An atom is made of three types of particles: protons and neutrons packed tightly in a central nucleus, and electrons moving rapidly through the space around it. Every element on the periodic table is defined by how many protons sit in its nucleus, a value called the atomic number. Understanding how these pieces fit together explains why elements behave differently, why some materials are radioactive, and how chemical reactions work at the most basic level.
The Three Subatomic Particles
Protons carry a positive electrical charge (+1) and have a relative mass of 1 atomic mass unit (amu). Neutrons have no charge at all and weigh almost exactly the same as a proton, so they also get assigned a mass of 1 amu. Both particles live inside the nucleus.
Electrons are a different story. Each electron carries a negative charge (-1) that perfectly balances one proton’s positive charge, but an electron weighs only about 1/1,850th as much as a proton. That means if you could pile 1,850 electrons on a scale, they’d barely match the mass of a single proton. Electrons move around the nucleus at speeds close to the speed of light, occupying regions of space called orbitals.
In a neutral atom (one with no overall electrical charge), the number of electrons equals the number of protons. A lithium atom, for example, has 3 protons and therefore 3 electrons. Change the number of protons and you change the element entirely.
Atomic Number and Mass Number
Two numbers define any atom. The atomic number (often written as Z) is simply the count of protons. Hydrogen has an atomic number of 1, carbon is 6, oxygen is 8. This number is what makes an element that element.
The mass number (A) is the total count of protons plus neutrons. Because electrons weigh almost nothing relative to the other particles, they don’t factor into mass number. If you know both values, you can figure out everything else. A lithium atom with Z = 3 and A = 7 has 3 protons, 3 electrons (since it’s neutral), and 4 neutrons (7 minus 3).
Inside the Nucleus
The nucleus is extraordinarily small compared to the atom as a whole. If the atom were the size of a football stadium, the nucleus would be roughly the size of a marble sitting at midfield. In actual numbers, the atom is about 100,000 times larger in diameter than its nucleus. Yet virtually all of the atom’s mass is concentrated in that tiny core, because protons and neutrons are so much heavier than electrons.
This raises an obvious question: protons are all positively charged, so they should repel each other violently at close range. What holds the nucleus together? The answer is a force called the strong nuclear force. It acts between all the particles in the nucleus, proton to proton, neutron to neutron, and proton to neutron, through the exchange of particles called mesons. This force is incredibly powerful but only works over extremely short distances, which is why the particles in the nucleus need to stay packed tightly together for it to hold.
Neutrons play a critical stabilizing role here. They add to the strong nuclear force without adding any electrical repulsion, so they help counterbalance the push that protons exert on each other. In heavier elements, the ratio of neutrons to protons increases to keep the nucleus stable. When that balance tips too far in either direction, the atom becomes radioactive.
Electron Shells and Energy Levels
Electrons don’t orbit the nucleus in neat circles like planets around a sun, though early models described them that way. The Danish physicist Niels Bohr proposed in 1913 that electrons travel in fixed orbits at specific distances from the nucleus, each orbit representing a set energy level. This model works reasonably well for simple atoms like hydrogen, but it breaks down for more complex ones.
The modern quantum mechanical model replaces those fixed paths with probability clouds. Instead of saying “the electron is here,” it says “the electron is most likely to be found in this region of space.” These regions are called orbitals, and they can have different shapes, from simple spheres to more complex dumbbell and clover-leaf patterns. The key shift is from certainty to probability: you can describe where an electron is likely to be, but you can never pin down its exact position and speed at the same time.
Despite this complexity, the concept of energy levels (or shells) remains useful. Each shell can hold a limited number of electrons, calculated with a simple formula: 2n², where n is the shell number. The first shell holds up to 2 electrons. The second holds up to 8. The third can accommodate 18, and the fourth maxes out at 32. Electrons fill the lowest energy shells first before moving to higher ones, which is why the periodic table is arranged in rows of increasing size.
Valence Electrons and Chemical Behavior
The electrons in the outermost shell of an atom are called valence electrons, and they’re the reason chemistry happens. Atoms tend to react with each other in ways that give them a full outer shell, a pattern known as the octet rule. Most atoms “want” eight electrons in their outermost shell, matching the stable configuration of noble gases like neon and argon. Hydrogen is an exception, satisfied with just two.
This drive toward a full outer shell explains all three major types of chemical bonding. Atoms can transfer electrons from one to another (forming ionic bonds, like in table salt), share electrons between them (forming covalent bonds, like in water), or pool their electrons in a communal “sea” (forming metallic bonds, like in copper wire). Every chemical reaction you’ve ever seen, from a rusting nail to a burning candle, is ultimately electrons rearranging to reach a more stable configuration.
Isotopes and Atomic Weight
While the number of protons defines an element, the number of neutrons can vary. Atoms of the same element with different neutron counts are called isotopes. Carbon, for instance, almost always has 6 neutrons (carbon-12), but some carbon atoms have 7 (carbon-13) or 8 (carbon-14). All three are still carbon because they all have 6 protons.
This is why the atomic weight listed on the periodic table is almost never a whole number. It’s not the mass of any single atom. Instead, it’s a weighted average that accounts for how common each isotope is in nature. Chlorine is a good example: about 75.8% of chlorine atoms have a mass number of 35, and 24.2% have a mass number of 37. Multiply each mass by its natural abundance and add them up, and you get 35.45, the decimal value you see on the periodic table. That number reflects the real-world mix of isotopes you’d encounter in any sample of chlorine.
Some isotopes are stable and will exist indefinitely. Others are unstable, meaning their nuclei eventually break apart and release energy, a process called radioactive decay. Carbon-14 is one of these unstable isotopes, and its predictable rate of decay is what makes carbon dating possible.
Why Atomic Structure Matters
Everything about how matter behaves traces back to atomic structure. The number of protons determines what element you’re dealing with. The number of neutrons determines which isotope it is and whether it’s stable or radioactive. The arrangement of electrons determines how that element bonds with others, what compounds it forms, and whether it conducts electricity, absorbs light, or dissolves in water. The periodic table itself is organized by atomic number and grouped by electron configuration, which is why elements in the same column share similar chemical properties. Sodium and potassium, for instance, both have a single electron in their outermost shell, which makes both of them soft, reactive metals that behave in strikingly similar ways.