Atoms, Sub-atomic Particles, and Elements

The word atom is derived from the Greek word atomos, which means “indivisible.” Atoms are the smallest unit that matter can be divided into and still retain the chemical properties of the bulk. Atoms are the smallest unit that matter can be divided into without the release of electrically charged particles. Atoms are actually made up of three types of particles: positively-charged protons, neutral neutrons, and negatively-charged electrons. Atoms can combine with other atoms to form molecules with chemical properties distinct from the composite elements.

Atomic theory has been refined over the centuries into the quantum model of the atom that we now use. In the fourth century B.C. the Greek philosopher Democritus considered what would happen if you broke a piece of matter in half, and then in half again, and again until you reached the smallest piece of matter possible. He called these most basic particles of matter atoms. Then, in 1803, John Dalton formulated the atomic theory of matter and the solid sphere model of the atom was born. In this model, atoms of a single element are identical and individual atoms cannot be divided, transformed, or destroyed.

We now know that an atom is indeed composed of smaller, sub-atomic particles. In 1887 J.J. Thomson discovered electrons, which he called “corpuscles,” during his work with cathode rays (electron beams), streams of electrons inside vacuum tubes. Thomson later proposed the plum pudding model of the atom which depicted an atom as a spherical cloud of positive charge in which negatively-charged electrons were scattered throughout. Ernest Rutherford, a former student of Thomson’s, discovered the nucleus in 1909. Rutherford introduced the nuclear model in which a diffuse cloud of electrons encompassed a small, compact, positively charged nucleus. The problem with this model was that it predicted all atoms were unstable—orbiting electrons would release radiation and spiral inwards collapsing on the nucleus.

In 1913 Neils Bohr amended Rutherford’s model by postulating that the electrons moved in stable orbits with quantized energy around the nucleus. His planetary model described the existence of electron shells or energy levels where electrons could be found. Electrons maintain these orbits as a result of attractive electrostatic forces. Electrons move between these energy levels by either absorbing or emitting energy. Then in 1926 Ernest Schrödinger gave us the quantum model of the atom which is in use today. In the quantum model, electrons move in clouds around the nucleus and their positions are not known exactly, but only in terms of probabilities, or orbitals.

  FIGURE in progress: Dalton Thomson Rutherford Bohr and Schrödinger models

At its most basic, an atom consists of a center nucleus, made up of protons and neutrons bundled together, and electrons that are in motion in orbit around the nucleus.

FIGURE in progress: electron cloud-style/quantum model atom showing protons and neutrons in the nucleus and a rim/cloud around showing areas of probable density. Such as a circle around the nucleus indicating average distance of the electron from the nucleus and a fuzzy cloud along the circle.

Sub-atomic Particles:
All fundamental particles fall into one of two camps, the fermions and the bosons. Fermions cannot occupy the same quantum state as one another. This means that they can’t occupy the same position in space, the same orbital state in an atom, or the same conduction band in a semiconductor. Famous fermions include electrons, protons, neutrons, and quarks, to name a few. Bosons don’t have this restriction and can pack infinitely many particles into the same quantum state. Some well-known bosons include helium, the Higgs boson, the photon, and other so-called force-carrying bosons.

Whether a particle is a fermion or a boson depends on a quantum property called spin. Bosons have integer spins like 0, 1, 2, etc. and fermions have half-integer spins like $\frac12, \frac32, \frac52,$ and so on. While beyond the scope of this wiki, the fact that fermions can’t occupy the same state is a result of the Pauli exclusion principle, which turns on the half-integer spin of fermions. The sub-atomic particles we will focus on here, protons, neutrons, and electrons, are all fermions.

Here's something interesting:

Depending on the total number of its protons and neutrons, the nucleus of an atom is a fermion or boson. This causes strange behavior in certain conditions. Helium’s nucleus consists of two neutrons and two protons making it a boson nucleus, and it never crystallizes, even when cooled to nearly absolute zero (the lowest temperature theoretically possible; 0 K, –273.15°C, or –459.67°F). Instead, it becomes a "superfluid," having zero viscosity and no surface tension.

Electrons:
Electrons are negatively charged elementary particles. When the number of electrons is equal to the number of protons, the atom is electrically neutral. Electrons are tiny—more than 1,800 times smaller than a proton or neutron—and nearly massless. An electron’s mass is just $9.109\,×\,10^{-31}$ kg; that’s a relative mass of 0.054% compared to the mass of a neutron. While electrons lack in mass, they make up most of the size of the atom as they can orbit quite far from the nucleus. How many electrons there are and how they are configured in space predicts atomic properties.

As mentioned earlier, it is not possible to know the position and momentum of an electron. The positions of the electrons are described in terms of probabilities and an electron can in theory be found at any distance from the nucleus. Still, there are certain regions that are much more probable to find an electron than others and these regions are called orbitals. Hund’s rule comes into play here—every orbital, where it is probable to find an electron, must be singly occupied before any one orbital can be doubly occupied; and all electrons in singly occupied orbitals have the same spin.

Protons & Neutrons:
Positively charged protons and electrically neutral neutrons comprise the small, dense nucleus of the atom—except hydrogen, whose nucleus is a single proton. But since protons are positively charged and thus repel each other, why doesn’t the nucleus fly apart? There are other forces at work in the nucleus, the strong force and the weak force. While the electromagnetic force that causes the like charges to repel each other works at large distances, the strong and weak forces in nucleus are only valid in such short distances of a couple of femtometers. The strong force holds the protons and neutrons together as it is more powerful than the electromagnetic force at small distances. The negatively-charged electrons around the nucleus are held in by the electromagnetic force, the Coulomb force—the force between two charged particles. Because the nucleus is positive and the electrons are negative, the force is attractive and the electrons stay close to the nucleus.

 FIGURE in progress: protons and neutrons in a nucleus and effective distance at which these forces are active

Protons and neutrons are themselves made up of other sub-atomic particles called quarks. As such, protons and neutrons are composite particles, meaning they are made up of other particle types. Quarks, like electrons, are called elementary or fundamental particles though it is not known whether or not they are composed of other particles.

Protons have a physical mass of $1.673\,×\,10^{-27}$ kg giving them slightly less mass than neutrons that weigh in at $1.6749\,×\,10^{-27}$ kg. Each proton is made up of two "up" quarks $\big($each has a $+\frac23$ charge$\big)$ and one "down" quark $\big($has a $-\frac13$ charge$\big)$ that are held together by gluons, another sub-atomic particle, that are massless.

Neutrons are neutral, sub-atomic particles found in every atomic nucleus except hydrogen. Neutrons are also composite particles made of quarks—one "up" quark and two "down" quarks. The protons and neutrons make up virtually all atomic mass as electrons are essentially massless.

If the weight of an electron is $9.109\,×\,10^{-31}$ kg and the weight of the proton is $1.673\,×\,10^{-27}$ kg, what percent of the mass of an atom would the electrons represent?

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Consider the element, chlorine. It has an atomic number of 17, meaning there are 17 protons, 17 neutrons, and 17 electrons. As there is a nearly negligible difference in mass between protons and neutrons, we can say that (mass of the proton) = (mass of the neutron).

So,

• 17 protons $\times 1.673\,×\,10^{-27}$ kg = $2.844\,×\,10^{-26}$ kg of protons
• 17 neutrons $\times 1.673\,×\,10^{-27}$ kg = $2.844\,×\,10^{-26}$ kg of neutrons
• 17 electrons $\times 9.109\,×\,10^{-31}$ kg = $1.549\,×\,10^{-29}$ kg of electrons.

The total mass of a chlorine atom is then

• $2.844\,×\,10^{-26} + 2.844\,×\,10^{-26} + 1.549\,×\,10^{-29} = 5.690\,×\,10^{-26}$ kg.

And the percentage of that mass that the electrons make up in a chlorine atom is

$$\frac{1.549\,×\,10^{-29}\text{ kg of electrons}}{5.690\,×\,10^{-26}\text{ kg total mass of a chlorine atom}} × 100 = 0.027\%.\ _\square$$

The number of protons in an atom defines what element it is and therefore what chemical behavior it exhibits. For example, hydrogen atoms have one proton, carbon atoms have six protons, and gold atoms have seventy-nine. All known elements are depicted on a chart called the periodic table of Elements. On the periodic table of Elements, the elements are arranged in order of increasing atomic number—the number of protons in an atom.

The number of neutrons in a nucleus determines the isotope of that element. Isotopes are atoms with the same number of protons, but differing numbers of neutrons. Isotopes are different forms of the same element but have different atomic weights.

 FIGURE in progress: Hydrogen has three known isotopes: protium (1H), deuterium (D or 2H), and tritium (T or 3H). Protium is just ordinary hydrogen - it has one proton, no neutrons, and one electron. Deuterium has one proton, one neutron, and one electron. Tritium has one proton, two neutrons, and one electron.