It has been said that during the 20th century, man
harnessed the "power of the atom." We made atomic
bombs and generated electricity
by nuclear
power. We even split the atom into smaller pieces called
subatomic particles.
But what exactly is an atom? What is it made of? What does
it look like? The pursuit of the structure of the atom has
married many areas of chemistry and physics in perhaps one of
the greatest contributions of modern science!
In this edition of How Stuff
Works, we will follow this fascinating story of how
discoveries in various fields of science resulted in our
modern view of the atom. We will look at the consequences of
knowing the atom's structure and how this structure will lead
to new technologies.
What is an Atom? The Legacy of Ancient Times Through
the 19th Century
Important Terms
atom - smallest piece of an element that
keeps its chemical properties
compound - substance that can be broken
into elements by chemical reactions
electron - particle orbiting the nucleus of
an atom with a negative charge (mass = 9.10 x
10-28 grams)
element - substance that cannot be broken
down by chemical reactions
ion - electrically charged atom (i.e.,
excess positive or negative charge)
molecule - smallest piece of a compound
that keeps its chemical properties (made of two or
more atoms)
neutron - particle in the nucleus of an
atom with no charge (mass = 1.675 x 10-24 grams)
nucleus - dense, central core of an atom
(made of protons and neutrons)
proton - particle in the nucleus of an atom
with a positive charge (mass = 1.673 x 10-24 grams)
The modern view of an atom has come from many fields of
chemistry and physics. The idea of an atom came from ancient
Greek science/philosophy and from the results of 18th and 19th
century chemistry:
concept of the atom
measurements of atomic mass
repeating or periodic relationship between the elements
Concept of the Atom From
the ancient Greeks through today, we have pondered what
ordinary matter is made of. To understand the problem, here is
a simple demonstration from a book entitled "The
Extraordinary Chemistry of Ordinary Things, 3rd Edition"
by Carl H. Snyder:
Take a pile of paper clips (all of the same size and
color).
Divide the pile into two equal piles.
Divide each of the smaller piles into two equal piles.
Repeat step 3 until you are down to a pile containing
only one paper clip. That one paper clip still does the job
of a paper clip (i.e., hold loose papers together).
Now, take a pair of scissors and cut that one paper clip
in half. Can half of the paper clip do the same job as the
single paper clip?
If you do the same thing with any
element, you will reach an indivisible part that has the same
properties of the element, like the single paper clip. This
indivisible part is called an atom.
The idea of the atom was first devised by Democritus
in 530 B.C. In 1808, an English school teacher and scientist
named John Dalton proposed the modern atomic theory.
Modern atomic theory simply states the following:
Every element is made of atoms - piles of paper
clips.
All atoms of any element are the same - all the
paper clips in the pile are the same size and color.
Atoms of different elements are different (size,
properties) - like different sizes and colors of paper
clips.
Atoms of different elements can combine to form
compounds - you can link different sizes and colors of
paper clips together to make new structures.
In chemical reactions, atoms are not made, destroyed,
or changed - no new paper clips appear, no paper clips
get lost and no paper clips change from one size/color to
another.
In any compound, the numbers and kinds of atoms
remain the same - the total number and types of paper
clips that you start with are the same as when you finish.
Dalton's atomic theory formed the groundwork of
chemistry at that time. Dalton envisioned atoms as tiny
spheres with hooks on them. With these hooks, one atom could
combine with another in definite proportions. But some
elements could combine to make different compounds (e.g.,
hydrogen + oxygen could make water or hydrogen peroxide). So,
he could not say anything about the numbers of each atom in
the molecules of specific substances. Did water have one
oxygen with one hydrogen or one oxygen with two hydrogens?
This point was resolved when chemists figured out how to weigh
atoms.
How Much Do Atoms Weigh? The ability to
weigh atoms came about by an observation from an Italian
chemist named Amadeo Avagadro. Avagadro was working
with gases (nitrogen, hydrogen, oxygen, chlorine) and noticed
that when temperature and pressure was the same, these gases
combined in definite volume ratios. For example:
One liter of nitrogen combined with three liters of
hydrogen to form ammonia (NH3)
One liter of hydrogen combined with one liter of
chlorine to make hydrogen chloride (HCl)
Avagadro
said that at the same temperature and pressure, equal volumes
of the gases had the same number of molecules. So, by weighing
the volumes of gases, he could determine the ratios of atomic
masses. For example, a liter of oxygen weighed 16 times more
than a liter of hydrogen, so an atom of oxygen must be 16
times the mass of an atom of hydrogen. Work of this type
resulted in a relative mass scale for elements in which all of
the elements related to carbon (chosen as the standard -12).
Once the relative mass scale was made, later experiments were
able to relate the mass in grams of a substance to the number
of atoms and an atomic mass unit (amu) was found; 1 amu
or Dalton is equal to 1.66 x 10-24 grams.
At this time, chemists knew the atomic masses of elements
and their chemical properties, and an astonishing phenomenon
jumped out at them!
The Properties of Elements Showed a Repeating
Pattern At the time that atomic masses had been
discovered, a Russian chemist named Dimitri Mendeleev
was writing a textbook. For his book, he began to organize
elements in terms of their properties by placing the elements
and their newly discovered atomic masses in cards. He arranged
the elements by increasing atomic mass and noticed that
elements with similar properties appeared at regular intervals
or periods. Mendeleev's table had two problems:
There were some gaps in his "periodic table."
When grouped by properties, most elements had increasing
atomic masses, but some were out of order.
To
explain the gaps, Mendeleev said that the gaps were due to
undiscovered elements. In fact, his table successfully
predicted the existence of gallium and germanium, which were
discovered later. However, Mendeleev was never able to explain
why some of the elements were out of order or why the elements
should show this periodic behavior. This would have to wait
until we knew about the structure of the atom.
In the next section, we will look at how we discovered the
inside of the atom!
The Structure of the Atom: Early 20th Century
Science
To know the structure of the atom, we must know the
following:
What are the parts of the atom?
How are these parts arranged?
Near the end of
the 19th century, the atom was thought to be nothing more than
a tiny indivisible sphere (Dalton's view). However, a series
of discoveries in the fields of chemistry, electricity and
magnetism, radioactivity, and quantum mechanics in the late
19th and early 20th centuries changed all of that. Here is
what these fields contributed:
The parts of the atom:
chemistry and electromagnetism --->
electron (first subatomic particle)
radioactivity ---> nucleus
proton
neutron
How the atom is arranged - quantum mechanics puts it
all together:
atomic spectra ---> Bohr model of the
atom
wave-particle duality ---> Quantum
model of the atom
Chemistry and Electromagnetism: Discovering the
Electron In the late 19th century, chemists and
physicists were studying the relationship between electricity
and matter. They were placing high voltage electric currents
through glass tubes filled with low-pressure gas (mercury,
neon, xenon) much like neon
lights. Electric current was carried from one electrode
(cathode) through the gas to the other electrode
(anode) by a beam called cathode rays. In 1897,
a British physicist, J. J. Thomson did a series of
experiments with the following results:
He found that if the tube was placed within an electric
or magnetic field, then the cathode rays could be
deflected or moved (this is how the the cathode ray
tube (CRT) on your television works).
By applying an electric field alone, a magnetic field
alone, or both in combination, Thomson could measure the
ratio of the electric charge to the mass of the cathode
rays.
He found the same charge to mass ratio of cathode
rays was seen regardless of what material was inside the
tube or what the cathode was made of.
Thomson
concluded the following:
Cathode rays were made of tiny, negatively charged
particles, which he called electrons.
The electrons had to come from inside the atoms
of the gas or metal electrode.
Because the charge to mass ratio was the same for any
substance, the electrons were a basic part of all
atoms.
Because the charge to mass ratio of the electron was
very high, the electron must be very small.
Later, an American Physicist named Robert
Milikan measured the electrical charge of an electron.
With these two numbers (charge, charge to mass ratio),
physicists calculated the mass of the electron as 9.10 x
10-28 grams. For comparison, a
U.S. penny has a mass of 2.5 grams; so, 2.7 x 1027 or 2.7 billion billion billion
electrons would weigh as much as a penny!
Two other conclusions came from the discovery of the
electron:
Because the electron was negatively charged and atoms
are electrically neutral, there must be a positive charge
somewhere in the atom.
Because electrons are so much smaller than atoms,
there must be other, more massive particles in the
atom.
From these results, Thomson proposed a
model of the atom that was like a watermelon. The red part was
the positive charge and the seeds were the electrons.
Radioactivity: Discovering the Nucleus, the Proton and
the Neutron About the same time as Thomson's
experiments with cathode rays, physicists such as by Henri
Becquerel, Marie Curie, Pierre Curie, and Ernest Rutherford
were studying radioactivity.
Radioactivity was characterized by three types of emitted rays
(see How
Radioactivity Works for details):
Alpha particles - positively charged and massive.
Ernest Rutherford showed that these particles were the
nucleus of a helium atom.
Beta particles - negatively charged and light
(later shown to be electrons).
Gamma rays - neutrally charged and no mass (i.e.,
energy).
The experiment from radioactivity that
contributed most to our knowledge of the structure of the atom
was done by Rutherford and his colleagues. Rutherford
bombarded a thin foil of gold with a beam of alpha particles
and looked at the beams on a fluorescent screen, he noticed
the following:
Most of the particles went straight through the foil and
struck the screen.
Some (0.1 percent) were deflected or scattered in front
(at various angles) of the foil, while others were scattered
behind the foil.
Rutherford concluded that the
gold atoms were mostly empty space, which allowed most
of the alpha particles through. However, some small region
of the atom must have been dense enough to deflect or
scatter the alpha particle. He called this dense region the
nucleus (see The
Rutherford Experiment for an excellent Java simulation of
this important experiment!); the nucleus comprised most of the
mass of the atom. Later, when Rutherford bombarded nitrogen
with alpha particles, a positively charged particle that was
lighter than the alpha particle was emitted. He called these
particles protons and realized that they were a
fundamental particle in the nucleus. Protons have a mass of
1.673 x 10-24 grams, about
1,835 times larger than an electron!
However, protons could not be the only particle in the
nucleus because the number of protons in any given element
(determined by the electrical charge) was less than the weight
of the nucleus. Therefore, a third, neutrally charged particle
must exist! It was James Chadwick, a British physicist
and co-worker of Rutherford, who discovered the third
subatomic particle, the neutron. Chadwick bombarded
beryllium foil with alpha particles and noticed a neutral
radiation coming out. This neutral radiation could in turn
knock protons out of the nuclei of other substances. Chadwick
concluded that this radiation was a stream of neutrally
charged particles with about the same mass as a proton; the
neutron has a mass of 1.675 x 10-24 grams.
Rutherford's view of the atom
Now that the parts of the atom were known, how were they
arranged to make an atom? Rutherford's gold foil experiment
indicated that the nucleus was in the center of the atom and
that the atom was mostly empty space. So, he envisioned the
atom as the positively charged nucleus in the center with the
negatively charged electrons circling around it much like a
planet with moons. Although he had no evidence that the
electrons circled the nucleus, his model seemed reasonable;
however, it presented a problem. As the electrons moved in a
circle, they would lose energy and give off light. The loss of
energy would slow the electrons down. Like any satellite,
the slowing electrons would fall into the nucleus. In fact, it
was calculated that a Rutherford atom would last only
billionths of a second before collapsing! Something was
missing!
Quantum Mechanics: Putting It All
Together At the same time that discoveries were
being made with radioactivity, physicists and chemists were
studying how light
interacted with matter. These studies began the field of
quantum mechanics and helped solve the structure of the
atom.
Quantum Mechanics
Branch of physics that deals with the motion of
particles by their wave properties at the atomic and
subatomic level.
Quantum Mechanics Sheds Light on the Atom:
The Bohr Model Physicists and chemists studied the
nature of the light that
was given off when electric currents were passed through tubes
containing gaseous elements (hydrogen, helium, neon) and when
elements were heated (e.g., sodium, potassium, calcium, etc.)
in a flame. They passed the light from these sources through a
spectrometer (a device containing a narrow slit and a glass
prism).
Photo courtesy NASA White light passing through a
prism.
Photo courtesy NASA Continuous spectrum of white
light.
Now, when you pass sunlight through a prism, you get a
continuous spectrum of colors like a rainbow. However, when
light from these various sources was passed through a prism,
they found a dark background with discrete lines.
Photo courtesy NASA Hydrogen spectrum
Photo courtesy NASA Helium spectrum
Each element had a unique spectrum and the wavelength
of each line within a spectrum had a specific energy (see How Light
Works for details on the relationship between wavelength
and energy).
In 1913, a Danish physicist named Niels Bohr put
Rutherford's findings together with the observed spectra to
come up with a new model of the atom in a real leap of
intuition. Bohr suggested that the electrons orbiting an atom
could only exist at certain energy levels (i.e., distances)
from the nucleus, not at continuous levels as might be
expected from Rutherford's model. When atoms in the gas tubes
absorbed the energy from the electric current, the electrons
became excited and jumped from low energy levels (close to the
nucleus) to high energy levels (farther out from the nucleus).
The excited electrons would fall back to their original levels
and emit energy as light. Because there were specific
differences between the energy levels, only specific
wavelengths of light were seen in the spectrum (i.e., lines).
Bohr models of various atoms.
The major advantage of the Bohr model was that it worked.
It explained several things:
Atomic spectra - discussed above
Periodic behavior of elements - elements with similar
properties had similar atomic spectra.
Each electron orbit of the same size or energy
(shell) could only hold so many electrons.
First shell = two electrons
Second shell = eight electrons
Third shell and higher = eight electrons
When one shell was filled, electrons were found at
higher levels.
Chemical properties were based on the number of
electrons in the outermost shell.
Elements with full outer shells do not react.
Other elements take or give up electrons to get a
full outer shell.
As it turns
out, Bohr's model is also useful for explaining the behavior
of lasers
although these devices were not invented until the middle of
the 20th century.
Bohr's model was the predominant model until new
discoveries in quantum mechanics were made.
Electrons Can Behave as Waves: The Quantum Model of
the Atom Although the Bohr model adequately
explained how atomic spectra worked, there were several
problems that bothered physicists and chemists:
Why should electrons be confined to only specified
energy levels?
Why don't electrons give off light all of the time?
As electrons change direction in their circular orbits
(i.e., accelerate), they should give off light.
The Bohr model could explain the spectra of atoms with
one electron in the outer shell very well, but was not very
good for those with more than one electron in the outer
shell.
Why could only two electrons fit in the first shell and
why eight electrons in each shell after that? What was so
special about two and eight?
Obviously, the Bohr
model was missing something!
In 1924, a French physicist named Louis de Broglie
suggested that, like light,
electrons could act as both particles and waves (see De
Broglie Phase Wave Animation for details). De Broglie's
hypothesis was soon confirmed in experiments that showed
electron beams could be diffracted or bent as they passed
through a slit much like light could.
So, the waves produced by an electron confined in its orbit
about the nucleus sets up a standing
wave of specific wavelength, energy and frequency (i.e.,
Bohr's energy levels) much like a guitar string sets up a
standing wave when plucked.
Another question quickly followed de Broglie's idea. If an
electron traveled as a wave, could you locate the precise
position of the electron within the wave? A German physicist,
Werner Heisenberg, answered no in what he called the
uncertainty principle:
To view an electron in its orbit, you must shine a
wavelength of light on it that is smaller than the
electron's wavelength.
The absorbed energy will change the electron's position.
We can never know both the momentum and
position of an electron in an atom. Therefore, Heisenberg
said that we shouldn't view electrons as moving in
well-defined orbits about the nucleus!
With de Broglie's hypothesis and Heisenberg's uncertainty
principle in mind, an Austrian physicist named Erwin
Schrodinger derived a set of equations or wave
functions in 1926 for electrons. According to Schrodinger,
electrons confined in their orbits would set up standing waves
and you could describe only the probability of where an
electron could be. The distributions of these probabilities
formed regions of space about the nucleus were called
orbitals. Orbitals could be described as electron
density clouds (see Atomic
& Molecular Orbitals for a look at various orbitals).
The densest area of the cloud is where you have the greatest
probability of finding the electron and the least dense area
is where you have the lowest probability of finding the
electron.
The wave function of each electron can be described as a
set of three quantum numbers:
Principal number (n) - describes the energy
level.
Altazimuth number (l) - how fast the electron
moves in its orbit (angular momentum); like how fast a CD spins
(rpm). This is related to the shape of the orbital.
Magnetic (m) - its orientation in space.
It was later suggested that no two electrons could
be in the exact same state, so a fourth quantum number was
added. This number was related to the direction that the
electron spins while it is moving in its orbit (i.e.,
clockwise, counterclockwise). Only two electrons could share
the same orbital, one spinning clockwise and the other
spinning counterclockwise.
The orbitals had different shapes and maximum numbers at
any level:
s (sharp) - spherical (max = 1)
p (principal) - dumb-bell shaped (max = 3)
d (diffuse) - four-lobe-shaped (max = 5)
f (fundamental) - six-lobe shaped (max = 7)
The names of the orbitals came from names of atomic
spectral features before quantum mechanics was formally
invented. Each orbital can hold only two electrons. Also, the
orbitals have a specific order of filling, generally:
s
p
d
f
However, there is some overlap (any chemistry textbook has
the details).
The resulting model of the atom is called the quantum
model of the atom.
Quantum model of a sodium
atom.
Sodium has 11 electrons distributed in the following energy
levels:
one s orbital - two electrons
one s orbital - two electrons and three p
orbitals (two electrons each)
one s orbital - one electron
Right now,
the quantum model is the most realistic vision of the overall
structure of the atom. It explains much of what we know about
chemistry and physics. Here are some examples:
The modern periodic table of the elements
(elements are ordered based on atomic number rather than
mass).
Chemistry:
The
Periodic Table - the Table's pattern and arrangement
reflects the arrangement of electrons in the atom.
Elements have different atomic numbers - the number
of protons or electrons increases up the table as
electrons fill the shells.
Elements have different atomic masses - the number
of protons plus neutrons increases up the table.
Rows - elements of each row have the same number of
energy levels (shells).
Columns - elements have the same number of electrons
in the outermost energy level or shell (one to eight).
Chemical reactions - exchange of electrons
between various atoms (giving, taking, or sharing).
Exchange involves electrons in the outermost energy level
in attempts to fill the outermost shell (i.e., most stable
form of the atom).
Physics
Radioactivity - changes in the nucleus (i.e.,
decay) emit radioactive particles.
Nuclear
bombs - splitting the nucleus (fission) or forming a
nucleus (fusion)
Atomic spectra - caused by excited electrons
changing energy levels (absorption or emission of energy
in the form of light photons).
Can We See Atoms? Atoms are so small that we
cannot see them with our eyes (i.e., microscopic). To give you
a feel for some sizes, these are approximate diameters of
various atoms and particles:
atom = 1 x 10-10 meters
nucleus = 1 x 10-15 to 1
x 10 -14 meters
neutron or proton = 1 x 10-15 meters
electron - not known exactly, but thought to be on the
order of 1 x 10-18 meters
You cannot see an atom with a light
microscope. However, in 1981, a type of microscope called
a scanning tunneling microscope (STM) was developed.
The STM consists of the following:
A very small, sharp tip that conducts electricity
(probe)
A rapid piezoelectric
scanning device to which the tip is mounted
Electronic components to supply current to the tip,
control the scanner and accept the signals from the motion
sensor
Computer to control the system and do data analysis
(data collection, processing, display)
The STM works
like this:
A current is supplied to the tip (probe) while the
scanner rapidly moves the tip across the surface of a
conducting sample.
When the tip encounters an atom, the flow of electrons
between the atom and the tip changes.
The computer registers the change in current with the
x,y-position of the atom.
The scanner continues to position the tip over each
x,y-point on the sample surface, registering a current for
each point.
The computer collects the data and plots a map of
current over the surface that corresponds to a map of the
atomic positions.
The process is much like an old
phonograph where the needle is the tip and the grooves in the
vinyl record are the atoms. The STM tip moves over the atomic
contour of the surface, using tunneling current as a
sensitive detector of atomic position.
The STM and new variations of this microscope allow us to
see atoms. In addition, the STM can be used to manipulate
atoms as shown here:
Photo courtesy NIST Photo source: IBM's
Almaden Research Labs Atoms
can be positioned on a surface using the STM tip,
creating a custom pattern on the
surface.
Atoms can be moved and molded to make various devices such
as molecular motors (see How
Nanotechnology Will Work for details).
In summary, science in the 20th century has revealed the
structure of the atom. Scientists are now conducting
experiments to reveal details of the structure of the nucleus
and the forces that hold it together.
