Protons, neutrons and electrons — mass, charge, location
Every atom has three sub-atomic particles. Protons (relative mass 1, charge +1) and neutrons (relative mass 1, charge 0) sit in the central nucleus; electrons (negligible mass, charge -1) occupy shells around it. The nucleus holds nearly all the mass (electrons are about 1836 times lighter), and a neutral atom has equal protons and electrons so the charges cancel. An atom is the smallest particle of an element; a molecule is two or more atoms joined by covalent bonds.
Atomic number, mass number and isotopes
Atomic number (Z) = protons, which defines the element. Mass number (A) = protons + neutrons. So neutrons = A minus Z, and a neutral atom has electrons = protons = Z. Isotopes are atoms of the same element with the same protons but different neutrons. With the same electron arrangement, isotopes share identical chemical properties and differ only in physical properties like density. Example: chlorine-35 and chlorine-37 both have 17 protons but 18 and 20 neutrons.
Relative atomic mass is a weighted average
Relative atomic mass (Ar) is the average mass of an atom relative to one-twelfth the mass of a carbon-12 atom, weighted by each isotope's natural abundance — so it is usually not a whole number. Calculate it as Ar = the sum of (isotope mass times percentage abundance), divided by the sum of the abundances (usually 100). The denominator is the total abundance, never the number of isotopes and never the sum of the mass numbers.
Drawn from real examiner reports.
Protons is not the mass number
Two slips when reading a nuclide. First, quoting the mass number as the number of protons — protons are the smaller atomic number, and neutrons come from subtracting: neutrons = A minus Z. Second, adding the electrons into the mass number, which counts only protons and neutrons (electrons have negligible mass). Fix the element from the atomic number first, then use A minus Z.
Flagged Jun 2024 P1CR Q1bi (protons confused with mass number; wrong neutron subtraction) and Q1bii (electrons wrongly added into the mass number); also Jun 2023 P1C Q7a.
Ar: divide by total abundance
The classic Ar error is dividing the sum of (mass times abundance) by the sum of the mass numbers (e.g. 79 + 81) instead of by the sum of the abundances (e.g. 52.8 + 47.2 = 100). A second error is copying the rounded Periodic-Table value (e.g. 80 for bromine) instead of the value the data give (79.9). Use the data-sheet value only as a sanity check, not as the answer.
Jun 2023 P1C Q7b — the most common error was dividing by the sum of the mass numbers instead of the sum of the abundances; candidates also rounded to the Periodic-Table value (80) rather than the calculated 79.9.
Isotopes: same protons, different neutrons
"Atoms that are different" or "atoms with different masses" is not enough for isotopes. The mark-scheme form is atoms of the same element (same number of protons) with different numbers of neutrons. Naming the shared protons AND the differing neutrons are two separate marking points — a one-element definition scores at most half the marks.
Ar definition must reference carbon-12
"The mass of an atom" scores nothing for relative atomic mass. The definition must be the average mass of an atom compared with one-twelfth the mass of a carbon-12 atom, taking isotopic abundance into account. Two parts are required: the comparison to carbon-12, and the weighting by abundance. A vague one-part definition is a common avoidable loss on this topic.
Isotope versus a different element
Changing the number of neutrons gives an isotope of the same element — the chemistry is unchanged. Changing the number of protons (the atomic number) makes a different element. Students who change the protons to "make an isotope", or who treat isotopes as different elements, lose the mark. Between isotopes, only the neutron count differs.
Ar is not a whole number
Do not confuse relative atomic mass with mass number. Mass number is a whole number for a single atom (protons + neutrons). Relative atomic mass is an average over all the isotopes weighted by abundance, so it is usually not a whole number (e.g. chlorine is 35.5). Divide by the total abundance, never by the number of isotopes.
Sanity-check Ar against the data sheet
After an Ar calculation, check the answer sits near the data-sheet value (about 35.5 for chlorine, about 80 for bromine). If not, you probably divided by the wrong number. Still give the calculated value as your answer — the data-sheet number is only a check.
Work Ar in four steps, round at the end
For any Ar calculation: multiply each isotope mass by its % abundance, add the products, divide by the total abundance (usually 100), and round only at the end. Keep full figures through the intermediate steps; truncating early loses the accuracy mark.
Read a nuclide safely
When counting particles, fix the element from the atomic number (smaller number = protons). The larger number is the mass number, so neutrons = mass number − atomic number, and in a neutral atom electrons = protons. Doing these in order stops the protons/mass-number mix-up.
Sub-atomic particles:
| Particle | Location | Relative mass | Relative charge |
|---|---|---|---|
| Proton | Nucleus | 1 | |
| Neutron | Nucleus | 1 | |
| Electron | Shells around the nucleus | (negligible) |
Useful relationships for a neutral atom:
The nucleus contains almost all of the atom's mass (protons + neutrons); the electrons contribute a negligible mass but occupy the space around the nucleus.
Define the term atom.
An atom of sodium has an atomic number of 11 and a mass number of 23.
State the number of protons, neutrons and electrons in this atom.
