That’s what I tell my students each year: matter is mostly empty space. The concept is hard to accept, especially if you have ever hit your head on something, but nonetheless true.
As an analogy, I have used a scale model for the hydrogen atom I picked up from somewhere now forgotten. Place an ant on the 50-yard line in a football stadium. The atom represents the proton. The electron is just outside the stadium, in the parking lot. In between electron and proton is space.
Now I have a new scale model to use. It’s online, to appeal to the digital generation. The creator has represented the electron as a one-pixel speck on the webpage, and far, far away the proton as a much larger ball. The distance and sizes of the particles are on the same scale.
The model, of course, is not to be taken literally. Atoms are not really miniature versions of solar systems. Electrons are strictly speaking not little specks traveling in neat little orbits around a spherical proton (or nucleus) that looks like Neptune. That conception (minus the Neptune part) dates back almost 100 years.
Time was, scientists did conceive of the atom as a solid ball. That picture changed in 1897, when J. J. Thomson identified the electon and proposed that it was part of the atom, rather than separate from it. He proposed a “plum pudding” model, in which the negatively charged electrons were embedded in a positively charged matrix.
In 1911, Ernest Rutherford tested this model with his famous gold foil experiment. Briefly, Rutherford bombarded an extremely thin sheet of gold foil with positively charged alpha particles, expecting that most would just pass through the “squishy” pudding.
Instead, some alpha particles veered off at substantial angles and a few came caroming back toward their radium source, as if they had hit something very hard. Rutherford said it was as surprising as if someone had shot a cannonball at a piece of tissue paper and had it bounce right back.
Rutherford proposed his own model, a planetary or nuclear one. The nucleus was positively charged and around it orbited the negatively charged electrons. This simple model is one that sticks in many minds as the way atoms “look.”
The Rutherford model had its problems, which Niels Bohr’s quantum model of the atom helped fix a year later. Bohr’s model was extremely successful in explaining the relationship between atoms and light emission and absorption, but there were two puzzling features.
For one, his model required the electron, a particle with mass, to disappear from one orbit (energy level) and instantaneously reappear in another. Not only did this behavior defy common sense, it violated a few laws of physics too.
Another issue was the requirement that the electron could only have discrete values for its total energy and angular momentum. Assuming this behavior yielded a successful model, but there was no physical explanation for that behavior.
In 1924, a French physics student proposed an explanation for both issues. Louis de Broglie hypothesized that a moving electron had wave-like properties, turning around earlier conclusions that light and energy had particle-like properties.
Waves have no intrinsic mass, so an electron wave can instantaneously disappear and reappear, solving the teleportation problem. As the electron wave changes its frequency, it also changes its energy. The different waves correspond to the Bohr model’s discrete orbits.
Each electron wave wraps around the nucleus and has to meet itself in the right way to create a standing wave. (A vibrating guitar string is a combination of several standing waves corresponding to the fundamental pitch and the harmonics.) The radius of each Bohr orbit then must be equal to a whole-number multiple of an electon wavelength, or the standing wave “dies out.” With this simple, though bizarre concept, de Broglie provided a physical explanation for Bohr’s discrete electron orbits: the first Bohr orbit corresponded to one wavelength, the second to two wavelengths, and so on.
Electron waves were detected in 1927. Two years later, de Broglie won the Nobel Prize in Physics. The first technology to harness electron waves was the electron microscope, developed soon afterward in the the 1930s.
Other physicists, Erwin Schroedinger, Werner Heisenberg, Paul Dirac, and Wolfgang Pauli, among others, refined the quantum model of the atom further. The Schroedinger equation enables us, for example, to predict the probable location of an electron around its parent nucleus.
The key word there is “probable.” Because of the electron’s dual wave-particle nature, we cannot pinpoint its location and speed simultaneously with the same precision (that’s from Heisenberg’s Principle of Uncertainty.)
So, the electron is not really a speck orbiting a much larger sphere. For that matter (pun intended!), the proton is not really a solid sphere, since it also has a dual wave-particle nature. A better picture would be of a dense, slightly fuzzy-around-the-edges cloud surrounded by a larger, fuzzier cloud that gradually thins out at its outer perimeter.
Further, there is no requirement that the electron cloud be spherical. Electron probability clouds can have a variety of shapes, including dumbbells, donuts, teardrops and spheres.
Have we ever seen atoms? Yes. Electron micrographs reveal atoms (actually their outer electron clouds) as, well, fuzzy balls.
So, not only are atoms mostly empty space, their constituent parts are not even solid! Intermolecular forces give us the illusion of solidity, and prevent our constituent molecules from ever actually touching any other molecules. Nevertheless, those forces can be pretty substantial as two surfaces approach each other, such as a bat hitting a ball, or a head hitting a table.
Atoms may be mostly empty space and mostly ephemeral, but they can still leave nasty bruises. Watch your head!