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Atoms consist of a nucleus containing protons and neutrons, surrounded by electrons in orbitals. Protons are positively charged, neutrons are neutral, and electrons are negatively charged. The identity of an element is determined by the number of protons (atomic number), while the mass is primarily from protons and neutrons.
Use Rutherford's scattering experiment to motivate the nuclear model. Build physical models of atoms. Draw Bohr models initially, then transition to orbital representations.
If you have studied matter classification, you know that elements are pure substances that cannot be broken down further by chemical means. But what makes one element different from another at the subatomic level? The answer lies in the structure of the atom — specifically, in the number of protons.
Every atom contains a nucleus — a tiny, dense core — surrounded by a diffuse cloud of electrons. The nucleus holds protons (positively charged) and neutrons (electrically neutral). Electrons (negatively charged) occupy the space around the nucleus. In a neutral atom, the number of electrons equals the number of protons, so the atom carries no net charge. The atomic number (Z) is the count of protons, and it completely defines which element you have: 1 proton = hydrogen, 6 protons = carbon, 79 protons = gold, no exceptions. If you change the proton count, you have a different element.
The mass number (A) counts the total number of protons plus neutrons. Since protons and neutrons each have roughly 1 atomic mass unit (amu) and electrons contribute almost nothing to mass, the mass number is a good approximation of atomic mass. The number of neutrons in an atom can vary without changing its element — atoms of the same element with different neutron counts are called isotopes. Carbon-12 has 6 protons and 6 neutrons; carbon-14 has 6 protons and 8 neutrons. Both are chemically carbon, but carbon-14 is radioactive, which is the basis of radiocarbon dating.
A persistent misconception, reinforced by textbook diagrams, is that electrons orbit the nucleus like tiny planets — this is the Bohr model, useful for introductory calculations but ultimately wrong. Quantum mechanics tells a different story: electrons exist as probability distributions, or orbitals, defined by wave functions. You cannot know simultaneously where an electron is and how fast it is moving (Heisenberg uncertainty principle). When we draw an orbital, we are drawing a region of space within which there is a high probability (typically ~90%) of finding the electron. The electron does not trace a path; it is delocalized.
The three particles differ sharply in mass. Protons and neutrons each weigh approximately 1.67 × 10⁻²⁷ kg. An electron weighs roughly 1/1836 as much — so tiny that electrons contribute essentially nothing to atomic mass, but everything to chemistry, since it is the electrons that participate in bonding, reactions, and energy absorption.