Basics — what an atom is made of
Before models and numbers, let's start with the three particles that make up every atom. Here are their definitions, one by one.
Atom
An atom is the smallest unit of a chemical element that retains all of its chemical properties. It consists of a small, dense nucleus at the center and electrons occupying the space around it. Almost all of the atom's mass is in the nucleus, but almost all of its volume is empty space around it.
Nucleus
The nucleus is a small, dense region at the center of the atom where almost all of its mass is located. It consists of two types of particles: protons and neutrons, bound tightly together by the strong nuclear force. Although the nucleus is only a tiny fraction of the size of the whole atom (like a pea in a football stadium), it carries over 99.9% of the mass.
Proton
A proton is a particle with a positive charge, located in the nucleus. The number of protons (called the atomic number, Z) determines which chemical element it is — 1 proton is always hydrogen, 6 protons is always carbon, 79 protons is always gold. Change the number of protons and you get a completely different element.
Neutron
A neutron is a particle with no charge (neutral), roughly the same mass as a proton, located together with the protons in the nucleus. The number of neutrons can vary among atoms of the same element — these variants are called isotopes (e.g. carbon-12 and carbon-14 have the same number of protons but a different number of neutrons).
Electron
An electron is a particle with a negative charge, thousands of times lighter than a proton or neutron. Electrons aren't found in the nucleus, but in the space around it, arranged in shells (energy levels). The number of electrons in the outer shell determines how the atom behaves chemically — whether it easily forms bonds with other atoms, and what kind.
In short — the three particles compared
| Particle | Charge | Location | Relative mass | Color in this app |
|---|---|---|---|---|
| Proton | Positive (+1) | Nucleus | 1 | |
| Neutron | Neutral (0) | Nucleus | ~1 | |
| Electron | Negative (-1) | Shells around the nucleus | ~1/1836 |
Numbers and units you'll encounter in the app
What does "10⁻¹⁰" or "10⁸" mean?
Atoms and their parts are so small (or some numbers in physics so large) that it would be impractical to write out all the zeros. That's why scientific (exponential) notation is used:
10³ = 1,000 (a 1 followed by 3 zeros)
10⁻³ = 0.001 (one divided by 1,000 — move the decimal point 3 places left)
So: a positive exponent (up and to the right of the 10) means a big number — you add zeros. A negative exponent means a small number, less than one — you move the decimal point to the left. The "more negative" the exponent, the smaller the number.
Example from the app: an atom is about 10⁻¹⁰ meters in size. That means: 0.0000000001 meters — ten zeros after the decimal point before the first digit appears.
What is a picometer (pm)?
The meter is too large a unit for measuring atoms, so much smaller units are used:
1 millimeter (mm) = 0.001 m
1 nanometer (nm) = 0.000000001 m (10⁻⁹ m)
1 picometer (pm) = 0.000000000001 m (10⁻¹² m) — a thousand times smaller than a nanometer
To picture it: when the app shows an atomic radius of, say, "70 pm" (oxygen) or "260 pm" (cesium), those are differences of tens of picometers — but even the largest atom is still unimaginably small. A human hair is about 70,000,000 pm thick — roughly a million times thicker than a single atom.
Atomic number vs. atomic mass — not the same thing
Atomic number (Z) = number of protons. This is the element's "ID number" — every element has exactly one, fixed atomic number (hydrogen is always 1, oxygen is always 8).
Atomic mass = the combined mass of protons and neutrons (electrons are too light to contribute significantly). That's why atomic mass is always greater than the atomic number (except for hydrogen, which has no neutrons) — the difference between them actually tells you the approximate number of neutrons.
Example: carbon has atomic number 6 (6 protons) and atomic mass ~12 (6 protons + 6 neutrons). When you see two numbers next to an element in the periodic table — one small (Z) and one larger (mass) — now you know the difference.
Natural vs. artificial (synthetic) elements
Natural elements occur in nature — in the Earth's crust, air, water, stars — and no one had to make them. Most of the 118 known elements (about 112) fall into this group.
Artificial (synthetic) elements don't occur in nature (or only appear in trace amounts) — scientists made them in a laboratory or nuclear reactor, usually by colliding atoms at high speed. These are: technetium (Z=43) and promethium (Z=61), plus every element from americium (Z=95) upward — all the way to oganesson (Z=118).
Artificial elements are often highly unstable — some last only a fraction of a second before decaying (radioactive decay). In the "Periodic Table" module you can spot them by the dashed border and the ⚗ mark on the cell.
Interactive atom model
Pick an element and see a simplified Bohr-style view of how its electrons are distributed across shells (K, L, M, N...). This is a didactic model — the real behavior of electrons is described in the "Quantum Concepts" module.
Zoomed in: inside the nucleus
The nucleus of the selected element, enlarged — each dot is one proton or one neutron.
Periodic table of elements
All 118 elements, colored by category. Click any element for details: atomic number, mass, electron configuration by shell, and atomic radius.
Comparing atom sizes
Atomic radius is measured in picometers (1 pm = 10⁻¹² m). Pick up to 6 elements to see their relative sizes, or look at the trend across the periodic table below.
Click a chip below to remove it. Maximum 6 elements at a time.
pm = picometer — the unit used to measure atomic radius. 1 pm = one thousandth of a nanometer = 10⁻¹² meters (0.000000000001 m). For comparison: a human hair is about 70,000,000 pm thick.
Why does atomic radius change?
Down a group (top to bottom): radius increases — each new row adds a new electron shell, so the outermost electron sits farther from the nucleus.
Across a period (left to right): radius decreases — the number of protons (and thus the nucleus's positive charge) grows faster than the number of shells, so the nucleus pulls electrons in more tightly. This effect is called effective nuclear charge.
That's why the largest atoms sit in the bottom-left corner of the periodic table (e.g. cesium, francium), and the smallest in the upper-right (e.g. fluorine, helium — excluding the noble gases, whose "radius" is theoretical/calculated since they don't form bonds).
Quantum concepts — from Bohr to today
A short, substantial overview of the key ideas describing how atoms behave at the quantum level.
1. From a solid ball to a cloud of probability
Dalton (~1803): the atom is an indivisible, solid sphere.
Thomson (1897): the "plum pudding" model — electrons scattered within a positively charged mass.
Rutherford (1911): the gold foil experiment showed that almost all of an atom's mass is concentrated in a small, positively charged nucleus — the atom is mostly empty space.
Bohr (1913): electrons orbit the nucleus in fixed, circular orbits with precisely defined energy levels — the first model to introduce quantization.
The quantum-mechanical model (from 1926, Schrödinger/Heisenberg): electrons can't be described as particles on a precisely defined path, but rather by wave functions that give the probability of finding an electron in a given region of space — the so-called orbital.
2. s, p, d, f orbitals
Each shell contains sublevels (orbitals) of different shapes. The number of electrons they can hold: s = 2, p = 6, d = 10, f = 14.
s — sphere
p — two lobes
d — four lobes
f — complex, multi-lobed
3. Why electrons "jump" between levels
An electron in an atom can't have just any energy — only precisely defined, discrete values. When an electron absorbs a photon of the right energy, it "jumps" to a higher level; when it drops back to a lower level, it emits a photon — this is the basis of atomic spectra (and, later, lasers).
4. Heisenberg's uncertainty principle
It's impossible to know both the position and momentum of an electron precisely at the same time. The more precisely we know one, the less precisely we can know the other. This isn't a limitation of measuring instruments — it's a fundamental property of nature at the quantum level.
5. Wave-particle duality
Electrons (and light) show properties of both particles and waves, depending on the experiment. The double-slit experiment shows that electrons, sent one at a time, form an interference pattern characteristic of waves — even when passing through "one by one".
6. The Pauli exclusion principle and spin
No two electrons in the same atom can have an identical set of four quantum numbers. That's why every orbital holds at most 2 electrons — and they must have opposite spin (an additional quantum property, often described as "up"/"down"). This principle determines how electrons are arranged across the entire periodic table.
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