Difficulty: Introductory | Prerequisites: Know what elements and the periodic table are (see Docs 1 and 2 in this set).
This topic takes you inside the atom. You learn that every atom is built from three subatomic particles (protons, neutrons, electrons), that the atomic number defines the element, and that the mass number counts the heavy particles in the nucleus. You then meet isotopes (same element, different neutron count) and learn to calculate a weighted average atomic mass from isotopic data. This is the quantitative backbone of everything that follows in stoichiometry and nuclear chemistry.
Atoms consist of protons and neutrons (in the nucleus) plus electrons (outside). The atomic number equals the proton count and defines the element. The mass number equals protons plus neutrons. Isotopes are atoms of the same element with different neutron counts, and the atomic mass on the periodic table is a weighted average of all naturally occurring isotopes.
Atomic number (Z)
The number of protons in the nucleus of an atom. This number defines which element it is. Every atom of iron has 26 protons, no exceptions. In simple terms, the atomic number is the element's identity card.
Mass number (A)
The total count of protons plus neutrons in the nucleus. Unlike atomic number, mass number can vary between atoms of the same element (that is what isotopes are). Think of it as the atom's "weight tag" for a specific version of that element.
Proton
A positively charged particle in the nucleus. Charge = +1. Relative mass = ~1 amu.
Neutron
A neutral particle in the nucleus (no charge). Relative mass = ~1 amu. Neutrons add mass without changing the element's identity.
Electron
A negatively charged particle found outside the nucleus in electron shells. Charge = -1. Mass is negligible compared to protons and neutrons (~1/1836 amu). In a neutral atom, the electron count equals the proton count.
Isotope
Atoms of the same element (same number of protons) that differ in their number of neutrons. For example, Ag-107 and Ag-109 are both silver, both with 47 protons, but one has 60 neutrons and the other has 62. Think of isotopes as siblings: same family (element), different weight.
Atomic mass (atomic weight)
The weighted average of the masses of all naturally occurring isotopes of an element, measured in atomic mass units (amu). This is the decimal number shown on the periodic table. It is not the same as the mass number of any single isotope.
Atomic mass unit (amu)
The unit used to express atomic and isotopic masses. Defined as one-twelfth the mass of a carbon-12 atom. One proton or one neutron is approximately 1 amu.
Nuclear symbol
The notation that shows an element's symbol with its mass number as a superscript and its atomic number as a subscript, both written to the left. For example, the silver-107 isotope is written with 107 on top and 47 on the bottom, to the left of Ag.
Protons (+ charge, ~1 amu) and neutrons (no charge, ~1 amu) sit in the nucleus.
Electrons (- charge, negligible mass) orbit the nucleus in shells.
In a neutral atom, the number of electrons equals the number of protons.
Atomic number (Z) = number of protons. This never changes for a given element.
Mass number (A) = protons + neutrons. This can differ between isotopes.
Number of neutrons = A - Z.
Iron (Fe): Z = 26, A = 56, so neutrons = 56 - 26 = 30, electrons = 26.
Aluminium (Al): Z = 13, A = 27, so neutrons = 27 - 13 = 14, electrons = 13.
Potassium (K): Z = 19, A = 39, so neutrons = 39 - 19 = 20, electrons = 19.
Bromine (Br): Z = 35, A = 80, so neutrons = 80 - 35 = 45, electrons = 35.
Gold (Au): Z = 79, A = 197, so neutrons = 197 - 79 = 118, electrons = 79.
Isotopes share the same atomic number but have different mass numbers.
Example: silver has two main isotopes, Ag-107 (47 protons, 60 neutrons) and Ag-109 (47 protons, 62 neutrons). Both are silver; only the neutron count differs.
The nuclear symbol places the mass number as a superscript and the atomic number as a subscript to the left of the element symbol.
The atomic mass on the periodic table is not the mass of any single isotope. It is a weighted average based on each isotope's mass and its natural abundance.
Method: multiply each isotope's mass by its fractional abundance (percent divided by 100), then add the results.
Silver example from Lab 6:
Ag-107: 106.9 amu, 51.84% abundance. Contribution: 106.9 x 0.5184 = 55.42 amu.
Ag-109: 108.9 amu, 48.15% abundance. Contribution: 108.9 x 0.4815 = 52.44 amu.
Weighted average: 55.42 + 52.44 = 107.9 amu (close to the periodic table value of 107.87 amu).
If told "a neutral atom has mass number 80 and 45 neutrons," find the atomic number: Z = 80 - 45 = 35. Element 35 is bromine (Br).
If that atom gains two protons and two electrons, Z becomes 37, which is rubidium (Rb).
Number of neutrons:
Neutrons = Mass number (A) - Atomic number (Z)
Weighted average atomic mass:
Atomic mass = (mass of isotope 1 x fractional abundance 1) + (mass of isotope 2 x fractional abundance 2) + ...
Where fractional abundance = percent abundance / 100.
Students often confuse atomic number with mass number. The atomic number is always the smaller of the two (it counts only protons), while the mass number counts protons and neutrons together.
A very common error is thinking that the atomic mass on the periodic table is the mass number of a single isotope. It is not. It is a weighted average across all naturally occurring isotopes, which is why it is almost always a decimal.
Students sometimes assume that isotopes of the same element have different chemical properties. They do not, at least not in any way that matters at this level. Chemical behaviour is governed by electrons (and therefore protons), not neutrons.
When calculating neutrons, some students subtract the atomic mass (the decimal from the periodic table) instead of the mass number (the whole number for a specific isotope). Use the mass number, which is always a whole number.
⚠️ Filling in a table of element name, symbol, Z, A, protons, neutrons, and electrons is one of the most common exam formats for this topic. Practise until it is automatic.
⚠️ Calculating weighted average atomic mass from isotopic data is a near-certain exam question. Show your working: multiply each isotope's mass by its fractional abundance, then sum.
⚠️ "Write the atomic symbol" questions give you a mass number and a neutron count (or similar pair of clues) and expect you to deduce the element. Z = A - neutrons, then look up the element.
⚠️ Remember that in a neutral atom, protons = electrons. If the question specifies an ion, that changes, but Lab 6 deals only with neutral atoms.
Isotopes have direct practical uses. Carbon-14 dating relies on the predictable decay of a carbon isotope to estimate the age of archaeological artefacts. Medical imaging (PET scans) uses positron-emitting isotopes injected into the body. Uranium enrichment for nuclear power is essentially sorting U-235 from U-238, two isotopes of the same element with very different nuclear properties.
True or false: The atomic number of an element can change.
False. The atomic number defines the element and never changes.
Fill in the blank: Neutrons = ______ minus ______.
Mass number minus atomic number.
True or false: The atomic mass on the periodic table is always a whole number.
False. It is a weighted average of isotopic masses and is almost always a decimal.
Fill in the blank: In a neutral atom, the number of electrons equals the number of ______.
Protons.
True or false: Two isotopes of the same element have different numbers of protons.
False. They have the same number of protons but different numbers of neutrons.
Q: An element has an atomic number of 19 and a mass number of 39. How many protons, neutrons, and electrons does a neutral atom have?
A: 19 protons, 20 neutrons (39 - 19), and 19 electrons.
Q: Chlorine has two isotopes: Cl-35 (mass 34.97 amu, 75.77% abundance) and Cl-37 (mass 36.97 amu, 24.23% abundance). Calculate the atomic mass.
A: (34.97 x 0.7577) + (36.97 x 0.2423) = 26.50 + 8.96 = 35.46 amu.
Q: A neutral atom has 80 nucleons and 45 neutrons. What element is it?
A: Protons = 80 - 45 = 35. Element 35 is bromine (Br).
Q: How do Ag-107 and Ag-109 differ?
A: Both have 47 protons and 47 electrons. Ag-107 has 60 neutrons; Ag-109 has 62. They differ only in neutron count.
Q: Why is the atomic mass of silver listed as 107.87 amu rather than a whole number?
A: Because 107.87 amu is the weighted average of the masses of all naturally occurring silver isotopes, reflecting that Ag-107 is slightly more abundant than Ag-109.
This material builds directly on the periodic table topic (Doc 2): atomic number is what places each element in its position on the table. It also connects forward to electron configuration (electrons fill shells in a pattern that mirrors the table's periods and groups) and to nuclear chemistry (isotope stability, radioactive decay, half-lives).
Stoichiometry relies on atomic mass for mole calculations, so getting comfortable with weighted averages now saves time later.
atomic structure, subatomic particles, proton, neutron, electron, atomic number, mass number, isotope, isotopic mass, atomic mass, atomic weight, amu, nuclear symbol, isotope notation, weighted average, percent abundance, natural abundance, Lab 6 chemistry, atoms and elements, MSU chemistry, Ag-107, Ag-109