Atoms are composed of a dense nucleus containing positively charged protons and electrically neutral neutrons, surrounded by negatively charged electrons that occupy defined orbitals. Protons and neutrons are about 2,000 times heavier than electrons, so they account for almost all of an atom’s mass. For every element, the number of protons in the nucleus is fixed—carbon, for example, always has six protons. In a neutral atom, the number of electrons equals the number of protons, but electrons can be gained or lost during chemical reactions. The neutron count varies among atoms of the same element, giving rise to isotopes. Mastering these concepts lets you calculate the subatomic composition of any isotope.
The mass number appears as a superscript before the element’s symbol or after a hyphen, e.g., 235U or U‑235. This number represents the total count of protons plus neutrons.
Locate the element’s atomic number on the periodic table. The atomic number is the count of protons in every atom of that element. Uranium (U) has an atomic number of 92, so every uranium nucleus contains 92 protons.
Check the symbol for a charge notation— a superscripted positive or negative number (e.g., 235U(4⁺)). A positive charge means the atom has lost that many electrons; a negative charge indicates gained electrons. If no charge is shown, the atom is neutral, and its electron count equals its proton count. For 235U(4⁺), the electron count is 92 – 4 = 88.
Subtract the proton count from the mass number: Neutrons = Mass Number – Protons. For 235U, that’s 235 – 92 = 143 neutrons.
