Atomic Structure: Protons, Neutrons, Electrons, and Isotopes

Everything around you — the air you breathe, the food you eat, the device you are reading this on — is made of atoms. Atoms are the fundamental building blocks of matter, and understanding their structure is the first step toward understanding all of chemistry. In this guide, we will explore what atoms are made of, how scientists developed our current model of the atom, and how concepts like atomic number, mass number, isotopes, and ions help us describe and distinguish the elements.

What is an atom?

An atom is the smallest unit of an element that retains the chemical properties of that element. Atoms are incredibly small: a single sheet of paper is about a million atoms thick. Despite their tiny size, atoms have an internal structure that determines how they behave in chemical reactions. Every atom consists of a dense central nucleus surrounded by a cloud of electrons. The nucleus contains protons and neutrons, while the electrons occupy the space around the nucleus in regions called orbitals.

Brief history of atomic models

Our understanding of the atom has evolved dramatically over the past two centuries. In 1803, John Dalton proposed that all matter is composed of indivisible atoms, that atoms of the same element are identical, and that chemical reactions involve the rearrangement of atoms. Dalton's model was groundbreaking but incomplete because it did not account for the internal structure of atoms.

In 1897, J.J. Thomson discovered the electron using cathode ray tubes and proposed the plum pudding model, in which negatively charged electrons were embedded in a positively charged sphere, like raisins in a pudding. This model was overturned in 1911 when Ernest Rutherford conducted his famous gold foil experiment. By firing alpha particles at a thin gold foil, Rutherford found that most particles passed straight through, but a small fraction were deflected at large angles. He concluded that the atom's positive charge and most of its mass are concentrated in a tiny, dense nucleus, with electrons orbiting at a distance.

In 1913, Niels Bohr refined Rutherford's model by proposing that electrons orbit the nucleus in specific energy levels, or shells, much like planets orbit the sun. Electrons could jump between energy levels by absorbing or emitting specific amounts of energy, which explained the line spectra of elements. The Bohr model works well for hydrogen but fails for more complex atoms.

The modern quantum mechanical model, developed in the 1920s by Schrödinger, Heisenberg, and others, replaces fixed orbits with probability distributions called orbitals. An orbital describes the region of space where an electron is most likely to be found. This model accurately predicts the behavior of all elements and is the foundation of modern chemistry. While the quantum model is mathematically complex, the key takeaway for introductory chemistry is that electrons exist in specific energy levels and sublevels (s, p, d, f orbitals) around the nucleus.

Subatomic particles: protons, neutrons, electrons

Atoms are composed of three subatomic particles. Protons are positively charged particles found in the nucleus. Each proton has a charge of +1 and a mass of approximately 1 atomic mass unit (amu), or 1.673 × 10⁻²⁴ grams. Neutrons are electrically neutral particles also found in the nucleus. Each neutron has a mass of approximately 1 amu, nearly identical to the proton. Electrons are negatively charged particles that occupy the space around the nucleus. Each electron has a charge of −1 and a mass of approximately 0.00055 amu, which is about 1/1836 the mass of a proton. Because electrons are so much lighter than protons and neutrons, the mass of an atom is almost entirely determined by its nucleus.

The number of protons in an atom defines which element it is. All carbon atoms have 6 protons, all oxygen atoms have 8 protons, and all gold atoms have 79 protons. If you change the number of protons, you change the element. In a neutral atom, the number of electrons equals the number of protons, so the overall charge is zero.

Atomic number and mass number

The atomic number (Z) of an element is the number of protons in the nucleus of each atom of that element. The atomic number uniquely identifies an element: hydrogen has Z = 1, helium has Z = 2, lithium has Z = 3, and so on. The periodic table is arranged in order of increasing atomic number.

The mass number (A) is the total number of protons and neutrons in the nucleus. It is always a whole number. For example, a carbon atom with 6 protons and 6 neutrons has a mass number of 12. A carbon atom with 6 protons and 7 neutrons has a mass number of 13. Both are carbon because they have the same number of protons, but they differ in their number of neutrons.

The standard notation for an isotope places the mass number as a superscript and the atomic number as a subscript to the left of the element symbol. For example, carbon-12 is written as ¹²C and carbon-13 as ¹³C. In many textbooks, the atomic number is omitted from the notation because the element symbol already tells you the atomic number.

How to calculate the number of each particle

Given the atomic number and mass number, you can calculate the number of each subatomic particle. The number of protons equals the atomic number. The number of neutrons equals the mass number minus the atomic number (N = A − Z). The number of electrons in a neutral atom equals the number of protons. For example, an atom of fluorine-19 has Z = 9 and A = 19. It has 9 protons, 19 − 9 = 10 neutrons, and 9 electrons. If the atom is an ion, the number of electrons changes: a cation has fewer electrons than protons, and an anion has more electrons than protons.

Isotopes and atomic mass

Isotopes are atoms of the same element that have different numbers of neutrons and therefore different mass numbers. For example, carbon has three naturally occurring isotopes: carbon-12 (6 protons, 6 neutrons), carbon-13 (6 protons, 7 neutrons), and carbon-14 (6 protons, 8 neutrons). All three isotopes are carbon because they all have 6 protons, but they differ in mass. Carbon-12 is the most abundant, making up about 98.9% of natural carbon. Carbon-13 accounts for about 1.1%, and carbon-14 is present in trace amounts.

The atomic mass listed on the periodic table for each element is a weighted average of the masses of all naturally occurring isotopes, taking into account their relative abundances. For carbon, the atomic mass is 12.011 amu, which is very close to 12 because carbon-12 is by far the most abundant isotope. For chlorine, the atomic mass is 35.45 amu because chlorine has two common isotopes: chlorine-35 (about 75.8% abundance) and chlorine-37 (about 24.2% abundance). The weighted average reflects the contribution of each isotope.

To calculate the atomic mass from isotope data, multiply the mass of each isotope by its fractional abundance and add the results. For chlorine: (34.969 × 0.758) + (36.966 × 0.242) = 26.527 + 8.946 = 35.47 amu, which rounds to the value on the periodic table.

Ions: cations and anions

An ion is an atom that has gained or lost one or more electrons, giving it a net electrical charge. When an atom loses electrons, it becomes a positively charged ion called a cation. Metals typically form cations. For example, sodium (Na) loses one electron to become Na⁺, and calcium (Ca) loses two electrons to become Ca²⁺. When an atom gains electrons, it becomes a negatively charged ion called an anion. Nonmetals typically form anions. For example, chlorine (Cl) gains one electron to become Cl⁻, and oxygen (O) gains two electrons to become O²⁻.

The charge of an ion tells you how many electrons were gained or lost. A +2 charge means the atom lost 2 electrons, so it has 2 fewer electrons than protons. A −1 charge means the atom gained 1 electron, so it has 1 more electron than protons. Ions form because achieving a full valence shell (often 8 electrons, following the octet rule) is energetically favorable.

The nucleus vs electron cloud

The nucleus is incredibly small compared to the overall size of the atom. If an atom were the size of a football stadium, the nucleus would be about the size of a marble at the center of the field. Despite its tiny size, the nucleus contains more than 99.9% of the atom's mass because protons and neutrons are nearly 2,000 times heavier than electrons.

The electron cloud is the region of space around the nucleus where electrons are likely to be found. In the quantum mechanical model, electrons do not travel in fixed circular paths. Instead, they exist in orbitals, which are three-dimensional probability distributions. The shape and energy of each orbital depend on the quantum numbers that describe it. The s orbitals are spherical, the p orbitals are dumbbell-shaped, the d orbitals have more complex shapes, and the f orbitals are even more intricate. The arrangement of electrons in these orbitals, known as the electron configuration, determines the chemical properties of the element.

How atomic structure determines chemical behavior

The chemical behavior of an element is determined primarily by its electron configuration, especially the number of valence electrons. Elements in the same group of the periodic table have the same number of valence electrons and therefore exhibit similar chemical properties. Sodium and potassium are both in group 1, both have one valence electron, and both are highly reactive metals that form +1 ions. Fluorine and chlorine are both in group 17, both have seven valence electrons, and both are highly reactive nonmetals that form −1 ions.

The tendency of atoms to gain, lose, or share electrons to achieve a stable electron configuration drives all chemical bonding. Metals with few valence electrons tend to lose them and form ionic bonds with nonmetals. Nonmetals with nearly full valence shells tend to share electrons with other nonmetals, forming covalent bonds. Understanding atomic structure is therefore the key to understanding why elements react the way they do and why the periodic table is organized the way it is.

Exploring elements in MyChemLab AI's periodic table

Learning about atomic structure becomes much more engaging when you can interact with the data. MyChemLab AI's interactive periodic table lets you click on any element to see its atomic number, mass number, electron configuration, common isotopes, and ion charges. You can compare elements side by side to see how their atomic structures differ and how those differences translate into different chemical behaviors. The AI tutor can walk you through calculating the number of subatomic particles for any isotope or ion, and it provides instant feedback so you can correct mistakes in real time. Whether you are studying for a quiz on atomic structure or exploring the elements out of curiosity, MyChemLab AI makes the invisible world of atoms visible and understandable.