected in an electric (or magnetic)
field. The heavier the particle, the less it is deflected by a force. By measuring the
amount of deflection in fields of known strength, Thomson was able to calculate the
ratio of the mass of the electron to its charge. He could not measure either the mass
or the charge separately. This is like knowing that each 1-ft length of a steel beam
has a mass of 25 lb. With this information alone, we cannot find either the total mass
or length of a beam. Once the beam’s mass or its length is known, it is easy to calculate the other from the known value and the 25 lb/1 ft ratio. Thomson was awarded
the Nobel Prize in Physics in 1906.
(⫺) (⫹)
High
voltage
Goldstein’s Experiment: Positive Particles
In 1886 German scientist Eugen Goldstein performed experiments with gas discharge tubes that had perforated cathodes (Figure 3.3). He found that although
electrons were formed and sped off toward the anode as usual, positive particles
were also formed and shot in the opposite direction toward the cathode. Some of
these positive particles went through the holes in the cathode. In 1907 a study of the
deflection of these particles in a magnetic field indicated that they were of varying
mass. The lightest particles, formed when there was a little hydrogen gas in the
tube, were later shown to have a mass 1837 times that of an electron.
Millikan’s Oil-Drop Experiment: Electron Charge
The charge on the electron was determined in 1909 by Robert A. Millikan
(1868–1953), a physicist at the University of Chicago. Millikan observed electrically
charged oil drops in an electric field. A diagram of his apparatus is shown in
⫹
⫹
Cathode
⫺
⫺
⫹
⫺
Anode
왖 Figure 3.3 Goldstein’s apparatus for the study of positive particles.
Some positive ions, attracted toward
the cathode, pass through the holes in
the cathode. The deflection of these
particles (not shown) in a magnetic
field can be studied in the region to the
left of the cathode.
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Atomic Structure
왘 Figure 3.4 The Millikan oildrop experiment. Oil drops irradiated
with X-rays pick up electrons and become negatively charged. Their fall due
to gravity can be balanced by adjusting
the voltage of the electric field. The
charge on the oil drop can be determined from the applied voltage and the
mass of the oil drop. The charge on
each drop is that of some whole number of electrons.
Atomizer
(⫹)
Positively charged
plate
Source of
ionizing
radiation
Telescope
(⫺)
You can find out more about
numbers such as 9.1 * 10-28
and 1 * 1027 in the Appendix.
However, in practice we seldom
use the actual mass and charge
of the electron. For many purposes, the electron is considered the unit of electrical charge.
The charge is shown as a superscript minus sign (meaning 1- ).
In indicating charges on ions, we
use - to indicate a net charge of
one electron, 2- to indicate a
charge of two electrons, and
so on.
Negatively charged
plate
Figure 3.4. A spray bottle is used to form tiny droplets of oil. Some acquire negative
charges by picking up electrons from the friction generated as the particles rub
against the opening of the spray nozzle and against each other. (The charge is static
electricity, just like the charge you get from walking across a nylon carpet.) The negative droplets can acquire one or more extra electrons by this process. Charges can
also be produced by irradiation with X-rays.
Some of the oil droplets pass into a chamber where they can be viewed through
a microscope. The negative plate at the bottom of the chamber repels the negatively
charged droplets; the positive plate attracts them. By manipulating the charge on
each plate and observing the behavior of the droplets, the charge on each droplet
can be determined. Millikan took the smallest possible difference in charge between
two droplets to be the charge of an individual electron. For his research, he received
the Nobel Prize in Physics in 1923.
From Millikan’s value for the charge and Thomson’s value for the mass-tocharge ratio, the mass of the electron was readily calculated. That mass was found
to be only 9.1 * 10-28 g. It would take more than 1 * 1027 electrons to weigh one
gram—that’s a 1 followed by 27 zeros, or a billion billion billion electrons. More important, that mass is much smaller than that of the lightest atom. This means that
electrons are much smaller than atoms.
Self-Assessment Questions
1. The use of electricity to cause a chemical reaction is called
a. anodizing
b. corrosion
c. electrolysis
d. hydrolysis
2. During electrolysis of molten sodium hydroxide, sodium cations move
to the
a. anode, which is negatively charged
b. anode, which is positively charged
c. cathode, which is negatively charged
d. cathode, which is positively charged
3. Cathode rays are not
a. composed of negatively charged particles
b. composed of particles with a mass of 1 u
c. deflected by electric fields
d. the same from element to element
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3.2 Serendipity in Science: X-Rays and Radioactivity
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4. Thomson used a cathode ray tube to discover that cathode rays
a. consist of radiation similar to X-rays
b. could be deflected by an electric field
c. could ionize air inside the tube
d. could ionize the air outside the far end of the tube
5. When J. J. Thomson discovered the electron, he measured its
a. atomic number
b. charge
c. mass
d. mass-to-charge ratio
6. In his oil-drop experiment, Millikan determined the charge of an electron by
a. measuring the diameter of the tiniest oil drops
b. noting that the drops had either a certain minimum charge or multiples
of that charge
c. using an atomizer that could spray out individual electrons
d. viewing the smallest charged drops microscopically and noting they
had one extra electron
Answers: 1, c; 2, c; 3, b; 4, b; 5, d; 6, b
3.2
Serendipity in Science:
X-Rays and Radioactivity
Let’s return now to the structure of the atom and look at a little scientific
serendipity—scientific discoveries that are “happy accidents.” Have you ever wondered why these accidents often seem to happen to scientists? It is probably because
scientists are trained observers. The same accident could happen right before the
eyes of an untrained person and go unnoticed (though often even scientists miss
important finds at first). Or, if noticed, its significance might not be grasped.
Roentgen: The Discovery of X-Rays
Two serendipitous discoveries in the last years of the nineteenth century profoundly changed the world. In 1895 German scientist Wilhelm Conrad Roentgen
(1845–1923) was working in a dark room, studying the glow produced in certain
substances by cathode rays. To his surprise, he noted this glow on a chemically
treated piece of paper some distance from the cathode ray tube. The paper even
glowed when taken into the next room. Roentgen had discovered a new type of ray
that could travel through walls. When he waved his hand between the radiation
source and the glowing paper, he suddenly was able to see the bones of his own
hand through the paper. He called these mysterious rays, which seemed to make
his flesh disappear, X-rays.
X-rays are a form of electromagnetic radiation—energy with electric and magnetic components. Among other types, electromagnetic radiation also includes visible light, radio waves, microwaves, infrared radiation, ultraviolet light, and gamma
rays (Section 3.3). These forms differ in energy, with radio waves lowest in energy,
followed by microwaves, infrared radiation, visible light, ultraviolet light, X-rays,
and gamma rays.
Today X-rays are one of the most widely used tools in the world for medical diagnosis. Not only are they employed for examining decayed teeth, broken bones, and
diseased lungs, but they are also the basis for such procedures as mammography and
computerized tomography (Chapter 11). In the United States alone, payment for various radiological procedures totals more than $20 billion each year. How ironic that
Roentgen himself made no profit at all from his discovery. He considered X-rays a “gift
to humanity” and refused to patent any part of the discovery. However, he did receive
much popular acclaim and in 1901 was awarded the first Nobel Prize in Physics.
왖 X-rays were used in medicine shortly
after they were discovered by Wilhelm
Roentgen (1845–1923) in 1895. Michael
Purpin of Columbia University made this
X-ray in 1896 to aid in the surgical removal
of gunshot pellets (the dark spots) from the
hand of a patient.
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The Discovery of Radioactivity
A N S W E R
2. What is radioactivity? Is
it contagious?
Radioactivity is the spontaneous decay (occurs by
itself) of an atom into one
or more other atoms. The
radiation—alpha, beta, or
gamma rays—given off by
radioactive material is
hazardous but you can no
more “catch” radiation
sickness from a person
exposed to radiation than
you can “catch” a sunburn
from your sunburned friend.
Self-Assessment Questions
1. When Roentgen took an X-ray photograph of a hand, the bones showed
up slightly dark because the X-rays were
a. absorbed by bone more so than by flesh
b. light that is less energetic than visible light
c. made up of fast-moving atomic nuclei
d. a type of radioactivity
2. Radioactivity arises from elements that
a. absorb radiofrequency energy and release stronger radiation
b. are unstable and release radiation
c. give off X-rays
d. undergo fission when they absorb energy
Answers: 1, a; 2, b
왖 Marie Sklodowska Curie in her laboratory
and Marie and Pierre Curie on a French
postage stamp.
Certain chemicals exhibit fluorescence after exposure to strong sunlight; they continue to glow even when taken into a dark room. In 1895 Antoine Henri Becquerel
(1852–1908), a French physicist, was studying fluorescence by wrapping photographic film in black paper, placing a few crystals of the fluorescing chemical on top
of the paper, and then placing the package in strong sunlight. If the glow was like
ordinary light, it would not pass through the paper. On the other hand, if similar to
X-rays, it would pass through the black paper and fog the film.
While working with a uranium compound, Becquerel made an important accidental discovery. When placed in sunlight, the compound fluoresced and fogged
the film. On several cloudy days when exposure to sunlight was not possible, he
prepared samples and placed them in a drawer. To his great surprise, the photographic film was fogged even though the uranium compound had not been exposed to sunlight. Further experiments showed that the radiation coming from the
uranium compound was unrelated to fluorescence but was a characteristic of the element uranium.
Other scientists immediately began to study this new radiation. Becquerel had
a graduate student from Poland, Marie Sklodowska, who gave the phenomenon a
name: radioactivity. Radioactivity is the spontaneous emission of radiation from
certain unstable elements. Marie later married Pierre Curie, a French physicist. Together they discovered the radioactive elements polonium and radium, and with
Becquerel they shared the 1903 Nobel Prize in Physics.
After her husband’s death in 1906, Marie Curie continued to work with radioactive substances, winning the Nobel Prize in Chemistry in 1911. For more than
50 years she was the only person ever to have received two Nobel Prizes.
3.3
Three Types of Radioactivity
Scientists soon showed that three types of radiation emanated from various radioactive elements. Ernest Rutherford (1871–1937), a New Zealander who spent his
career in Canada and Great Britain, chose the names alpha, beta, and gamma for the
three types of radiation. When passed through a strong magnetic or electric field,
the alpha form was deflected in a manner indicating that it consisted of a beam of
positive particles (Figure 3.5). Later experiments showed that an alpha particle has
a mass four times that of a hydrogen atom and a charge twice the magnitude of, but
opposite in sign to, that of an electron. An alpha particle is in fact identical to the
nucleus (Section 3.4) of a helium atom and is often symbolized by He 2 + .
The beta radiation was shown to be made up of negatively charged particles
identical to those of cathode rays. Therefore, a beta particle is an electron.
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3.4 Rutherford’s Experiment: The Nuclear Model of the Atom
Gamma rays are not deflected by a magnetic field.
They are a form of electromagnetic radiation, much like
the X-rays used in medical work but even more energetic
and more penetrating. The three types of radioactivity
are summarized in Table 3.1.
The discoveries of the late nineteenth century paved
the way for an entirely new picture of the atom, which
developed rapidly during the early years of the twentieth
century.
Radioactive
material
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Photographic plate
(⫺)
α
γ
β
Table 3.1 Types of Radioactivity
Lead block
Name
Symbol
Mass (u)
Charge
Alpha
a
4
2+
Beta
b
1-
Gamma
g
1
1837
0
(⫹)
왖 Figure 3.5 Behavior of radioactive rays in an electric field.
0
Self-Assessment Questions
1. The mass of an alpha particle is about
a. four times that of a hydrogen atom
b. the same as that of an electron
c. the same as that of a hydrogen atom
d. the same as that of a beta particle
2. Compared to the charge on an electron (designated -1), that on an alpha
particle is about
a. -4
b. -1
c. +1
d. +2
e. +4
3. A beta particle is a(n)
a. boron nucleus
b. electron
c. helium nucleus
d. proton
4. The mass of a gamma ray is
a. -1 u
b. 0 u
c. 1 u
d. 4 u
Answers: 1, a; 2, d; 3, b; 4, b
3.4
Rutherford’s Experiment:
The Nuclear Model of the Atom
At Rutherford’s suggestion, two of his coworkers, Hans Geiger (1882–1945), a German physicist, and Ernest Marsden (1889–1970), an English undergraduate student,
bombarded very thin metal foils with alpha particles from a radioactive source
(Figure 3.6). In an experiment with gold foil, most of the particles behaved as
Rutherford expected, going right through the foil with little or no scattering. However, a few particles were deflected sharply. Occasionally one was sent right back in
the direction from which it had come! Rutherford had assumed the positive charge
to be spread evenly over all the space occupied by the atom, but obviously it was
not. To explain the experiment, Rutherford concluded that all the positive charge
Alpha
source
Alpha
radiation
Zinc sulfide screen for
detecting radiation
Gold
foil
왖 Figure 3.6 Rutherford’s goldfoil experiment. Most alpha particles
passed right through the gold foil, but
now and then a particle was deflected.
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Atomic Structure
왘 Figure 3.7 Model explaining
the results of Rutherford’s gold-foil experiment. Most of the alpha particles
pass right through the foil because it is
mainly empty space. But some alpha
particles are deflected as they pass
close to a dense, positively charged
atomic nucleus. Once in a while an
alpha particle approaches an atomic
nucleus head-on and is knocked back
in the direction from which it came.
Most of an atom
is empty space
Alpha source
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Tiny, extremely dense
nucleus deflects
only a few alpha
particles
and nearly all the mass of an atom are concentrated at the center of the atom in a
tiny core called the nucleus.
When an alpha particle, which is positively charged, directly approached the
positively charged nucleus, it was strongly repelled and, therefore, sharply deflected (Figure 3.7). Because only a few alpha particles were deflected, Rutherford concluded that the nucleus must occupy only a tiny fraction of the volume of an atom.
Most of the alpha particles passed right through because most of an atom is empty
space. The space outside the nucleus isn’t completely empty, however. It contains
the negatively charged electrons. Rutherford concluded that the electrons had so little mass that they were no match for the alpha “bullets.” It would be analogous to a
mouse trying to stop the charge of a bull elephant.
Rutherford’s nuclear theory of the atom, set forth in 1911, was revolutionary.
He postulated that all the positive charge and nearly all the mass of an atom are
concentrated in a tiny, tiny nucleus. The negatively charged electrons have almost
no mass, yet they occupy nearly all the volume of an atom. To picture Rutherford’s
model, visualize a sphere as big as a giant indoor football stadium. The nucleus at
the middle of the sphere is as small as a pea but weighs several million tons. A few
flies flitting here and there throughout the sphere represent the electrons.
Self-Assessment Questions
1. The most persuasive evidence found in Rutherford’s gold foil experiment
was that
a. most of the alpha particles went right through the foil, but a tiny fraction
was deflected right back toward the source
b. some alpha particles were deflected a bit
c. some gold atoms were converted to lead atoms
d. some gold atoms were split
2. Rutherford’s gold foil experiment showed that
a. the gold atoms consisted of protons in a “pudding” of electrons
b. the nucleus was huge and was positively charged
c. the nucleus was tiny and positively charged
d. some gold atoms were shattered by alpha particles
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3.5 The Atomic Nucleus
3. From his gold foil experiment, Rutherford concluded that atoms consist
mainly of
a. electrons
b. empty space
c. neutrons
d. protons
4. The positively charged part of the atom is its
a. cation
b. electron cloud
c. nucleus
d. valence shell
Answers: 1, a; 2, c; 3, b; 4, c
3.5
The Atomic Nucleus
In 1914 Rutherford suggested that the smallest positive particle (the one formed
when there is hydrogen gas in the Goldstein apparatus—see Figure 3.3) is the unit
of positive charge in the nucleus. This particle, called a proton, has a charge equal
in magnitude to that of the electron and has nearly the same mass as a hydrogen
atom. Rutherford’s suggestion was that protons constitute the positively charged
matter in all atoms. The nucleus of a hydrogen atom consists of one proton, and the
nuclei of larger atoms contain greater numbers of protons.
Except for hydrogen atoms, atomic nuclei are heavier than indicated by the number
of positive charges (number of protons). For example, the helium nucleus has a charge
of 2+ (and, therefore, two protons, according to Rutherford’s theory), but its mass is four
times that of hydrogen. This excess mass puzzled scientists at first. But in 1932 English
physicist James Chadwick (1891–1974) discovered a particle with about the same mass
as a proton but with no electric charge. It was called a neutron, and its existence explains
the unexpectedly high mass of the helium nucleus. Whereas the hydrogen nucleus contains only one proton of mass 1 u, the helium nucleus contains not only two protons
(2 u) but also two neutrons (2 u), giving the nucleus a total mass of 4 u.
With the discovery of the neutron, the list of “building blocks” we will need for
“constructing” atoms is complete. The properties of these particles are summarized
in Table 3.2. (There are dozens of other subatomic particles, but most exist only momentarily and are not important in our discussion here.)
The number of protons in the nucleus of an atom of any element is the atomic
number (Z) of that element. This number determines the kind of atom—that is, the
identity of the element—and it is found on any periodic table or list of the elements.
An element, then, is a substance in which all the atoms have the same atomic number. Dalton had said that the mass of an atom determines the element. We now
know it is not the mass but the number of protons that determines the identity of an
element. For example, an atom with 26 protons (one whose atomic number Z = 26)
is an atom of iron (Fe). An atom with 50 protons 1Z = 502 is an atom of tin (Sn). In
a neutral atom (without an electric charge) the positive charge of the protons is
exactly neutralized by the negative charge of the electrons. The attractive forces
between the unlike charges help hold the atom together.
Table 3.2 Subatomic Particles
Particle
Symbol
Mass (u)
Charge
Proton
p+
1
1+
Nucleus
Neutron
n
1
0
Nucleus
e-
1
1837
1-
Electron
Location in Atom
Outside nucleus
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Atomic Structure
A proton and a neutron have almost the same mass, 1.0073 u and 1.0087 u, respectively. This is equivalent to saying that two different people weigh 100.7 kg and
100.9 kg. The difference is so small that it usually can be ignored. Thus, for many
purposes, we assume the masses of the proton and the neutron to be the same, 1 u.
The proton has a charge equal in magnitude but opposite in sign to that of an electron. This charge on a proton is written as 1+. The electron has a charge of 1- and a
mass of 0.00055 u. The electrons in an atom contribute so little to its total mass that
their mass is usually disregarded and treated as if it were 0.
Isotopes
Atoms of a given element can have different numbers of neutrons in their nuclei.
For example, most hydrogen atoms have a nucleus consisting of a single proton and
no neutrons. However, about 1 hydrogen atom in 6700 has a neutron as well as a
proton in the nucleus. This heavier hydrogen atom is called deuterium. Whether it
has one neutron, two neutrons, or none, any atom with Z = 1—that is, with one
proton—is a hydrogen atom. Atoms that have this sort of relationship—having the
same number of protons but different numbers of neutrons—are called isotopes
(Figure 3.8). A third, rare isotope of hydrogen is tritium, which has two neutrons
and one proton in the nucleus.
Most, but not all, elements exist in nature in isotopic forms. For example, tin
(Sn) is present in nature in 10 different isotopic forms. It also has 15 radioactive isotopes that do not occur in nature. The existence of isotopes requires a modification
of Dalton’s original theory. He said that all atoms of the same element are alike. We
now say that all atoms of the same element have the same number of protons. Different isotopes of an element have atoms with the same number of protons but with
different numbers of neutrons (and, therefore, different masses).
Isotopes usually are of little importance in ordinary chemical reactions. All
three hydrogen isotopes react with oxygen to form water. Because the isotopes differ in mass, compounds formed with different hydrogen isotopes have different
physical properties, but such differences are usually slight. For example, water in
which both hydrogen atoms are deuterium is called heavy water, often represented
as D2O. Heavy water boils at 101.4 °C (instead of 100 °C) and freezes at 3.8 °C
(instead of 0 °C). In nuclear reactions, however, isotopes are of utmost importance,
as we shall see in Chapter 11.
Symbols for Isotopes
Collectively, the two principal nuclear particles, protons and neutrons, are called
nucleons. Isotopes are represented by symbols with subscripts and superscripts.
A
ZX
In this general symbol, Z is the nuclear charge (atomic number or number of
protons), and A is the mass number or the nucleon number because it is the num-
e⫺
왘 Figure 3.8 The three isotopes
of hydrogen. Each has one proton and
one electron, but they differ in the
number of neutrons in the nucleus.
QUESTION: What is the atomic
number and mass number of
each of the isotopes? What is
the mass of each, in atomic
mass units, to the nearest
whole number?
e⫺
e⫺
p⫹
p⫹ n
p⫹
n n
Protium
(ordinary hydrogen)
Deuterium
(heavy hydrogen)
Tritium
(radioactive hydrogen)