Atoms are the building blocks of all matter. They're made up of three subatomic particles: protons, neutrons, and electrons, each arranged in a specific way. For astronomy, the most important thing about atoms is this: when electrons jump between energy levels, they absorb or release light at very specific wavelengths. That's what produces the spectra we use to study stars, nebulae, and galaxies from enormous distances.

Structure of Atoms

Every atom has a nucleus at its center containing protons (positive charge) and neutrons (no charge). The nucleus holds about 99.9% of the atom's mass, even though it's incredibly small compared to the atom as a whole.

Electrons are negatively charged particles that exist in regions around the nucleus called energy levels or shells. In a neutral atom, the number of electrons equals the number of protons, so the charges balance out.

The atomic number (number of protons) uniquely identifies each element. Hydrogen has 1 proton, helium has 2, carbon has 6, and so on.
The number of electrons in the outermost shell (called valence electrons) determines how an atom bonds with other atoms.
The chemical properties of an element depend on its electron configuration, but for astronomy, we care most about how electrons interact with light.

Electron Energy Transitions

This is the section that connects atoms to spectra, so it's worth understanding well.

Electrons occupy discrete energy levels around the nucleus, labeled with integers: $n = 1, 2, 3, ...$ The level closest to the nucleus ( $n = 1$ ) has the lowest energy, and energy increases as $n$ gets larger. Electrons can't exist between these levels. They can only occupy specific allowed energies.

When an electron moves between levels, it must absorb or emit a precise amount of energy:

Absorption: An electron jumps from a lower level to a higher one by absorbing a photon. The photon's energy must exactly match the energy difference between the two levels.
Emission: An electron drops from a higher level to a lower one, releasing a photon. The photon carries away energy equal to the difference between those levels.

The energy of the photon determines its wavelength and color. Because each element has a unique set of energy levels, each element absorbs and emits its own characteristic wavelengths of light. This is why electron transitions produce atomic spectra:

Emission spectra are the bright lines of color you see when an element is heated or energized. For example, hydrogen produces a distinctive pattern of red, blue-green, blue, and violet lines.
Absorption spectra are dark lines that appear in an otherwise continuous spectrum. These dark lines show up at the exact same wavelengths as the emission lines for that element. The Fraunhofer lines in the solar spectrum are a classic example: cooler gas in the Sun's outer atmosphere absorbs specific wavelengths from the continuous light below.

This is how astronomers identify what distant stars and nebulae are made of without ever visiting them.

Isotopes and Atomic Properties

Isotopes are atoms of the same element that have different numbers of neutrons. They have the same number of protons (so they're the same element) and the same number of electrons (so they behave the same chemically), but their masses differ.

Isotopes are identified by their mass number ( $A$ ), which equals the number of protons plus neutrons. For example:

Carbon-12: 6 protons + 6 neutrons ( $A = 12$ )
Carbon-14: 6 protons + 8 neutrons ( $A = 14$ )

Some isotopes are unstable and undergo radioactive decay, transforming into other elements or isotopes over time. The three main types of decay are:

Alpha decay: the nucleus emits an alpha particle (essentially a helium-4 nucleus, with 2 protons and 2 neutrons)
Beta decay: a neutron converts into a proton, emitting an electron and an antineutrino
Gamma decay: the nucleus releases energy as a high-energy photon (gamma ray), without changing its composition

The average atomic mass of an element listed on the periodic table is a weighted average based on the natural abundance of its isotopes. Chlorine, for instance, averages about 35.5 atomic mass units because chlorine-35 makes up roughly 75.8% of natural chlorine and chlorine-37 makes up about 24.2%.

Atomic Models and Quantum Mechanics

Our understanding of the atom has evolved through several models, each improving on the last:

The Rutherford model (1911) established that atoms have a small, dense, positively charged nucleus with electrons somewhere around it. But it couldn't explain why electrons don't spiral into the nucleus or why atoms emit light at only specific wavelengths.
The Bohr model (1913) solved this by proposing that electrons travel in fixed orbits at discrete energy levels. This worked well for hydrogen and correctly predicted its spectral lines, but it couldn't accurately handle atoms with more than one electron.
Quantum mechanics replaced the Bohr model with a more complete picture. Instead of neat orbits, electrons exist in orbitals, which are probability clouds describing where an electron is likely to be found. Quantum mechanics also introduced wave-particle duality: electrons (and photons) behave as both particles and waves depending on the situation.

For this course, the Bohr model is a useful way to visualize energy levels and electron transitions. Just keep in mind that the real picture is more complex, and quantum mechanics is what actually governs how atoms behave.

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