Atomic Number, Mass Number, and Notation

High School⏱ 8 min readAtomic Structure

Atomic number defines the element, mass number counts nucleons, and nuclear notation packs both into one symbol.

Quick answer

Atomic number (Z) is the number of protons — it defines the element. Mass number (A) is protons + neutrons in a specific atom. Neutrons = A − Z. Nuclear notation ᴬZX packs both numbers around the element symbol.

Example: ¹⁴₆C: 6 protons, 14 − 6 = 8 neutrons, and (neutral) 6 electrons — carbon-14.

What you will learn

Define atomic number (Z) and mass number (A) and compute neutron counts.
Write and interpret nuclear notation (ᴬZX) including charge.
Determine particle counts for neutral atoms and ions.
Explain why periodic-table atomic weights are weighted averages, not whole numbers.

Atomic number: the element's identity

The atomic number, symbol $Z$ , is the number of protons in an atom's nucleus. It is the element's fingerprint: every carbon atom has $Z = 6$ , every oxygen atom $Z = 8$ . The periodic table is ordered by atomic number, and no chemical process can change it — altering $Z$ means transmuting one element into another, which only nuclear reactions do.

Because a neutral atom has no net charge, its electron count equals its proton count. A neutral carbon atom ( $Z = 6$ ) holds six electrons. This equality breaks for ions: a Ca²⁺ ion still has $Z = 20$ protons but only 18 electrons — the proton count defines it as calcium regardless of charge.

Atomic number also fixes an element's address on the periodic table and, through electron count, its entire chemistry. Everything downstream — bonding, reactivity, periodic trends — traces back to $Z$ .

Mass number and nuclear notation

The mass number, symbol $A$ , counts the nucleons: protons plus neutrons in one specific atom. Neutrons follow immediately:

$\text{neutrons} = A - Z$

For $^{23}_{11}Na$ : $Z = 11$ protons, $A = 23$ , so neutrons $= 23 - 11 = 12$ ; neutral sodium has 11 electrons. For the ion $^{23}_{11}Na^+$ : still 11 protons and 12 neutrons, but only 10 electrons.

Nuclear notation writes this compactly as $^{A}_{Z}X$ , with the mass number top-left and atomic number bottom-left of the element symbol. The atomic number is technically redundant (the symbol already implies it) but is included for clarity in nuclear contexts. Charge, when present, goes top-right: $^{16}_{8}O^{2-}$ is the oxide ion — 8 protons, 8 neutrons, 10 electrons.

\text{neutrons} = A - Z

Neutron count from mass and atomic numbers

Why atomic weights are not whole numbers

Here the macroscopic and particulate levels meet. The periodic table lists chlorine's atomic weight as 35.45 — not a whole number. That is because natural chlorine is a mixture: about 76% $^{35}Cl$ and 24% $^{37}Cl$ . The table value is the weighted average over isotopes:

$0.76(35) + 0.24(37) \approx 35.5$

Mass number describes one atom (always whole); atomic weight describes the natural mixture (rarely whole). This distinction resolves endless confusion: the whole numbers belong to individual isotopes, the decimals to bulk samples. When you weigh chlorine in the lab, you weigh the mixture — so the mixture's average is what you use.

Worked examples in both directions

Given $^{31}_{15}P^{3-}$ (phosphide ion): protons $= Z = 15$ ; neutrons $= 31 - 15 = 16$ ; electrons $= 15 + 3 = 18$ (the 3− charge means three extra electrons). The element is phosphorus ( $Z = 15$ ), the isotope is phosphorus-31, and the charge makes it an anion.

Reverse the direction: an atom has 12 protons and 14 neutrons. Then $Z = 12$ (magnesium), $A = 26$ , notation $^{26}_{12}Mg$ . If it loses two electrons it becomes $^{26}_{12}Mg^{2+}$ with 10 electrons — still magnesium, still mass number 26, because neither $Z$ nor $A$ changed.

Fluency in both directions — notation to counts, counts to notation — is the grammar of everything that follows in atomic structure.

Common mistakes

✗ The atomic number can change in a chemical reaction.

Chemical reactions only rearrange electrons; the nucleus is untouched, so proton count never changes. Only nuclear reactions — decay, fission, fusion — alter the atomic number, and those transmute the element.

✗ Mass number and atomic mass are the same thing.

Mass number (A) is a whole-number count of nucleons in one specific isotope. Atomic mass (atomic weight) is the weighted average over all natural isotopes, usually not a whole number. A = 35 for one chlorine atom; atomic weight 35.45 for chlorine in nature.

✗ Electrons are included in the mass number.

Electrons are far too light (~1/1836 of a proton) to matter for mass counts. Mass number counts only protons + neutrons. Electron mass is simply neglected in these whole-number bookkeeping sums.

Real-world applications

Identifying elements from particle counts in mass spectrometry
Writing nuclear notation for isotopes used in medicine and dating
Predicting ion charges from an element's position and electron count
Interpreting periodic table entries, which are organized by atomic number

Frequently asked questions

What is the difference between atomic number and mass number?

Atomic number (Z) is the proton count — it defines the element and never changes in chemistry. Mass number (A) is protons + neutrons — it defines the specific isotope. Carbon always has Z = 6, but A can be 12, 13, or 14 depending on neutrons.

Does the electron count always equal the atomic number?

In a neutral atom, electrons equal protons (charges balance). Ions are the exception: cations have fewer electrons than protons, anions have more. So for neutral atoms, electron count = Z; for ions, adjust by the charge.

How do I read nuclear notation like ¹⁴₆C?

Top-left is the mass number (A), bottom-left is the atomic number (Z): ᴬZX. For carbon-14: ¹⁴₆C. Some tables write the atomic number at bottom-left and the charge at top-right (¹⁴₆C has no charge shown because it is neutral).

Why aren't atomic weights on the periodic table whole numbers?

Because most elements are mixtures of isotopes with different masses. Chlorine is ~76% Cl-35 and ~24% Cl-37, so its average is 35.45 — closer to 35 but not a whole number. Only elements with a single dominant isotope have near-integer atomic weights.