It has been said that during the 20th century, man harnessed the "power of the
atom." We made
atomic bombs and generated
nuclear power. We even split the atom into smaller pieces called
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
- atom - smallest piece of an element that keeps its chemical
- compound - substance that can be broken into elements by chemical
- 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
- 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
- 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
"The Extraordinary Chemistry of Ordinary Things, 3rd Edition" by Carl H.
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.
- 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
- 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?
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:
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.
- 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
How Much Do Atoms Weigh?
The ability to weigh atoms came about by an observation from an Italian chemist
named Amadeo Avogadro. Avogadro 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:
Avogadro 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.
- One liter of nitrogen combined with three liters of hydrogen to form ammonia
- One liter of hydrogen combined with one liter of chlorine to make hydrogen
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
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:
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.
- There were some gaps in his "periodic table."
- When grouped by properties, most elements had increasing atomic masses, but
some were out of order.
In the next section, we will look at how we discovered the inside of the atom!
The Structure of the Atom: Early 20th Century
To know the structure of the atom, we must know the following:
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:
- What are the parts of the atom?
- How are these parts arranged?
- The parts of the atom:
- chemistry and electromagnetism ---> electron (first
- radioactivity ---> nucleus
- 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
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
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:
Thomson concluded the following:
- 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
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!
- 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
- 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.
Two other conclusions came from the discovery of the electron:
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
- 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.
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
radioactivity. Radioactivity was characterized by three types of emitted
How Radioactivity Works for details):
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:
- 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
- Gamma rays - neutrally charged and no mass (i.e., energy).
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!
- 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.
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
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 Sheds Light on the Atom: The Bohr
Branch of physics that deals with the motion of particles by their wave
properties at the atomic and subatomic level.
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
Photo courtesy NASA
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
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
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
- Atomic spectra - discussed above
- Periodic behavior of elements - elements with similar properties had similar
- 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
- Elements with full outer shells do not react.
- Other elements take or give up electrons to get a full outer shell.
Bohr's model was the predominant model until new discoveries in quantum
mechanics were made.
Electrons Can Behave as Waves: The Quantum Model of
Although the Bohr model adequately explained how atomic spectra worked, there
were several problems that bothered physicists and chemists:
Obviously, the Bohr model was missing something!
- 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?
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
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!
- To view an electron in its orbit, you must shine a wavelength of light on it
that is smaller than the electron's wavelength.
small wavelength of light has a high energy.
- The electron will absorb that energy.
- The absorbed energy will change the electron's position.
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
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.
- 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.
The orbitals had different shapes and maximum numbers at any level:
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 (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)
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:
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:
- one s orbital - two electrons
- one s orbital - two electrons and three p orbitals (two
- one s orbital - one electron
The modern periodic table of the elements (elements are ordered
based on atomic number rather than mass).
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
- Radioactivity - changes in the nucleus (i.e., decay) emit radioactive
Nuclear reactors - splitting the nucleus (fission)
Nuclear bombs - splitting the nucleus (fission) or forming a nucleus
- 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
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
- atom = 1 x 10-10 meters
- nucleus = 1 x 10-15 to 1 x 10
- neutron or proton = 1 x 10-15 meters
- electron - not known exactly, but thought to be on the order of 1 x 10-18
The STM works like this:
- 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,
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.
- 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
- 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 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
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.
Lots More Information!
Other Great Links
Scanning Tunneling Microscopy