How Atoms Work

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

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:

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:

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.

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.

At this time, chemists knew the atomic masses of elements and their chemical properties, and an astonishing phenomenon jumped out at them!

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.

In the next section, we will look at how we discovered the inside of the atom!

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:

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:

Thomson concluded the following:

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:

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.

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):

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:

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.

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!

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 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).

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.

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).

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.

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.

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!

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:

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:

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:

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:

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.

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:

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:

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:

The STM works like this:

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:

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.