CHEMISTRY · LESSON 01 OF 12

Atomic structure

Atoms are made of protons, neutrons and electrons. Z (protons) fixes the element; A counts protons plus neutrons; the charge counts missing or extra electrons.

What this lesson explains

Every chemical property begins with atomic structure. The number of protons decides which element you have; the electrons decide how it bonds and reacts; the neutrons affect mass and nuclear stability. Isotopic masses give the atomic masses used in every mole calculation in chemistry and chemical engineering.

Before you begin

Elements, compounds and mixtures

A pure substance has a fixed composition: an element contains one kind of atom (O₂, Fe); a compound contains two or more elements chemically joined in a fixed ratio (H₂O, NaCl). A mixture combines substances without fixing their ratio; it can be uniform (air, salt water) or not (sand in water) and is separated physically. A physical change (melting, dissolving) keeps the substances the same; a chemical change (burning, rusting) makes new substances. A uniform appearance does not prove a substance is pure: salt water looks like water.

Charges and signs

A proton has charge +1 (in units of e = 1.602 × 10⁻¹⁹ C), an electron −1. Net charge = (number of protons) − (number of electrons).

Weighted average

A weighted average multiplies each value by its fraction (abundance/100) and adds: Σ(fraction × value). The fractions must add to 1.

The idea, made visible

An atom has a tiny, dense nucleus containing positively charged protons and uncharged neutrons, surrounded by negatively charged electrons. Almost all the mass is in the nucleus; almost all the volume is the space occupied by electrons. A nucleus is about 10⁻¹⁵ m across, an atom about 10⁻¹⁰ m — a factor of 100 000.

The atomic number Z is the number of protons. It defines the element: every atom with 6 protons is carbon. The mass number A is the number of protons plus neutrons. A neutral atom has as many electrons as protons.

Reading ²⁷₁₃Al³⁺: 13 protons (so aluminium), 14 neutrons and, because the charge is 3+, three fewer electrons than protons: 10.

Isotopes are atoms of the same element (same Z) with different numbers of neutrons (different A), for example carbon-12 and carbon-14. Isotopes have almost identical chemistry because chemistry is governed by electrons.

An ion has gained or lost electrons; the number of protons never changes in a chemical reaction. A cation (positive) has lost electrons: Mg²⁺ has 12 protons and 10 electrons. An anion (negative) has gained electrons: Cl⁻ has 17 protons and 18 electrons.

The atomic mass printed on the periodic table is a weighted average of the masses of the naturally occurring isotopes, weighted by their abundances. That is why chlorine’s atomic mass is 35.45 u even though no chlorine atom has that mass.

Counting particles
species	Z (protons)	A	neutrons	electrons
²³Na	11	23	12	11
²³Na⁺	11	23	12	10
³⁵Cl⁻	17	35	18	18
²⁴Mg²⁺	12	24	12	10
¹⁶O²⁻	8	16	8	10
¹⁴C	6	14	8	6

Na⁺, Mg²⁺ and O²⁻ all have 10 electrons: they are isoelectronic with neon, but they are different elements because their proton numbers differ.

Key terms

Atomic number Z: Number of protons in the nucleus; identifies the element.
Mass number A: Number of protons + neutrons (a whole number). Neutrons n = A − Z.
Isotopes: Atoms with the same Z but different A (different numbers of neutrons).
Ion: An atom or group of atoms with a net charge due to gained or lost electrons. Cation: positive; anion: negative.
Atomic mass unit (u): 1/12 of the mass of a carbon-12 atom, 1.6605 × 10⁻²⁷ kg. Proton and neutron masses are each about 1 u; an electron is about 1/1836 u.
Average atomic mass: Abundance-weighted average of isotope masses, as printed on the periodic table.

The formulas and what they mean

Symbols, meanings and units
Symbol	Meaning	Unit
Z	atomic number (protons)	—
A	mass number (protons + neutrons)	—
n	number of neutrons	—
^A_ZX^q	nuclide symbol: element X with mass number A, atomic number Z, charge q	—
u	atomic mass unit	1.6605 × 10⁻²⁷ kg

Neutrons

n = A − Z

Conditions and limits: A is the mass number of a specific isotope, not the average atomic mass.

Electrons in an ion

electrons = Z − q

Conditions and limits: q is the charge with its sign: Mg²⁺ (q = +2) has 12 − 2 = 10; Cl⁻ (q = −1) has 17 + 1 = 18.

Average atomic mass

M_avg = Σ (fractional abundance × isotopic mass)

Conditions and limits: Abundances as fractions that add to 1 (divide percentages by 100).

Why it works: Why chlorine’s atomic mass is 35.45 u

Natural chlorine is 75.78% chlorine-35 (34.969 u) and 24.22% chlorine-37 (36.966 u).

0.7578 × 34.969 = 26.50 u.
Contribution of Cl-35.
0.2422 × 36.966 = 8.953 u.
Contribution of Cl-37.
Sum: 26.50 + 8.953 = 35.45 u.
The weighted average.

The average lies closer to 35 because Cl-35 is about three times as abundant. No individual chlorine atom has mass 35.45 u.

A first worked example

How to count particles

Find Z from the element symbol (periodic table): protons = Z.
Neutrons = A − Z, using the mass number of the stated isotope.
Electrons = Z − charge (subtract a positive charge, add the size of a negative charge).
For average atomic mass, convert percentages to fractions and form Σ(fraction × mass).
Check: a cation has fewer electrons than protons; an anion has more.

Neutrons in an isotope

Problem. An atom has mass number 23 and atomic number 11. How many neutrons does it have, and what element is it?

n = A − Z = 23 − 11 = 12.
Mass number counts protons and neutrons.
Z = 11 is sodium.
The proton number identifies the element.

Result: 12 neutrons; the atom is sodium-23.

What it means: Any atom with 11 protons is sodium, whatever its neutron number.

A different case

Electrons in an ion

Problem. How many electrons are in Mg²⁺ (Z = 12) and in S²⁻ (Z = 16)?

Mg²⁺: 12 − 2 = 10 electrons.
A 2+ charge means two electrons fewer than protons.
S²⁻: 16 + 2 = 18 electrons.
A 2− charge means two extra electrons.

Result: Mg²⁺: 10 electrons; S²⁻: 18 electrons.

What it means: Ionisation never changes the proton number or the element.

More worked cases

Each case below uses a different skill. Every step and result is shown.

Average atomic mass

Problem. Boron is 19.9% ¹⁰B (10.013 u) and 80.1% ¹¹B (11.009 u). Find its average atomic mass.

0.199 × 10.013 + 0.801 × 11.009.
Weighted average with fractional abundances.
= 1.993 + 8.818 = 10.81 u.
Add the contributions.

Result: 10.81 u (as on the periodic table).

What it means: The average is closer to 11 because boron-11 is about four times more abundant.

Abundance from the average mass

Problem. Copper has isotopes ⁶³Cu (62.930 u) and ⁶⁵Cu (64.928 u); its average atomic mass is 63.546 u. Find the abundance of ⁶³Cu.

Let x be the fraction of ⁶³Cu; then 1 − x is ⁶⁵Cu.
Two isotopes whose fractions add to 1.
62.930x + 64.928(1 − x) = 63.546.
Weighted-average equation.
64.928 − 1.998x = 63.546 ⇒ x = 1.382 / 1.998 = 0.692.
Solve for x.

Result: About 69.2% ⁶³Cu (and 30.8% ⁶⁵Cu).

What it means: The accepted natural abundance of ⁶³Cu is 69.15%, which agrees.

Common misunderstandings

Misunderstanding: Using the periodic-table atomic mass as A to find neutrons.
Correct idea: A is a whole number for a specific isotope. Chlorine’s 35.45 is an average, not a mass number.
Misunderstanding: Changing the number of protons when an ion forms.
Correct idea: Only electrons are gained or lost in chemical changes.
Misunderstanding: Adding electrons for a positive charge.
Correct idea: Positive means electrons were lost: electrons = Z − (positive charge).
Misunderstanding: “Isotopes are different elements.”
Correct idea: Isotopes have the same Z, so they are the same element with different masses.

Keep in mind

n = A − ZNeutrons
electrons = Z − qElectrons in an ion
M_avg = Σ (fractional abundance × isotopic mass)Average atomic mass

Scope of this lesson

Nuclear reactions and radioactive decay are not part of Semester 1 Chemistry.
Electron arrangement in atoms is developed in “Quantum theory” and “Periodic table”.

Next: Quantum theory. Light comes in photons with energy E = hν. Electrons in atoms can only have certain energies, so atoms absorb and emit light only at particular wavelengths. Electrons are described by orbitals labelled by quantum numbers.

An older bookmark may point here. That section is now its own lesson: Moles, molar mass and chemical formulas.

Original study text. Sources and credits.

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