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The Basics of Atomic Structure

This guide walks through the atom one layer at a time, from the particles inside the nucleus to the way electrons are arranged around it, with worked examples so you can apply the ideas to any element on the periodic table.

  1. The location and charge of protons, neutrons, and electrons

    An atom consists of a nucleus and the electrons surrounding it. Inside the nucleus, positively (+) charged protons and uncharged neutrons are packed together, while negatively (−) charged electrons move rapidly around the outside of the nucleus. The nucleus takes up only a tiny fraction of the atom's total volume but accounts for most of its mass. In a neutral atom the number of electrons equals the number of protons, so the positive and negative charges cancel out exactly.

  2. The meaning of atomic number and mass number

    The atomic number indicates the number of protons in the nucleus, and this is the unique value that determines which element it is. The mass number is the sum of the number of protons and neutrons, giving an approximate figure for the mass of the nucleus. For example, carbon has an atomic number of 6, meaning it has 6 protons, and its most common isotope, carbon-12, has a mass number of 12 (6 protons plus 6 neutrons). Change the proton count and you have a different element; change only the neutron count and you still have the same element.

  3. Reading an element symbol: how to find protons, neutrons and electrons

    Isotope notation puts the mass number at the upper left of the symbol and the atomic number at the lower left. Protons equal the atomic number, neutrons equal the mass number minus the atomic number, and electrons equal the protons in a neutral atom. Chlorine-35 (atomic number 17) therefore has 17 protons, 18 neutrons and 17 electrons. Uranium-238 (atomic number 92) has 92 protons, 146 neutrons and 92 electrons. For an ion, adjust the electron count by the charge: a Mg²⁺ ion made from magnesium (12 protons) has 12 protons and 10 electrons.

  4. The concept of electron shells (energy levels)

    Electrons don't orbit the nucleus just anywhere — they fill a set number of positions within fixed energy-level orbits (electron shells), such as the K, L, M, and N shells. The maximum number of electrons a given shell can hold follows the formula 2n² (where n is the shell number), which gives 2, 8, 18 and 32 for the first four shells. The outermost electrons in the outermost shell, called valence electrons, determine chemical bonding and reactivity.

  5. Filling the shells: oxygen, sodium and neon as examples

    Electrons fill the lowest-energy shells first. Oxygen, with 8 electrons, is arranged 2 in the first shell and 6 in the second, so it is two electrons short of a full second shell and readily forms bonds. Neon, with 10 electrons, is arranged 2 and 8, a full outer shell, which is why it hardly reacts at all. Sodium, with 11 electrons, is arranged 2, 8 and 1, leaving a single loosely held outer electron that it gives up easily. The simple 2n² capacity works neatly for the first 18 elements; beyond that, the order of filling becomes more complicated, so treat the shell model as a first approximation.

  6. The concept of isotopes

    Isotopes are atoms that have the same number of protons (atomic number) but a different number of neutrons, giving them different mass numbers. Their chemical properties are nearly identical, but their mass and stability can differ, and some isotopes are unstable, decaying while emitting radiation as radioactive isotopes. For example, carbon-12 and carbon-14 are isotopes of the same element, carbon, but with different numbers of neutrons. Hydrogen has three well-known isotopes: protium (no neutron), deuterium (one neutron) and tritium (two neutrons, radioactive).

  7. Why the atomic mass on the periodic table is not a whole number

    The number printed under an element's symbol is a weighted average of the masses of its naturally occurring isotopes, not the mass of a single atom. Chlorine is the classic example: about three quarters of natural chlorine atoms are chlorine-35 and about one quarter are chlorine-37, so the average comes out near 35.45 rather than a whole number. Elements with only one stable isotope, such as fluorine, sit very close to a whole number.

  8. Unstable nuclei and radioactive decay

    A nucleus is stable only when its protons and neutrons are in a workable balance. Very heavy nuclei, or ones with too many or too few neutrons, can shed particles or energy to reach a more stable state. Alpha decay emits a helium nucleus, beta decay converts a neutron into a proton (or the reverse) and emits an electron or positron, and gamma emission releases energy only. The time for half of a sample to decay, its half-life, is characteristic of each isotope: carbon-14 has a half-life of about 5,730 years, which is why it is useful for dating organic remains from the last tens of thousands of years.

  9. The difference between an atom and a molecule

    An atom is the smallest basic unit that makes up matter, while a molecule is a particle formed when two or more atoms are joined by chemical bonds, giving it independent properties. For example, a single oxygen atom (O) is highly unstable, but an oxygen molecule (O₂), formed by two oxygen atoms bonding together, exists as the stable gas we breathe. When atoms of different elements bond together, as in a water molecule (H₂O), the result is called a compound.

  10. The principle behind how ions form

    When an atom loses or gains electrons, the number of protons and electrons no longer match, and the atom becomes a charged ion. Losing an electron creates a positively (+) charged cation, while gaining an electron creates a negatively (−) charged anion. When a sodium atom loses one electron to become Na⁺ and a chlorine atom gains one electron to become Cl⁻, and the two bond together, sodium chloride (table salt) is formed. The proton count never changes during this process, which is why a sodium ion is still sodium.

  11. The historical development of atomic models

    In 1897 J. J. Thomson showed that atoms contain tiny negative particles, later named electrons. In 1911, Rutherford discovered that a small, heavy nucleus exists at the center of an atom through an experiment in which his team fired alpha particles at thin gold foil: most passed straight through, but a tiny fraction bounced back sharply. Afterward, in 1913, Bohr proposed a model in which electrons exist only in specific orbits (energy levels), and James Chadwick identified the neutron in 1932. Modern quantum mechanics uses an electron cloud (orbital) model that represents the probability of finding an electron rather than its exact location.

Why do we need to understand atoms?

Every substance around us is made up of extremely small atoms invisible to the eye. Understanding the internal structure and properties of atoms allows us to logically explain how substances bond to form molecules and why different elements have different chemical properties. It also underlies practical topics such as batteries, medical imaging, nuclear power, semiconductor chips and the way sunscreens, medicines and metals behave. Once you are comfortable with atomic number and outer electrons, take a look at our guide to reading the periodic table as well, because its layout follows directly from atomic structure.

How big is an atom, really?

A typical atom is about one ten-billionth of a meter across, roughly 0.1 nanometers. The nucleus is smaller still: its diameter is on the order of 10,000 to 100,000 times smaller than the atom as a whole, depending on the element. If a nucleus were the size of a marble, the electrons would be found hundreds of meters away, and the space in between is not filled with matter in the everyday sense. Solid objects feel solid because of the electric repulsion between the electron clouds of neighboring atoms, not because atoms are packed full of material.

Common misconceptions

The most common mistake is to picture electrons as little planets on neat circular tracks. The Bohr picture is a useful teaching model, and it correctly explains why electrons occupy discrete energy levels, but modern physics describes regions of probability instead of paths. A second mistake is to think that the mass number is the same as the atomic mass on the periodic table; the first is a whole number for one specific isotope, the second is an average over natural isotopes. A third is to say that an atom loses protons to become an ion: ions form by gaining or losing electrons, and the nucleus stays the same in ordinary chemical reactions.

Frequently Asked Questions

How small a portion of an atom does the nucleus take up?

The nucleus is extremely small, roughly between one ten-thousandth and one hundred-thousandth of the atom's diameter depending on the element, but it accounts for most of the atom's mass (over 99.9%). Most of an atom's space is effectively the near-empty region where electrons are active.

Are there atoms without any neutrons?

Yes — the most common hydrogen atom (hydrogen-1) has just one proton and one electron, with no neutrons. Deuterium, an isotope of hydrogen, has one neutron, and tritium has two.

Why do isotopes of the same element behave almost identically in chemistry?

Chemical behavior is controlled mainly by the electrons, and the number and arrangement of electrons depends on the proton count, which isotopes share. Extra neutrons change the mass and sometimes the stability of the nucleus, so isotopes differ in properties such as reaction rate (slightly) and radioactivity, but the kinds of bonds they form are the same.

What is the difference between an atom, an ion and a molecule?

An atom is a single unit of an element with equal numbers of protons and electrons. An ion is an atom (or a group of atoms) that has gained or lost electrons and so carries a net charge. A molecule is two or more atoms held together by shared electrons, such as O₂ or H₂O. Table salt is made of ions arranged in a crystal rather than of discrete molecules.