Introduction to the Atom
Protons, Neutrons, Electrons & Atomic Structure
History of Atomic Theory
The idea that matter is made of indivisible particles dates back to ancient Greece, but the scientific atomic theory developed only over the last two centuries — through a series of experiments that progressively revealed the atom’s true structure.
Each model below was a major advance over its predecessor, and each was eventually shown to be incomplete. This is how science works — models improve as new evidence emerges.
BC
Proposed that all matter is made of tiny, indivisible particles called atomos (Greek: uncuttable). Purely philosophical — no experimental evidence. The idea was largely ignored for 2,000 years in favour of Aristotle’s four-element theory.
The first scientific atomic theory. Proposed that: all matter is made of atoms; atoms of the same element are identical; atoms cannot be created or destroyed; atoms combine in fixed ratios to form compounds. Pictured atoms as solid, indestructible spheres — like billiard balls.
Discovered the electron through cathode ray experiments — proving atoms are not indivisible. Proposed a model of negatively charged electrons embedded in a diffuse sphere of positive charge, like plums in a pudding. First model to include subatomic particles.
Gold foil experiment: fired alpha particles at a thin gold foil and found that most passed straight through, but a few deflected sharply — even bouncing back. Concluded that the atom has a small, dense, positively charged nucleus at its centre, surrounded by mostly empty space where electrons orbit.
Proposed that electrons orbit the nucleus in fixed energy levels (shells) — like planets around a sun. Electrons can jump between shells by absorbing or emitting specific amounts of energy (photons). Explained the hydrogen emission spectrum and forms the basis of our Tier I understanding.
The modern model. Electrons do not travel in defined orbits — they exist in regions of probability called orbitals. The exact position and momentum of an electron cannot be simultaneously known (Heisenberg Uncertainty Principle). This model will be explored fully in Tier IV.
Atomic Structure
An atom consists of two regions: a tiny, dense nucleus at the centre, and an electron cloud (or electron shells) surrounding it. The nucleus contains protons and neutrons; electrons occupy the space outside.
The nucleus is extraordinarily small relative to the whole atom. If an atom were the size of a football stadium, the nucleus would be about the size of a marble — yet it contains almost all of the atom’s mass.
A single hydrogen atom has a diameter of about 0.1 nanometres (1 × 10⁻¹⁰ m). That’s so small that over 1 million hydrogen atoms placed side by side would fit across the width of a human hair. The nucleus is about 100,000 times smaller than the atom itself.
Subatomic Particles
Atoms are composed of three types of subatomic particles, each with a characteristic mass, charge, and location. Knowing these is the absolute foundation of all further chemistry.
| Particle | Symbol | Charge | Relative Mass | Location |
|---|---|---|---|---|
| Proton | p⁺ | +1 | 1 (1.673 × 10⁻²⁷ kg) | Nucleus |
| Neutron | n⁰ | 0 (neutral) | 1 (1.675 × 10⁻²⁷ kg) | Nucleus |
| Electron | e⁻ | −1 | 1/1836 (≈ negligible) | Electron shells (outside nucleus) |
The number of protons in the nucleus defines what element an atom is — this is the atomic number. Carbon always has 6 protons; oxygen always has 8. Change the number of protons and you change the element entirely.
In a neutral atom, the number of electrons equals the number of protons, so the positive and negative charges cancel out. The atom carries no overall charge.
Neutrons have no charge, but they play two critical roles: they add mass to the nucleus, and they act as a kind of “glue” — the strong nuclear force between neutrons and protons holds the nucleus together against the repulsion between the positively charged protons. Without neutrons, most nuclei would fly apart.
Atomic Number & Mass Number
Two numbers fully describe the composition of any atom’s nucleus. Understanding them lets you calculate the number of protons, neutrons, and electrons in any atom.
The number of protons in the nucleus. This number uniquely identifies the element — it never changes for a given element. Carbon always has Z = 6.
The total number of protons + neutrons in the nucleus. Because electron mass is negligible, this approximates the atom’s mass. A = Z + N (where N = neutrons).
Neutrons (N) = Mass number (A) − Atomic number (Z)
For a neutral atom: Electrons = Protons = Atomic number (Z)
For an ion: electrons = Z ± charge (add electrons for negative ions, remove for positive)
Isotopes
Isotopes are atoms of the same element that have the same number of protons but a different number of neutrons. Because the atomic number (proton count) is the same, they are the same element — but their mass numbers differ.
Isotopes have virtually identical chemical properties (because chemistry is governed by electrons, which are the same in all isotopes of an element), but they have different physical properties such as mass and nuclear stability.
Neutrons: 0
Abundance: 99.98%
Stable
Neutrons: 1
Abundance: 0.02%
Stable
Neutrons: 2
Abundance: trace
Radioactive
The three isotopes of hydrogen are the most well-known example. All three have 1 proton — making them hydrogen — but they have 0, 1, and 2 neutrons respectively. Deuterium is used in nuclear reactors and heavy water; tritium is radioactive and used in nuclear weapons research and self-luminous devices.
Because most elements exist as a mixture of isotopes in nature, the atomic mass shown on the periodic table is a weighted average of all naturally occurring isotopes, based on their abundance. This is why chlorine has an atomic mass of 35.5 — it’s a mix of ³⁵Cl (75%) and ³⁷Cl (25%), giving a weighted average that falls between the two.
Ions
A neutral atom has equal numbers of protons and electrons. When an atom gains or loses electrons, it becomes electrically charged — it becomes an ion.
A positively charged ion formed when an atom loses one or more electrons. Fewer electrons than protons. Metals typically form cations. E.g. Na⁺, Ca²⁺, Fe³⁺.
A negatively charged ion formed when an atom gains one or more electrons. More electrons than protons. Non-metals typically form anions. E.g. Cl⁻, O²⁻, N³⁻.
Ions are critical to chemistry — they are responsible for ionic bonding, electrical conductivity in solutions, nerve impulses in biology, and a vast range of chemical reactions. You will encounter them throughout this curriculum.
Isotopes differ in the number of neutrons — the number of protons (and therefore the element) stays the same. Ions differ in the number of electrons — the number of protons stays the same. Neither changes the element’s identity. Only changing the proton count changes the element.
Worked Examples
Method: Protons = atomic number (Z). Neutrons = A − Z. Electrons = protons ± ion charge.
Method: Multiply each isotope’s mass by its fractional abundance, then sum.
= 26.496 + 8.958
= 35.45 u
This matches the atomic mass of chlorine on the periodic table (35.45). The value falls between the two isotope masses, closer to ³⁵Cl because it is more abundant.
Practice Questions
Q1. What determines which element an atom belongs to?
Q2. An atom has atomic number 8 and mass number 18. How many neutrons does it have?
Q3. Which of the following correctly describes isotopes?
Q4. A Ca²⁺ ion has atomic number 20 and mass number 40. How many electrons does it have?
Q5. Rutherford’s gold foil experiment led to the conclusion that:
Key Takeaways
- Atomic theory developed progressively: Dalton (solid spheres) → Thomson (plum pudding) → Rutherford (nuclear model) → Bohr (electron shells) → Quantum model.
- An atom has a small, dense nucleus containing protons (+) and neutrons (neutral), surrounded by electrons (−) in shells.
- Protons define the element (atomic number Z). Neutrons add mass. Electrons determine chemical behaviour.
- Atomic number (Z) = number of protons. Mass number (A) = protons + neutrons. Neutrons = A − Z.
- In a neutral atom, electrons = protons. Gaining electrons gives a negative ion (anion); losing electrons gives a positive ion (cation).
- Isotopes are atoms of the same element with different numbers of neutrons — same chemical behaviour, different masses.
- The atomic mass on the periodic table is a weighted average of all naturally occurring isotopes.