Atom
| I | INTRODUCTION |
Atom, tiny basic building block of matter. All the
material on Earth is composed of various combinations of atoms. Atoms are the
smallest particles of a chemical element that still exhibit all the chemical
properties unique to that element. A row of 100 million atoms would be only
about a centimeter long. See also Chemical Element.
Understanding atoms is key to understanding the
physical world. More than 100 different elements exist in nature, each with its
own unique atomic makeup. The atoms of these elements react with one another and
combine in different ways to form a virtually unlimited number of chemical
compounds. When two or more atoms combine, they form a molecule. For example,
two atoms of the element hydrogen (abbreviated H) combine with one atom of the
element oxygen (O) to form a molecule of water (H20).
Since all matter—from its formation in the
early universe to present-day biological systems—consists of atoms,
understanding their structure and properties plays a vital role in physics,
chemistry, and medicine. In fact, knowledge of atoms is essential to the modern
scientific understanding of the complex systems that govern the physical and
biological worlds. Atoms and the compounds they form play a part in almost all
processes that occur on Earth and in space. All organisms rely on a set of
chemical compounds and chemical reactions to digest food, transport energy, and
reproduce. Stars such as the Sun rely on reactions in atomic nuclei to produce
energy. Scientists duplicate these reactions in laboratories on Earth and study
them to learn about processes that occur throughout the universe.
Throughout history, people have sought to
explain the world in terms of its most basic parts. Ancient Greek philosophers
conceived of the idea of the atom, which they defined as the smallest possible
piece of a substance. The word atom comes from the Greek word meaning
“not divisible.” The ancient Greeks also believed this fundamental particle was
indestructible. Scientists have since learned that atoms are not indivisible but
made of smaller particles, and atoms of different elements contain different
numbers of each type of these smaller particles.
| II | THE STRUCTURE OF THE ATOM |
Atoms are made of smaller particles, called
electrons, protons, and neutrons. An atom consists of a cloud of electrons
surrounding a small, dense nucleus of protons and neutrons. Electrons and
protons have a property called electric charge, which affects the way they
interact with each other and with other electrically charged particles.
Electrons carry a negative electric charge, while protons have a positive
electric charge. The negative charge is the opposite of the positive charge,
and, like the opposite poles of a magnet, these opposite electric charges
attract one another. Conversely, like charges (negative and negative, or
positive and positive) repel one another. The attraction between an atom’s
electrons and its protons holds the atom together. Normally, an atom is
electrically neutral, which means that the negative charge of its electrons is
exactly equaled by the positive charge of its protons.
The nucleus contains nearly all of the mass of
the atom, but it occupies only a tiny fraction of the space inside the atom. The
diameter of a typical nucleus is only about 1 × 10-14 m (4 ×
10-13 in), or about 1/100,000 of the diameter of the entire atom. The
electron cloud makes up the rest of the atom’s overall size. If an atom were
magnified until it was as large as a football stadium, the nucleus would be
about the size of a grape.
| A | Electrons |
Electrons are tiny, negatively charged
particles that form a cloud around the nucleus of an atom. Each electron carries
a single fundamental unit of negative electric charge, or –1.
The electron is one of the lightest
particles with a known mass. A droplet of water weighs about a billion, billion,
billion times more than an electron. Physicists believe that electrons are one
of the fundamental particles of physics, which means they cannot be split into
anything smaller. Physicists also believe that electrons do not have any real
size, but are instead true points in space—that is, an electron has a radius of
zero.
Electrons act differently than everyday
objects because electrons can behave as both particles and waves. Actually, all
objects have this property, but the wavelike behavior of larger objects, such as
sand, marbles, or even people, is too small to measure. In very small particles
wave behavior is measurable and important. Electrons travel around the nucleus
of an atom, but because they behave like waves, they do not follow a specific
path like a planet orbiting the Sun does. Instead they form regions of negative
electric charge around the nucleus. These regions are called orbitals, and they
correspond to the space in which the electron is most likely to be found. As we
will discuss later, orbitals have different sizes and shapes, depending on the
energy of the electrons occupying them.
| B | Protons and Neutrons |
Protons carry a positive charge of +1,
exactly the opposite electric charge as electrons. The number of protons in the
nucleus determines the total quantity of positive charge in the atom. In an
electrically neutral atom, the number of the protons and the number of electrons
are equal, so that the positive and negative charges balance out to zero. The
proton is very small, but it is fairly massive compared to the other particles
that make up matter. A proton’s mass is about 1,840 times the mass of an
electron.
Neutrons are about the same size as protons
but their mass is slightly greater. Without neutrons present, the repulsion
among the positively charged protons would cause the nucleus to fly apart.
Consider the element helium, which has two protons in its nucleus. If the
nucleus did not contain neutrons as well, it would be unstable because of the
electrical repulsion between the protons. (The process by which neutrons hold
the nucleus together is explained below in the Strong Force section of this
article.) A helium nucleus needs either one or two neutrons to be stable. Most
atoms are stable and exist for a long period of time, but some atoms are
unstable and spontaneously break apart and change, or decay, into other
atoms.
Unlike electrons, which are fundamental
particles, protons and neutrons are made up of other, smaller particles called
quarks. Physicists know of six different quarks. Neutrons and protons are made
up of up quarks and down quarks—two of the six different kinds of
quarks. The fanciful names of quarks have nothing to do with their properties;
the names are simply labels to distinguish one quark from another.
Quarks are unique among all elementary
particles in that they have electric charges that are fractions of the
fundamental charge. All other particles have electric charges of zero or of
whole multiples of the fundamental charge. Up quarks have electric charges of
+’. Down quarks have
charges of -€. A proton
is made up of two up quarks and a down quark, so its electric charge is ’ + ’ - €, for a total charge of +1.
A neutron is made up of an up quark and two down quarks, so its electric charge
is ’ - € - €, for a net charge of zero.
Physicists believe that quarks are true fundamental particles, so they have no
internal structure and cannot be split into something smaller.
| III | PROPERTIES OF ATOMS |
Atoms have several properties that help
distinguish one type of atom from another and determine how atoms change under
certain conditions.
| A | Atomic Number |
Each element has a unique number of
protons in its atoms. This number is called the atomic number (abbreviated Z).
Because atoms are normally electrically neutral, the atomic number also
specifies how many electrons an atom will have. The number of electrons, in
turn, determines many of the chemical and physical properties of the atom. The
lightest atom, hydrogen, has an atomic number equal to one, contains one proton,
and (if electrically neutral) one electron. The most massive stable atom found
in nature is bismuth (Z = 83). More massive unstable atoms also exist in nature,
but they break apart and change into other atoms over time. Scientists have
produced even more massive unstable elements in laboratories.
| B | Mass Number |
The total number of protons and neutrons
in the nucleus of an atom is the mass number of the atom (abbreviated A). The
mass number of an atom is an approximation of the mass of the atom. The
electrons contribute very little mass to the atom, so they are not included in
the mass number. A stable helium atom can have a mass number equal to three (two
protons plus one neutron) or equal to four (two protons plus two neutrons).
Bismuth, with 83 protons, requires 126 neutrons for stability, so its mass
number is 209 (83 protons plus 126 neutrons).
| C | Atomic Mass and Weight |
Scientists usually measure the mass of an
atom in terms of a unit called the atomic mass unit (abbreviated amu). They
define an amu as exactly 1/12 the mass of an atom of carbon with six protons and
six neutrons. On this scale, the mass of a proton is 1.00728 amu and the mass of
a neutron is 1.00866 amu. The mass of an atom measured in amu is nearly equal to
its mass number.
Scientists can use a device called a mass
spectrometer to measure atomic mass. A mass spectrometer removes one or more
electrons from an atom. The electrons are so light that removing them hardly
changes the mass of the atom at all. The spectrometer then sends the atom
through a magnetic field, a region of space that exerts a force on magnetic or
electrically charged particles. Because of the missing electrons, the atom has
more protons than electrons and hence a net positive charge. The magnetic field
bends the path of the positively charged atom as it moves through the field. The
amount of bending depends on the atom’s mass. Lighter atoms will be affected
more strongly than heavier atoms. By measuring how much the atom’s path curves,
a scientist can determine the atom’s mass.
The atomic mass of an atom, which depends
on the number of protons and neutrons present, also relates to the atomic weight
of an element. Weight usually refers to the force of gravity on an object, but
atomic weight is really just another way to express mass. An element’s atomic
weight is given in grams. It represents the mass of one mole (6.02 ×
1023 atoms) of that element. Numerically, the atomic weight and the
atomic mass of an element are the same, but the first is expressed in grams and
the second is in atomic mass units. So, the atomic weight of hydrogen is 1 gram
and the atomic mass of hydrogen is 1 amu.
| D | Isotopes |
Atoms of the same element that differ in
mass number are called isotopes. Since all atoms of a given element have the
same number of protons in their nucleus, isotopes must have different numbers of
neutrons. Helium, for example, has an atomic number of 2 because of the two
protons in its nucleus. But helium has two stable isotopes—one with one neutron
in the nucleus and a mass number equal to three and another with two neutrons
and a mass number equal to four.
Scientists attach the mass number to an
element’s name to differentiate between isotopes. Under this convention, helium
with a mass number of three is called helium-3, and helium with a mass number of
four is called helium-4. Helium in its natural form on Earth is a mixture of
these two isotopes. The percentage of each isotope found in nature is called the
isotope’s isotopic abundance. The isotopic abundance of helium-3 is very small,
only 0.00014 percent, while the abundance of helium-4 is 99.99986 percent. This
means that only about one of every 1 million helium atoms is helium-3, and the
rest are all helium-4. Bismuth has only one naturally occurring stable isotope,
bismuth-209. Bismuth-209’s isotopic abundance is therefore 100 percent. The
element with the largest number of stable isotopes found in nature is tin, which
has ten stable isotopes.
All elements also have unstable isotopes,
which are more susceptible to breaking down, or decaying, than are the other
isotopes of an element. When atoms decay, the number of protons in their nucleus
changes. Since the number of protons in the nucleus of an atom determines what
element that atom belongs to, this decay changes one element into another.
Different isotopes decay at different rates. One way to measure the decay rate
of an isotope is to find its half-life. An isotope’s half-life is the
time that passes until half of a sample of an isotope has decayed.
The various isotopes of a given element
have nearly identical chemical properties and many similar physical properties.
They differ, of course, in their mass. The mass of a helium-3 atom, for example,
is 3.016 amu, while the mass of a helium-4 atom is 4.003 amu.
Usually scientists do not specify the
atomic weight of an element in terms of one isotope or another. Instead, they
express atomic weight as an average of all of the naturally occurring isotopes
of the element, taking into account the isotopic abundance of each. For example,
the element copper has two naturally occurring isotopes: copper-63, with a mass
of 62.930 amu and an isotopic abundance of 69.2 percent, and copper-65, with a
mass of 64.928 amu and an abundance of 30.8 percent. The average mass of
naturally occurring copper atoms is equal to the sum of the atomic mass for each
isotope multiplied by its isotopic abundance. For copper, it would be (62.930
amu x 0.692) + (64.928 amu x 0.308) = 63.545 amu. The atomic weight of copper is
therefore 63.545 g.
| E | Radioactivity |
About 300 combinations of protons and
neutrons in nuclei are stable enough to exist in nature. Scientists can produce
another 3,000 nuclei in the laboratory. These nuclei tend to be extremely
unstable because they have too many protons or neutrons to stay in one piece for
long. Unstable nuclei, whether naturally occurring or created in the laboratory,
break apart or change into stable nuclei through a variety of processes known as
radioactive decays (see Radioactivity).
Some nuclei with an excess of protons
simply eject a proton. A similar process can occur in nuclei with an excess of
neutrons. A more common process of decay is for a nucleus to simultaneously
eject a cluster of 2 protons and 2 neutrons. This cluster is actually the
nucleus of an atom of helium-4, and this decay process is called alpha decay.
Before scientists identified the ejected particle as a helium-4 nucleus, they
called it an alpha particle. Helium-4 nuclei are still sometimes called alpha
particles.
The most common way for a nucleus to get
rid of excess protons or neutrons is to convert a proton into a neutron or a
neutron into a proton. This process is known as beta decay. The total electric
charge before and after the decay must remain the same. Because protons are
electrically charged and neutrons are not, the reaction must involve other
charged particles. For example, a neutron can decay into a proton, an electron,
and another particle called an electron antineutrino. The neutron has no charge,
so the charge at the beginning of the reaction is zero. The proton has an
electric charge of +1 and the electron has an electric charge of –1. The
antineutrino is a tiny particle with no electric charge. The electric charges of
the proton and electron cancel each other, leaving a net charge of zero. The
electron is the most easily detected product of this type of beta decay, and
scientists called these products beta particles before they identified them as
electrons.
Beta decay also results when a proton
changes to a neutron. The end result of this decay must have a charge of +1 to
balance the charge of the initial proton. The proton changes into a neutron, an
anti-electron (also called a positron), and an electron neutrino. A positron is
identical to an electron, except the positron has an electric charge of +1. The
electron neutrino is a tiny, electrically neutral particle. The difference
between the antineutrino in neutron-proton beta decay and the neutrino in
proton-neutron beta decay is very subtle—so subtle that scientists have yet to
prove that a difference actually exists.
While scientists often create unstable
nuclei in the laboratory, several radioactive isotopes also occur naturally.
These atoms decay more slowly than most of the radioactive isotopes created in
laboratories. If they decayed too rapidly, they wouldn’t stay around long enough
for scientists to find them. The heavy radioactive isotopes found on Earth
formed in the interiors of stars more than 5 billion years ago. They were part
of the cloud of gas and dust that formed our solar system and, as such, are
reminders of the origin of Earth and the other planets. In addition, the decay
of radioactive material provides much of the energy that heats Earth’s
core.
The most common naturally occurring
radioactive isotopes are potassium-40 (see Potassium), thorium-232
(see Thorium), and uranium-238 (see Uranium). Atoms of these
isotopes last, on average, for billions of years before undergoing alpha or beta
decay. The steady decay of these isotopes and other, more stable atoms allows
scientists to determine the age of minerals in which these isotopes occur.
Scientists begin by estimating the amount of isotope that was present when the
mineral formed, then measure how much has decayed. Knowing the rate at which the
isotope decays, they can determine how much time has passed. This process, known
as radioactive dating (see Dating Methods), allows scientists to measure
the age of Earth. The currently accepted value for Earth’s age is about 4.5
billion years. Scientists have also examined rocks from the Moon and other
objects in the solar system and have found that they have similar ages.
| IV | FORCES ACTING INSIDE ATOMS |
In physics, a force is a push or pull on an
object. There are four fundamental forces, three of which—the electromagnetic
force, the strong force, and the weak force—are involved in keeping stable atoms
in one piece and determining how unstable atoms will decay. The electromagnetic
force keeps electrons attached to their atom. The strong force holds the protons
and neutrons together in the nucleus. The weak force governs how atoms decay
when they have excess protons or neutrons. The fourth fundamental force,
gravity, only becomes apparent with objects much larger than subatomic
particles.
| A | Electromagnetic Force |
The most familiar of the forces at work
inside the atom is the electromagnetic force. This is the same force that causes
people’s hair to stick to a brush or comb when they have a buildup of static
electricity. The electromagnetic force causes opposite electric charges to
attract each other. Because of this force, the negatively charged electrons in
an atom are attracted to the positively charged protons in the atom’s nucleus.
This force of attraction binds the electrons to the atom. The electromagnetic
force becomes stronger as the distance between charges becomes smaller. This
property usually causes oppositely charged particles to come as close to each
other as possible. For many years, scientists wondered why electrons didn’t just
spiral into the nucleus of an atom, getting as close as possible to the protons.
Physicists eventually learned that particles as small as electrons can behave
like waves, and this property keeps electrons at set distances from the atom’s
nucleus. The wavelike nature of electrons is discussed below in the Quantum Atom
section of this article.
The electromagnetic force also causes like
charges to repel each other. The negatively charged electrons repel one another
and tend to move far apart from each other, but the positively charged nucleus
exerts enough electromagnetic force to keep the electrons attached to the atom.
Protons in the nucleus also repel one other, but, as described below, the strong
force overcomes the electromagnetic force in the nucleus to hold the protons
together.
| B | Strong Force |
Protons and neutrons in the nuclei of
atoms are held together by the strong force. This force must overcome the
electromagnetic force of repulsion the protons in a nucleus exert on one
another. The strong force that occurs between protons alone, however, is not
enough to hold them together. Other particles that add to the strong force, but
not to the electromagnetic force, must be present to make a nucleus stable. The
particles that provide this additional force are neutrons. Neutrons add to the
strong force of attraction but have no electric charge and so do not increase
the electromagnetic repulsion.
| B1 | Range of the Strong Force |
The strong force only operates at very
short range—about 2 femtometers (abbreviated fm), or 2 × 10-15 m (8 ×
10-14 in). Physicists also use the word fermi (also abbreviated fm)
for this unit in honor of Italian-born American physicist Enrico Fermi. The
short-range property of the strong force makes it very different from the
electromagnetic and gravitational forces. These latter forces become weaker as
distance increases, but they continue to affect objects millions of light-years
away from each other. Conversely, the strong force has such limited range that
not even all protons and neutrons in the same nucleus feel each other’s strong
force. Because the diameter of even a small nucleus is about 5 to 6 fm, protons
and neutrons on opposite sides of a nucleus only feel the strong force from
their nearest neighbors.
The strong force differs from
electromagnetic and gravitational forces in another important way—the way it
changes with distance. Electromagnetic and gravitational forces of attraction
increase as particles move closer to one another, no matter how close the
particles get. This increase causes particles to move as close together as
possible. The strong force, on the other hand, remains roughly constant as
protons and neutrons move closer together than about 2 fm. If the particles are
forced much closer together, the attractive nuclear force suddenly turns
repulsive. This property causes nuclei to form with the same average
spacing—about 2 fm—between the protons and neutrons, no matter how many protons
and neutrons there are in the nucleus.
The unique nature of the strong force
determines the relative number of protons and neutrons in the nucleus. If a
nucleus has too many protons, the strong force cannot overcome the
electromagnetic repulsion of the protons. If the nucleus has too many neutrons,
the excess strong force tries to crowd the protons and neutrons too close
together. Most stable atomic nuclei fall between these extremes. Lighter nuclei,
such as carbon-12 and oxygen-16, are made up of 50 percent protons and 50
percent neutrons. More massive nuclei, such as bismuth-209, contain about 40
percent protons and 60 percent neutrons.
| B2 | Pions |
Particle physicists explain the behavior
of the strong force by introducing another type of particle, called a pion.
Protons and neutrons interact in the nucleus by exchanging pions. Exchanging
pions pulls protons and neutrons together. The process is similar to two people
having a game of catch with a heavy ball, but with each person attached to the
ball by a spring. As one person throws the ball to the other, the spring pulls
the thrower toward the ball. If the players exchange the ball rapidly enough,
the ball and springs become just a blur to an observer, and it appears as if the
two throwers are simply pulled toward one another. This is what occurs in the
nuclei of atoms. The protons and neutrons in the nucleus are the people, pions
act as the ball, and the strong force acts as the springs holding everything
together.
Pions in the nucleus exist only for the
briefest instant of time, no more than 1 × 10-23 seconds, but even
during their short existence they can provide the attraction that holds the
nucleus together. Pions can also exist as independent particles outside of the
nucleus of an atom. Scientists have created them by striking high-speed protons
against a target. Even though the free pions also live only for a short period
of time (about 1 × 10-8 seconds), scientists have been able study
their properties.
| C | Weak Force |
The weak force lives up to its name—it is
much weaker than the electromagnetic and strong forces. Like the strong force,
it only acts over a short distance, about .01 fm. Unlike these other forces,
however, the weak force affects all the particles in an atom. The
electromagnetic force only affects the electrons and protons, and the strong
force only affects the protons and neutrons. When a nucleus has too many protons
to hold together or so many neutrons that the strong force squeezes too tightly,
the weak force actually changes one type of particle into another. When an atom
undergoes one type of decay, for example, the weak force causes a neutron to
change into a proton, an electron, and an electron antineutrino. The total
electric charge and the total energy of the particles remain the same before and
after the change.
| V | THE QUANTUM ATOM |
Scientists of the early 20th century found
they could not explain the behavior of atoms using their current knowledge of
matter. They had to develop a new view of matter and energy to accurately
describe how atoms behaved. They called this theory quantum theory, or quantum
mechanics. Quantum theory describes matter as acting both as a particle and as a
wave. In the visible objects encountered in everyday life, the wavelike nature
of matter is too small to be apparent. Wavelike nature becomes important,
however, in microscopic particles such as electrons. As we have discussed,
electrons in atoms behave like waves. They exist as a fuzzy cloud of negative
charge around the nucleus, instead of as a particle located at a single
point.
| A | Wave Behavior |
In order to understand the quantum model of
the atom, we must know some basic facts about waves. Waves are vibrations that
repeat regularly over and over again. A familiar example of waves occurs when
one end of a rope is tied to a fixed object and someone moves the other end up
and down. This action creates waves that travel along the rope. The highest
point that the rope reaches is called the crest of the wave. The lowest point is
called the trough of the wave. Troughs and crests follow each other in a regular
sequence. The distance from one trough to the next trough, or from one crest to
the next crest, is called a wavelength. The number of wavelengths that pass a
certain point in a given amount of time is called the wave’s frequency.
In physics, the word wave usually
means the entire pattern, which may consist of many individual troughs and
crests. For example, when the person holding the loose end of the rope moves it
up and down very fast, many troughs and crests occupy the rope at once. A
physicist would use the word wave to describe the entire set of troughs
and crests on the rope.
When two waves meet each other, they merge
in a process called interference. Interference creates a new wave pattern. If
two waves with the same wavelength and frequency come together, the resulting
pattern depends on the relative position of the waves’ crests. If the crests and
troughs of the two waves coincide, the waves are said to be in phase. Waves in
phase with each other will merge to produce higher crests and lower troughs.
Physicists call this type of interference constructive interference.
Sometimes waves with the same wavelength
and frequency are out of phase, meaning they meet in such a way that their
respective crests and troughs do not coincide. In these cases the waves produce
destructive interference. If two identical waves are exactly half a wavelength
out of phase, the crests of one wave line up with the troughs of the other.
These waves cancel each other out completely, and no wave will appear. If two
waves meet that are not exactly in phase and not exactly one-half wavelength out
of phase, they will interfere constructively in some places and destructively in
others, producing a complicated new wave. See also Wave Motion.
| B | Electrons as Waves |
Electrons behave as both particles and
waves in atoms. This characteristic is called wave-particle duality.
Wave-particle duality actually affects all particles and collections of
particles, including protons, neutrons, and atoms themselves. But in terms of
the structure of the atom, the wavelike nature of the electron is the most
important.
As waves, electrons have wavelengths and
frequencies. The wavelength of an electron depends on the electron’s energy.
Since the energy of electrons is kinetic (energy related to motion), an
electron’s wavelength depends on how fast it is moving. The more energy an
electron has, the shorter its wavelength is. Electron waves can interfere with
each other, just as waves along a rope do.
Because of the electron’s wave-particle
duality, physicists cannot define an electron’s exact location in an atom. If
the electron were just a particle, measuring its location would be relatively
simple. As soon as physicists try to measure its location, however, the
electron’s wavelike nature becomes apparent, and they cannot pinpoint an exact
location. Instead, physicists calculate the probability that the electron is
located in a certain place. Adding up all these probabilities, physicists can
produce a picture of the electron that resembles a fuzzy cloud around the
nucleus. The densest part of this cloud represents the place where the electron
is most likely to be located.
| C | Electron Orbitals and Shells |
Physicists call the region of space an
electron occupies in an atom the electron’s orbital. Similar orbitals constitute
groups called shells. The electrons in the orbitals of a particular shell have
similar levels of energy. This energy is in the form of both kinetic energy and
potential energy. Lower shells are close to the nucleus and higher shells are
farther from the nucleus. Electrons occupying orbitals in higher shells
generally have more energy than electrons occupying orbitals in lower
shells.
| C1 | Differences Between Orbitals |
The wavelike nature of electrons sets
boundaries for their possible locations and determines what shape their orbital,
or cloud of probability, will form. Orbitals differ from each other in size,
angular momentum, and magnetic properties. In general, angular momentum is the
energy an object contains based on how fast the object is revolving, the
object’s mass, and the object’s distance from the axis around which it is
revolving. The angular momentum of a whirling ball tied to a string, for
example, would be greater if the ball was heavier, the string was longer, or the
whirling was faster. In atoms, the angular momentum of an electron orbital
depends on the size and shape of the orbital. Orbitals with the same size and
shape all have the same angular momentum. Some orbitals, however, can differ in
shape but still have the same angular momentum. The magnetic properties of an
orbital describe how it would behave in a magnetic field. Magnetic properties
also depend on the size and shape of the orbital, as well as on the orbital’s
orientation in space.
The orbitals in an atom must occur at
certain distances from the nucleus to create a stable atom. At these distances,
the orbitals allow the electron wave to complete one or more half-wavelengths
(y, 1, 1y, 2, 2y, and so on) as it travels
around the nucleus. The electron wave can then double back on itself and
constructively interfere with itself in a way that reinforces the wave. Any
other distance would cause the electron to interfere with its own wave in an
unpredictable and unstable way, creating an unstable atom.
| C2 | Principal and Secondary Quantum Numbers |
Physicists call the number of
half-wavelengths that an orbital allows the orbital’s principal quantum
number (abbreviated n). In general, this number determines the size of the
orbital. Larger orbitals allow more half-wavelengths and therefore have higher
principal quantum numbers. The orbital that allows one half-wavelength has a
principal quantum number of one. Only one orbital allows one half-wavelength.
More than one orbital can allow two or more half-wavelengths. These orbitals may
have the same principal quantum number, but they differ from each other in their
angular momentum and their magnetic properties. The orbitals that allow one
wavelength have a principal quantum number of 2 (n = 2), the orbitals that allow
one and a half wavelengths have a principal quantum number of 3 (n = 3), and so
on. The set of orbitals with the same principal quantum number make up a
shell.
Physicists use a second number to
describe the angular momentum of an orbital. This number is called the orbital’s
secondary quantum number, or its angular momentum quantum number (abbreviated
l). The number of possible values an orbital can have for its angular momentum
is one less than the number of half-wavelengths it allows. This means that an
orbital with a principal quantum number of n can have n-1 possible values for
its secondary quantum number.
Physicists customarily use letters to
indicate orbitals with certain secondary quantum numbers. In order of increasing
angular momentum, the orbitals with the six lowest secondary quantum numbers are
indicated by the letters s, p, d, f, g, and
h. The letter s corresponds to the secondary quantum number 0, the
letter p corresponds to the secondary quantum number 1, and so on. In
general, the angular momentum of an orbital depends on its shape. An s-orbital,
with a secondary quantum number of 0, is spherical. A p-orbital, with a
secondary quantum number of 1, resembles two hemispheres, facing one another.
The possible combinations of principal and secondary quantum numbers for the
first five shells are listed below.
| C3 | Subshells |
More than one orbital can allow the same
number of half-wavelengths and have the same angular momentum. Physicists call
orbitals in a shell that all have the same angular momentum a subshell. They
designate a subshell with the subshell’s principal and secondary quantum
numbers. For example, the 1s subshell is the group of orbitals in the first
shell with an angular momentum described by the letter s. The 2p subshell
is the group of orbitals in the second shell with an angular momentum described
by the letter p.
Orbitals within a subshell differ from
each other in their magnetic properties. The magnetic properties of an orbital
depend on its shape and orientation in space. For example, a p-orbital can have
three different orientations in space: one situated up and down, one from side
to side, and a third from front to back.
| C4 | Magnetic Quantum Number and Spin |
Physicists describe the magnetic
properties of an orbital with a third quantum number called the orbital’s
magnetic quantum number (abbreviated m). The magnetic quantum number determines
how orbitals with the same size and angular momentum are oriented in space. An
orbital’s magnetic quantum number can only have whole number values ranging from
the value of the orbital’s secondary quantum number down to the negative value
of the secondary quantum number. A p-orbital, for example, has a secondary
quantum number of 1 (l = 1), so the magnetic quantum number has three possible
values: +1, 0, and -1. This means the p-orbital has three possible orientations
in space. An s-orbital has a secondary quantum number of 0 (l = 0), so the
magnetic quantum number has only one possibility: 0. This orbital is a sphere,
and a sphere can only have one orientation in space. For a d-orbital, the
secondary quantum number is 2 (l = 2), so the magnetic quantum number has five
possible values: -2, -1, 0, +1, and +2. A d-orbital has four possible
orientations in space, as well as a fifth orbital that differs in shape from the
other four. Together, the principal, secondary, and magnetic quantum numbers
specify a particular orbital in an atom.
Electrons are a type of particle known as
a fermion. Austrian-American physicist Wolfgang Pauli discovered that no two
fermions can have the exact same quantum numbers. This principle is called the
Pauli exclusion principle, which states that two or more identical electrons
cannot occupy the same orbital in an atom. Scientists know, however, that each
orbital can hold two electrons. Electrons have another property, called spin,
that differentiates the two electrons in each orbital. An electron’s spin has
two possible values: +y
(called spin-up) or -y
(called spin-down). These two possible values mean that two electrons can occupy
the same orbital, as long as their spins are different. Physicists call spin the
fourth quantum number of an electron orbital (abbreviated ms). Spin,
in addition to the other three quantum numbers, uniquely describes a particular
electron’s orbital.
| C5 | Filling Orbitals |
When electrons collect around an atom’s
nucleus, they fill up orbitals in a definite pattern. They seek the first
available orbital that takes the least amount of energy to occupy. Generally, it
takes more energy to occupy orbitals with higher quantum numbers. It takes the
same energy to occupy all the orbitals in a subshell. The lowest energy orbital
is the one closest to the nucleus. It has a principal quantum number of 1, a
secondary quantum number of 0, and a magnetic quantum number of 0. The first two
electrons—with opposite spins—occupy this orbital.
If an atom has more than two electrons,
the electrons begin filling orbitals in the next subshell with one electron each
until all the orbitals in the subshell have one electron. The electrons that are
left then go back and fill each orbital in the subshell with a second electron
with opposite spin. They follow this order because it takes less energy to add
an electron to an empty orbital than to complete a pair of electrons in an
orbital. The electrons fill all the subshells in a shell, then go on to the next
shell. As the subshells and shells increase, the order of energy for orbitals
becomes more complicated. For example, it takes slightly less energy to occupy
the s-subshell in the fourth shell than it does to occupy the d-subshell in the
third shell. Electrons will therefore fill the orbitals in the 4s subshell
before they fill the orbitals in the 3d subshell, even though the 3d subshell is
in a lower shell.
| D | Atomic Properties |
The atom’s electron cloud, that is, the
arrangement of electrons around an atom, determines most of the atom’s physical
and chemical properties. Scientists can therefore predict how atoms will
interact with other atoms by studying their electron clouds. The electrons in
the outermost shell largely determine the chemical properties of an atom. If
this shell is full, meaning all the orbitals in the shell have two electrons,
then the atom is stable, and it won’t react readily with other atoms. If the
shell is not full, the atom will chemically react with other atoms, exchanging
or sharing electrons in order to fill its outer shell. Atoms bond with other
atoms to fill their outer shells because it requires less energy to exist in
this bonded state. Atoms always seek to exist in the lowest energy state
possible.
| D1 | Valence Shells |
Physicists call the outer shell of an
atom its valence shell. The valence shell determines the atom’s chemical
behavior, or how it reacts with other elements. The fullness of an atom’s
valence shell affects how the atom reacts with other atoms. Atoms with valence
shells that are completely full are not likely to interact with other atoms. Six
gaseous elements—helium, neon, argon, krypton, xenon, and radon—have full
valence shells. These six elements are often called the noble gases because they
do not normally form compounds with other elements. The noble gases are
chemically inert because their atoms are in a state of low energy. A full
valence shell, like that of atoms of noble gases, provides the lowest and most
stable energy for an atom.
Atoms that do not have a full valence
shell try to lower their energy by filling up their valence shell. They can do
this in several ways: Two atoms can share electrons to complete the valence
shell of both atoms, an atom can shed or take on electrons to create a full
valence shell, or a large number of atoms can share a common pool of electrons
to complete their valence shells.
| D2 | Covalent Bonds |
When two atoms share a pair of electrons,
they form a covalent bond. When atoms bond covalently, they form molecules. A
molecule can be made up of two or more atoms, all joined with covalent bonds.
Each atom can share its electrons with one or more other atoms. Some molecules
contain chains of thousands of covalently bonded atoms.
Carbon is an important example of an
element that readily forms covalent bonds. Carbon has a total of six electrons.
Two of the electrons fill up the first orbital, the 1s orbital, which is the
only orbital in the first shell. The rest of the electrons partially fill
carbon’s valence shell. Two fill up the next orbital, the 2s orbital, which
forms the 2s subshell. Carbon’s valence shell still has the 2p subshell,
containing three p-orbitals. The two remaining electrons each fill half of the
two orbitals in the 2p subshell. The carbon atom thus has two half-full orbitals
and one empty orbital in its valence shell. A carbon atom fills its valence
shell by sharing electrons with other atoms, creating covalent bonds. The carbon
atom can bond with other atoms through any of the three unfilled orbitals in its
valence shell. The three available orbitals in carbon’s valence shell enable
carbon to bond with other atoms in many different ways. This flexibility allows
carbon to form a great variety of molecules, which can have a similarly great
variety of geometrical shapes. This diversity of carbon-based molecules is
responsible for the importance of carbon in molecules that form the basis for
living things (see Organic Chemistry).
| D3 | Ionic Bonds |
Atoms can also lose or gain electrons to
complete their valence shell. An atom will tend to lose electrons if it has just
a few electrons in its valence shell. After losing the electrons, the next lower
shell, which is full, becomes its valence shell. An atom will tend to steal
electrons away from other atoms if it only needs a few more electrons to
complete the shell. Losing or gaining electrons gives an atom a net electric
charge because the number of electrons in the atom is no longer the same as the
number of protons. Atoms with net electric charge are called ions. Scientists
call atoms with a net positive electric charge cations (pronounced CAT-eye-uhns)
and atoms with a net negative electric charge anions (pronounced
AN-eye-uhns).
The oppositely charged cations and anions
are attracted to each other by electromagnetic force and form ionic bonds. When
these ions come together, they form crystals. A crystal is a solid material made
up of repeating patterns of atoms. Alternating positive and negative ions build
up into a solid lattice, or framework. Crystals are also called ionic compounds,
or salts.
The element sodium is an example of an
atom that has a single electron in its valence shell. It will easily lose this
electron and become a cation. Chlorine atoms are just one electron away from
completing their valence shell. They will tend to steal an electron away from
another atom, forming an anion. When sodium and chlorine atoms come together,
the sodium atoms readily give up their outer electron to the chlorine atoms. The
oppositely charged ions bond with each other to form the crystal known as sodium
chloride, or table salt. See also Chemical Reaction.
| D4 | Metallic Bonds |
Atoms can complete their valence shells
in a third way: by bonding together in such a way so that all the atoms in the
substance share each other’s outer electrons. This is the way metallic elements
bond and fill their valence shells. Metals form crystal lattice structures
similar to salts, but the outer electrons in their atoms do not belong to any
atom in particular. Instead, the outer electrons belong to all the atoms in the
crystal, and they are free to move throughout the crystal. This property makes
metals good conductors of electricity.
| D5 | The Periodic Table |
The organization of the periodic table
reflects the way elements fill their orbitals with electrons. Scientists first
developed this chart by grouping together elements that behave similarly in
order of increasing atomic number. Scientists eventually realized that the
chemical and physical behavior of elements was dependant on the electron clouds
of the atoms of each element. The periodic table does not have a simple
rectangular shape. Each column lists elements that share chemical properties,
properties that depend on the arrangement of electrons in the orbitals of atoms.
These elements have the same number of electrons in their valence shells.
Different numbers of elements have similar valence shells, so the columns of the
periodic table differ in height. The noble gases are all located in the
rightmost column of the periodic table, labeled column 18 in Encarta’s periodic
table. The noble gases all have full valence shells and are extremely stable.
The column labeled 11 holds the elements copper, silver, and gold. These
elements are metals that have partially filled valence shells and conduct
electricity well.
| E | Electron Energy Levels |
Each electron in an atom has a particular
energy. This energy depends on the electron’s speed, the presence of other
electrons, the electron’s distance from the nucleus, and the positive charge of
the nucleus. For atoms with more than one electron, calculating the energy of
each electron becomes too complicated to be practical. However, the order and
relative energies of electrons follows the order of the electron orbitals, as
discussed in the Electron Orbital and Shell section of this article. Physicists
call the energy an electron has in a particular orbital the energy state
of the electron. For example, the 1s orbital holds the two electrons with the
lowest possible energies in an atom. These electrons are in the lowest energy
state of any electrons in the atom.
When an atom gains or loses energy, it does
so by adding energy to, or removing energy from, its electrons. This change in
energy causes the electrons to move from one orbital, or allowed energy state,
to another. Under ordinary conditions, all electrons in an atom are in their
lowest possible energy states, given that only two electrons can occupy each
orbital. Atoms gain energy by absorbing it from light or from a collision with
another particle, or they gain it by entering an electric or magnetic field.
When an atom absorbs energy, one or more of its electrons moves to a higher, or
more energetic, orbital. Usually atoms can only hold energy for a very short
amount of time—typically 1 × 10-12 seconds or less. When electrons
drop back down to their original energy states, they release their extra energy
in the form of a photon (a packet of radiation). Sometimes this radiation
is in the form of visible light. The light emitted by a fluorescent lamp is an
example of this process.
The outer electrons in an atom are easier
to move to higher orbitals than the electrons in lower orbitals. The inner
electrons require more energy to move because they are closer to the nucleus and
therefore experience a stronger electromagnetic pull toward the nucleus than the
outer electrons. When an inner electron absorbs energy and then falls back down,
the photon it emits has more energy than the photon an outer electron would
emit. The emitted energy relates directly to the wavelength of the photon.
Photons with more energy are made of radiation with a shorter wavelength. When
inner electrons drop down, they emit high-energy radiation, in the range of an X
ray. X rays have much shorter wavelengths than visible light. When outer
electrons drop down, they emit light with longer wavelengths, in the range of
visible light.
| VI | STUDYING ATOMS |
Physicists and chemists first learned about
the properties of atoms indirectly, by studying the way that atoms join together
in molecules or how atoms and molecules make up solids, liquids, and gases.
Modern devices such as electron microscopes, particle traps, spectroscopes, and
particle accelerators allow scientists to perform experiments on small groups of
atoms and even on individual atoms. Scientists use these experiments to study
the properties of atoms more directly.
| A | Electron Microscopes |
One of the most direct ways to study an
object is to take its photograph. Scientists take photographs of atoms by using
an electron microscope. An electron microscope imitates a normal camera, but it
uses electrons instead of visible light to form an image. In photography, light
reflects off of an object and is recorded on film or some other kind of
detector. Taking a photograph of an atom with light is difficult because atoms
are so tiny. Light, like all waves, tends to diffract, or bend around objects in
its path (see Diffraction). In order to take a sharp photograph of any
object, the wavelength of the light that bounces off the object must be much
smaller than the size of the object. If the object is about the same size as or
smaller than the light’s wavelength, the light will bend around the object and
produce a fuzzy image.
Atoms are so small that even the shortest
wavelengths of visible light will diffract around them. Therefore, capturing
photographic images of atoms requires the use of waves that are shorter than
those of visible light. X rays are a type of electromagnetic radiation like
visible light, but they have very short wavelengths—much too short to be visible
to human eyes. X-ray wavelengths are small enough to prevent the waves from
diffracting around atoms. X rays, however, have so much energy that when they
bounce off an atom, they knock electrons away from the atom. Scientists,
therefore, cannot use X rays to take a picture of an atom without changing the
atom. They must use a different method to get an accurate picture.
Electron microscopes provide scientists
with an alternate method. Scientists shine electrons, instead of light, on an
atom. As discussed in the Electrons as Waves section of this article, electrons
have wavelike properties, so they can behave like light waves. The simplest type
of electron microscope focuses the electrons reflected off of an object and
translates the pattern formed by the reflected electrons into a visible display.
Scientists have used this technique to create images of tiny insects and even
individual living cells, but they have not been able to use it to make a clear
image of objects smaller than about 10 nanometers (abbreviated nm), or 1 ×
10-8 m (4 × 10-7 in).
To get to the level of individual atoms,
scientists must use a more powerful type of electron microscope called a
scanning tunneling microscope (STM). An STM uses a tiny probe, the tip of which
can be as small as a single atom, to scan an object. An STM takes advantage of
another wavelike property of electrons called tunneling. Tunneling allows
electrons emitted from the probe of the microscope to penetrate, or tunnel into,
the surface of the object being examined. The rate at which the electrons tunnel
from the probe to the surface is related to the distance between the probe and
the surface. These moving electrons generate a tiny electric current that the
STM measures. The STM constantly adjusts the height of the probe to keep the
current constant. By tracking how the height of the probe changes as the probe
moves over the surface, scientists can get a detailed map of the surface. The
map can be so detailed that individual atoms on the surface are visible.
| B | Particle Traps |
Studying single atoms or small samples of
atoms can help scientists understand atomic structure. However, all atoms, even
atoms that are part of a solid material, are constantly in motion. This constant
motion makes them difficult to examine. To study single atoms, scientists must
slow the atoms down and confine them to one place. Scientists can slow and trap
atoms using devices called particle traps.
Slowing down atoms is actually the same as
cooling them. This is because an atom’s rate of motion is directly related to
its temperature. Atoms that are moving very quickly cause a substance to have a
high temperature. Atoms moving more slowly create a lower temperature.
Scientists therefore build traps that cool atoms down to a very low
temperature.
Several different types of particle traps
exist. Some traps are designed to slow down ions, while others are designed to
slow electrically neutral atoms. Traps for ions often use electric and magnetic
fields to influence the movement of the particle, confining it in a small space
or slowing it down. Traps for neutral atoms often use lasers, beams of light in
which the light waves are uniform and consistent. Light has no mass, but it
moves so quickly that it does have momentum. This property allows the light to
affect other particles, or “bump” into them. When laser light collides with
atoms, the momentum of the light forces the atoms to change speed and
direction.
Scientists use trapped and cooled atoms
for a variety of experiments, including those that precisely measure the
properties of individual atoms and those in which scientists construct extremely
accurate atomic clocks. Atomic clocks keep track of time by counting waves of
radiation emitted by atoms in traps inside the clock. Because the traps hold the
atoms at low temperatures, the mechanisms inside the clock can exercise more
control over the atom, reducing the possibility of error. Scientists can also
use isolated atoms to measure the force of gravity in an area with extreme
accuracy. These measurements are useful in oil exploration, among other things.
A deposit of oil or other substance beneath Earth’s surface has a different
density than the material surrounding it. The strength of the pull of gravity in
an area depends on the density of material in the area, so these changes in
density produce changes in the local strength of gravity. Advances in the
manipulation of atoms have also raised the possibility of using atoms to etch
electronic circuits. This would help make the circuits smaller and thereby allow
more circuits to fit in a tinier area.
In 1995 American physicists used particle
traps to cool a sample of rubidium atoms to a temperature near absolute zero
(-273°C, or –459°F). Absolute zero is the temperature at which all motion stops.
When the scientists cooled the rubidium atoms to such a low temperature, the
atoms slowed almost to a stop. The scientists knew that the momentum of the
atoms, which is related to their speed, was close to zero. At this point, a
special rule of quantum physics, called the uncertainty principle, greatly
affected the positions of the atoms. This rule states that the momentum and
position of a particle both cannot have precise values at the same time. The
scientists had a fairly precise value for the atom’s momentum (nearly zero), so
the positions of the atoms became very imprecise. The position of each atom
could be described as a large, fuzzy cloud of probability. The atoms were very
close together in the trap, so the probability clouds of many atoms overlapped
one another. It was impossible for the scientists to tell where one atom ended
and another began. In effect, the atoms formed one huge particle. This new state
of matter is called a Bose-Einstein condensate.
| C | Spectroscopes |
Spectroscopy is the study of the
radiation, or energy, that atoms, ions, molecules, and atomic nuclei emit. This
emitted energy is usually in the form of electromagnetic radiation—vibrating
electric and magnetic waves. Electromagnetic waves can have a variety of
wavelengths, including those of visible light. X rays, ultraviolet radiation,
and infrared radiation are also forms of electromagnetic radiation. Scientists
use spectroscopes to measure this emitted radiation.
| C1 | Characteristic Radiation of Atoms |
Atoms emit radiation when their
electrons lose energy and drop down to lower orbitals, or energy states, as
described in the Electron Energy Levels section above. The difference in energy
between the orbitals determines the wavelength of the emitted radiation. This
radiation can be in the form of visible light for outer electrons, or it can be
radiation of shorter wavelengths, such as X-ray radiation, for inner electrons.
Because the energies of the orbitals are strictly defined and differ from
element to element, atoms of a particular element can only emit certain
wavelengths of radiation. By studying the wavelengths of radiation emitted by a
substance, scientists can identify the element or elements comprising the
substance. For example, the outer electrons in a sodium atom emit a
characteristic yellow light when they return to lower orbitals. This is why
street lamps that use sodium vapor have a yellowish glow (See also
Sodium-Vapor Lamp).
Chemists often use a procedure called a
flame test to identify elements. In a flame test, the chemist burns a sample of
the element. The heat excites the outer electrons in the element’s atoms, making
the electrons jump to higher energy orbitals. When the electrons drop back down
to their original orbitals, they emit light characteristic of that element. This
light colors the flame and allows the chemist to identify the element.
The inner electrons of atoms also emit
radiation that can help scientists identify elements. The energy it takes to
boost an inner electron to a higher orbital is directly related to the positive
charge of the nucleus and the pull this charge exerts on the electron. When the
electron drops back to its original level, it emits the same amount of energy it
absorbed, so the emitted energy is also related to the nucleus’s charge. The
charge on the nucleus is equal to the atom’s atomic number.
Scientists measure the energy of the
emitted radiation by measuring the radiation’s wavelength. The radiation’s
energy is directly related to its wavelength, which usually resembles that of an
X ray for the inner electrons. By measuring the wavelength of the radiation that
an atom’s inner electron emits, scientists can identify the atom by its atomic
number. Scientists used this method in the 1910s to identify the atomic number
of the elements and to place the elements in their correct order in the periodic
table. The method is still used today to identify particularly heavy elements
(those with atomic numbers greater than 100) that are produced a few atoms at a
time in large accelerators (see Transuranium Elements).
| C2 | Radiation Released by Radioactivity |
Atomic nuclei emit radiation when they
undergo radioactive decay, as discussed in the Radioactivity section above.
Nuclei usually emit radiation with very short wavelengths (and therefore high
energy) when they decay. Often this radiation is in the form of gamma rays, a
form of electromagnetic radiation with wavelengths even shorter than X rays.
Once again, nuclei of different elements emit radiation of characteristic
wavelengths. Scientists can identify nuclei by measuring this radiation. This
method is especially useful in neutron activation analysis, a technique
scientists use for identifying the presence of tiny amounts of elements.
Scientists bombard samples that they wish to identify with neutrons. Some of the
neutrons join the nuclei, making them radioactive. When the nuclei decay, they
emit radiation that allows the scientists to identify the substance.
Environmental scientists use neutron activation analysis in studying air and
water pollution. Forensic scientists, who study evidence related to crimes, use
this technique to identify gunshot residue and traces of poisons.
| D | Particle Accelerators |
Particle accelerators are devices that
increase the speed of a beam of elementary particles such as protons and
electrons. Scientists use the accelerated beam to study collisions between
particles. The beam can collide with a target of stationary particles, or it can
collide with another accelerated beam of particles moving in the opposite
direction. If physicists use the nucleus of an atom as the target, the particles
and radiation produced in the collision can help them learn about the nucleus.
The faster the particles move, the higher the energy they contain. If collisions
occur at very high energy, it is possible to create particles never before
detected. In certain circumstances, energy can be converted to matter, resulting
in heavier particles after the collision.
Cyclotrons and linear accelerators are two
of the most important kinds of particle accelerators. In a cyclotron, a magnetic
field holds a beam of charged particles in a circular path. An electric field
interacts with the particles’ electric charge to give them a boost of energy and
speed each time the beam goes around. In linear accelerators, charged particles
move in a straight line. They receive many small boosts of energy from electric
fields as they move through the accelerator.
Bombarding nuclei with beams of neutrons
forces the nuclei to absorb some of the neutrons and become unstable. The
unstable nuclei then decay radioactively. The way atoms decay tells scientists
about the original structure of the atom. Scientists can also deduce the size
and shape of nuclei from the way particles scatter from nuclei when they
collide. Another use of particle accelerators is to create new and exotic
isotopes, including atoms of elements with very high atomic numbers that are not
found in nature.
At higher energy levels, using particles
moving at much higher speeds, scientists can use accelerators to look inside
protons and neutrons to examine their internal structure. At these energy
levels, accelerators can produce new types of particles. Some of these particles
are similar to protons or neutrons but have larger masses and are very unstable.
Others have a structure similar to the pion, the particle that is exchanged
between the proton and neutron as part of the strong force that binds the
nucleus together. By creating new particles and studying their properties,
physicists have been able to deduce their common internal structure and to
classify them using the theory of quarks. High-energy collisions between one
particle and another often produce hundreds of particles. Experimenters have the
challenging task of identifying and measuring all of these particles, some of
which exist for only the tiniest fraction of a second.
| VII | HISTORY OF ATOMIC THEORY |
Beginning with Democritus, who lived during
the late 5th and early 4th centuries bc, Greek philosophers developed a
theory of matter that was not based on experimental evidence, but on their
attempts to understand the universe in philosophical terms. According to this
theory, all matter was composed of tiny, indivisible particles called atoms
(from the Greek word atomos, meaning “indivisible”). If a sample of a
pure element was divided into smaller and smaller parts, eventually a point
would be reached at which no further cutting would be possible—this was the atom
of that element, the smallest possible bit of that element.
According to the ancient Greeks, atoms were
all made of the same basic material, but atoms of different elements had
different sizes and shapes. The sizes, shapes, and arrangements of a material’s
atoms determined the material’s properties. For example, the atoms of a fluid
were smooth so that they could easily slide over one another, while the atoms of
a solid were rough and jagged so that they could attach to one another. Other
than the atoms, matter was empty space. Atoms and empty space were believed to
be the ultimate reality.
Although the notion of atoms as tiny bits
of elemental matter is consistent with modern atomic theory, the researchers of
prior eras did not understand the nature of atoms or their interactions in
materials. For centuries scientists did not have the methods or technology to
test their theories about the basic structure of matter, so people accepted the
ancient Greek view.
| A | The Birth of the Modern Atomic Theory |
The work of British chemist John Dalton
at the beginning of the 19th century revealed some of the first clues about the
true nature of atoms. Dalton studied how quantities of different elements, such
as hydrogen and oxygen, could combine to make other substances, such as water.
In his book A New System of Chemical Philosophy (1808), Dalton made two
assertions about atoms: (1) atoms of each element are all identical to one
another but different from the atoms of all other elements, and (2) atoms of
different elements can combine to form more complex substances.
Dalton’s idea that different elements had
different atoms was unlike the Greek idea of atoms. The characteristics of
Dalton’s atoms determined the chemical and physical properties of a substance,
no matter what the substance’s form. For example, carbon atoms can form both
hard diamonds and soft graphite. In the Greek theory of atoms, diamond atoms
would be very different from graphite atoms. In Dalton’s theory, diamond atoms
would be very similar to graphite atoms because both substances are composed of
the same chemical element.
While developing his theory of atoms,
Dalton observed that two elements can combine in more than one way. For example,
modern scientists know that carbon monoxide (CO) and carbon dioxide
(CO2) are both compounds of carbon and oxygen. According to Dalton’s
experiments, the quantities of an element needed to form different compounds are
always whole-number multiples of one another. For example, two times as much
oxygen is needed to form a liter of CO2 than is needed to form a
liter of CO. Dalton correctly concluded that compounds were created when atoms
of pure elements joined together in fixed proportions to form units that
scientists today call molecules.
| A1 | States of Matter |
Scientists in the early 19th century
struggled in another area of atomic theory. They tried to understand how atoms
of a single element could exist in solid, liquid, and gaseous forms. Scientists
correctly proposed that atoms in a solid attract each other with enough force to
hold the solid together, but they did not understand why the atoms of liquids
and gases did not attract each other as strongly. Some scientists theorized that
the forces between atoms were attractive at short distances (such as when the
atoms were packed very close together to form a solid) and repulsive at larger
distances (such as in a gas, where the atoms are on the average relatively far
apart).
Scientists had difficulty solving the
problem of states of matter because they did not adequately understand the
nature of heat. Today scientists recognize that heat is a form of energy, and
that different amounts of this energy in a substance lead to different states of
matter. In the 19th century, however, people believed that heat was a material
substance, called caloric, that could be transferred from one object to another.
This explanation of heat was called the caloric theory. Dalton used the caloric
theory to propose that each molecule of a gas is surrounded by caloric, which
exerts a repulsive force on other molecules. According to Dalton’s theory, as a
gas is heated, more caloric is added to the gas, which increases the repulsive
force between the molecules. More caloric would also cause the gas to exert a
greater pressure on the walls of its container, in accordance with scientists’
experiments.
This early explanation of heat and
states of matter broke down when experiments in the middle of the 19th century
showed that heat could change into energy of motion. The laws of physics state
that the amount of energy in a system cannot increase, so scientists had to
accept that heat must be energy, not a substance. This revelation required a new
theory of how atoms in different states of matter behave.
| A2 | Behavior of Gases |
In the early 19th century Italian
chemist Amedeo Avogadro made an important advance in the understanding of how
atoms and molecules in a gas behave. Avogadro began his work from a theory
developed by Dalton. Dalton’s theory proposed that a gaseous compound, formed by
combining equal numbers of atoms of two elements, should have the same number of
molecules as the atoms in one of the original elements. For example, ten atoms
of the element hydrogen (H) combine with ten atoms of chlorine (Cl) to form ten
gaseous hydrogen chloride (HCl) molecules.
In 1811 Avogadro developed a law of
physics that seemed to contradict Dalton’s theory. Avogadro’s law states that
equal volumes of different gases contain the same number of particles (atoms or
molecules) if both gases are at the same temperature and pressure. In Dalton’s
experiment, the volume of the original vessels containing the hydrogen or
chlorine gases was the same as the volume of the vessel containing the hydrogen
chloride gas. The pressures of the original hydrogen and chlorine gases were
equal, but the pressure of the hydrochloric gas was twice as great as either of
the original gases. According to Avogadro’s law, this doubled pressure would
mean that there were twice as many hydrogen chloride gas particles than there
had been chlorine particles prior to their combination.
To reconcile the results of Dalton’s
experiment with his new rule, Avogadro was forced to conclude that the original
vessels of hydrogen or chlorine contained only half as many particles as Dalton
had thought. Dalton, however, knew the total weight of each gas in the vessels,
as well as the weight of an individual atom of each gas, so he knew the total
number of atoms of each gas that was present in the vessels. Avogadro reconciled
the fact that there were twice as many atoms as there were particles in the
vessels by proposing that gases such as hydrogen and chlorine are really made up
of molecules of hydrogen and chlorine, with two atoms in each molecule. Today
scientists write the chemical symbols for hydrogen and chlorine as H2
and Cl2, respectively, indicating that there are two atoms in each
molecule. One molecule of hydrogen and one molecule of chlorine combine to form
two molecules of hydrogen chlorine (H2 + Cl2 → 2HCl). The
sample of hydrogen chloride contains twice the number of particles as either the
hydrogen or chlorine because two molecules of hydrogen chloride form when a
molecule of hydrogen combines with a molecule of chlorine.
| B | Electrical Forces in Atoms |
The work of Dalton and Avogadro led to a
consistent view of the quantities of different gases that could be combined to
form compounds, but scientists still did not understand the nature of the forces
that attracted the atoms to one another in compounds and molecules. Scientists
suspected that electrical forces might have something to do with that
attraction, but they found it difficult to understand how electrical forces
could allow two identical, neutral hydrogen atoms to attract one another to form
a hydrogen molecule.
In the 1830s, British physicist Michael
Faraday took the first significant step toward appreciating the importance of
electrical forces in compounds. Faraday placed two electrodes connected to
opposite terminals of a battery into a solution of water containing a dissolved
compound. As the electric current flowed through the solution, Faraday observed
that one of the elements that comprised the dissolved compound became deposited
on one electrode while the other element became deposited on the other
electrode. The electric current provided by the electrodes undid the coupling of
atoms in the compound. Faraday also observed that the quantity of each element
deposited on an electrode was directly proportional to the total quantity of
current that flowed through the solution—the stronger the current, the more
material became deposited on the electrode. This discovery made it clear that
electrical forces must be in some way responsible for the joining of atoms in
compounds.
Despite these significant discoveries,
most scientists did not immediately accept that atoms as described by Dalton,
Faraday, and Avogadro were responsible for the chemical and physical behavior of
substances. Before the end of the 19th century, many scientists believed that
all chemical and physical properties could be determined by the rules of heat,
an understanding of atoms closer to that of the Greek philosophers. The
development of the science of thermodynamics (the scientific study of
heat) and the recognition that heat was a form of energy eliminated the role of
caloric in atomic theory and made atomic theory more acceptable. The new theory
of heat, called the kinetic theory, said that the atoms or molecules of a
substance move faster, or gain kinetic energy, as heat energy is added to the
substance. Nevertheless, a small but powerful group of scientists still did not
accept the existence of atoms—they regarded atoms as convenient mathematical
devices that explained the chemistry of compounds, not as real entities.
In 1905 French chemist Jean-Baptiste
Perrin performed the final experiments that helped prove the atomic theory of
matter. Perrin observed the irregular wiggling of pollen grains suspended in a
liquid (a phenomenon called Brownian motion) and correctly explained that the
wiggling was the result of atoms of the fluid colliding with the pollen grains.
This experiment showed that the idea that materials were composed of real atoms
in thermal motion was in fact correct.
As scientists began to accept atomic
theory, researchers turned their efforts to understanding the electrical
properties of the atom. Several scientists, most notably British scientist Sir
William Crookes, studied the effects of sending electric current through a gas.
The scientists placed a very small amount of gas in a sealed glass tube. The
tube had electrodes at either end. When an electric current was applied to the
gas, a stream of electrically charged particles flowed from one of the
electrodes. This electrode was called the cathode, and the particles were called
cathode rays.
At first scientists believed that the
rays were composed of charged atoms or molecules, but experiments showed that
the cathode rays could penetrate thin sheets of material, which would not be
possible for a particle as large as an atom or a molecule. British physicist Sir
Joseph John Thomson measured the velocity of the cathode rays and showed that
they were much too fast to be atoms or molecules. No known force could
accelerate a particle as heavy as an atom or a molecule to such a high speed.
Thomson also measured the ratio of the charge of a cathode ray to the mass of
the cathode ray. The value he measured was about 1,000 times larger than any
previous measurement associated with charged atoms or molecules, indicating that
within cathode rays particularly tiny masses carried relatively large amounts of
charge. Thomson studied different gases and always found the same value for the
charge-to-mass ratio. He concluded that he was observing a new type of particle,
which carried a negative electric charge but was about a thousand times less
massive than the lightest known atom. He also concluded that these particles
were constituents of all atoms. Today scientists know these particles as
electrons, and Thomson is credited with their discovery.
| C | Rutherford’s Nuclear Atom |
Scientists realized that if all atoms
contain electrons but are electrically neutral, atoms must also contain an equal
quantity of positive charge to balance the electrons’ negative charge.
Furthermore, if electrons are indeed much less massive than even the lightest
atom, then this positive charge must account for most of the mass of the atom.
Thomson proposed a model by which this phenomenon could occur: He suggested that
the atom was a sphere of positive charge into which the negative electrons were
imbedded, like raisins in a loaf of raisin bread. In 1911 British scientist
Ernest Rutherford set out to test Thomson’s proposal by firing a beam of charged
particles at atoms.
Rutherford chose alpha particles for his
beam. Alpha particles are heavy particles with twice the positive charge of a
proton. Alpha particles are now known to be the nuclei of helium atoms, which
contain two protons and two neutrons. If Thomson’s model of the atom was
correct, Rutherford theorized that the electric charge and the mass of the atoms
would be too spread out to significantly deflect the alpha particles. Rutherford
was quite surprised to observe something very different. Most of the alpha
particles did indeed change their paths by a small angle, and occasionally an
alpha particle bounced back in the opposite direction. The alpha particles that
bounced back must have struck something at least as heavy as themselves. This
led Rutherford to propose a very different model for the atom. Instead of
supposing that the positive charge and mass were spread throughout the volume of
the atom, he theorized that it was concentrated in the center of the atom.
Rutherford called this concentrated region of electric charge the nucleus of the
atom.
In the span of 100 years, from Dalton to
Rutherford, the basic ideas of atomic structure evolved from very primitive
concepts of how atoms combined with one another to an understanding of the
constituents of atoms—a positively charged nucleus surrounded by negatively
charged electrons. The interactions between the nucleus and the electrons still
required study. It was natural for physicists to model the atom, in which tiny
electrons orbit a much more massive nucleus, after a familiar structure such as
the solar system, in which planets orbit around a much more massive Sun.
Rutherford’s model of the atom did indeed resemble a tiny solar system. The only
difference between early models of the nuclear atom and the solar system was
that atoms were held together by electromagnetic force, while gravitational
force holds together the solar system.
| D | The Bohr Model |
Danish physicist Niels Bohr used new
knowledge about the radiation emitted from atoms to develop a model of the atom
significantly different from Rutherford’s model. Scientists of the 19th century
discovered that when an electrical discharge passes through a small quantity of
a gas in a glass tube, the atoms in the gas emit light. This radiation occurs
only at certain discrete wavelengths, and different elements and compounds emit
different wavelengths. Bohr, working in Rutherford’s laboratory, set out to
understand the emission of radiation at these wavelengths based on the nuclear
model of the atom.
Using Rutherford’s model of the atom as
a miniature solar system, Bohr developed a theory by which he could predict the
same wavelengths scientists had measured radiating from atoms with a single
electron. However, when conceiving this theory, Bohr was forced to make some
startling conclusions. He concluded that because atoms emit light only at
discrete wavelengths, electrons could only orbit at certain designated radii,
and light could be emitted only when an electron jumped from one of these
designated orbits to another. Both of these conclusions were in disagreement
with classical physics, which imposed no strict rules on the size of orbits. To
make his theory work, Bohr had to propose special rules that violated the rules
of classical physics. He concluded that, on the atomic scale, certain preferred
states of motion were especially stable. In these states of motion an orbiting
electron (contrary to the laws of electromagnetism) would not radiate
energy.
At the same time that Bohr and
Rutherford were developing the nuclear model of the atom, other experiments
indicated similar failures of classical physics. These experiments included the
emission of radiation from hot, glowing objects (called thermal radiation) and
the release of electrons from metal surfaces illuminated with ultraviolet light
(the photoelectric effect). Classical physics could not account for these
observations, and scientists began to realize that they needed to take a new
approach. They called this new approach quantum mechanics (see Quantum
Theory), and they developed a mathematical basis for it in the 1920s. The laws
of classical physics work perfectly well on the scale of everyday objects, but
on the tiny atomic scale, the laws of quantum mechanics apply.
| E | Quantum Theory of Atoms |
The quantum mechanical view of atomic
structure maintains some of Rutherford and Bohr’s ideas. The nucleus is still at
the center of the atom and provides the electrical attraction that binds the
electrons to the atom. Contrary to Bohr’s theory, however, the electrons do not
circulate in definite planet-like orbits. The quantum-mechanical approach
acknowledges the wavelike character of electrons and provides the framework for
viewing the electrons as fuzzy clouds of negative charge. Electrons still have
assigned states of motion, but these states of motion do not correspond to fixed
orbits. Instead, they tell us something about the geometry of the electron
cloud—its size and shape and whether it is spherical or bunched in lobes like a
figure eight. Physicists called these states of motion orbitals. Quantum
mechanics also provides the mathematical basis for understanding how atoms that
join together in molecules share electrons. Nearly 100 years after Faraday’s
pioneering experiments, the quantum theory confirmed that it is indeed
electrical forces that are responsible for the structure of molecules.
Two of the rules of quantum theory that
are most important to explaining the atom are the idea of wave-particle duality
and the exclusion principle. French physicist Louis de Broglie first suggested
that particles could be described as waves in 1924. In the same decade, Austrian
physicist Erwin Schrödinger and German physicist Werner Heisenberg expanded de
Broglie’s ideas into formal, mathematical descriptions of quantum mechanics. The
exclusion principle was developed by Austrian-born American physicist Wolfgang
Pauli in 1925. The Pauli exclusion principle states that no two electrons in an
atom can have exactly the same characteristics.
The combination of wave-particle duality
and the Pauli exclusion principle sets up the rules for filling electron
orbitals in atoms. The way electrons fill up orbitals determines the number of
electrons that end up in the atom’s valence shell. This in turn determines an
atom’s chemical and physical properties, such as how it reacts with other atoms
and how well it conducts electricity. These rules explain why atoms with similar
numbers of electrons can have very different properties, and why chemical
properties reappear again and again in a regular pattern among the elements.
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