Introduction: Why start with atoms?

Atoms are the building blocks of everything from the air we breathe to the cells that make us. Understanding atoms and elements isn’t just a chemistry requirement; it’s a gateway to grasping how matter behaves, bonds, and transforms. This lecture introduces the key concepts that underpin all chemical understanding: what elements are, how atoms are structured, and how the periodic table organises this knowledge into a coherent, predictive system.

Whether you’re revising for exams, teaching others, or building outreach materials, this guide offers a modular, remixable foundation with live links to deepen your study and spark curiosity.

1. What is an element?

An element is a pure substance that cannot be broken down into simpler substances by chemical means. Each element is defined by the type of atom it contains, and every atom of a given element has the same number of protons in its nucleus.

• Elements exist naturally as solids, liquids, or gases at room temperature.
• There are 118 known elements:
  • 94 occur naturally (e.g. oxygen, iron, gold)
  • 24 are synthetic (e.g. technetium, einsteinium)
• Each element is assigned a symbol, often derived from its Latin name:

• Oxygen → O
• Carbon → C
• Aluminium → Al

Explore the full list interactively at Ptable.com, a dynamic periodic table with live data and visualisations.

2. What is an atom?

An atom is the smallest unit of an element that retains its chemical properties. Atoms are made of three subatomic particles:

• Protons are positively charged, found in the nucleus
• Neutrons are neutral, also in the nucleus
• Electrons are negatively charged, orbiting the nucleus in shells

Each element is made of one type of atom, and the number of protons in the nucleus determines the element’s identity. For example:

• Carbon atoms always have 6 protons
• Oxygen atoms always have 8 protons

Learn more about atomic structure at the Khan Academy’s atomic model tutorial.

3. Atomic mass unit (amu)

Atoms are incredibly small, so chemists use a relative scale to compare their masses. This scale is based on the carbon-12 isotope, which is assigned a mass of exactly 12 atomic mass units (amu).

• 1 amu = 1/12 the mass of a carbon-12 atom
• Hydrogen has a mass of ~1 amu
• Oxygen has a mass of ~16 amu

This system allows chemists to compare atomic masses without needing to weigh individual atoms.

Explore atomic masses and isotopes at the Royal Society of Chemistry’s periodic table.

4. Atomic number and mass number

Each atom is defined by two key numbers:

• Atomic number (Z): number of protons in the nucleus
• Mass number (A): total number of protons and neutrons

For example, chlorine has:

• Atomic number = 17
• Mass number = 35
• Neutrons = 35 − 17 = 18

This notation is often written as:

This is called nuclide notation, and it helps identify isotopes of the same element with different numbers of neutrons.

Learn more about isotopes and nuclide notation at Chemguide’s isotope page.

5. Electron shells and orbitals

Electrons orbit the nucleus in energy levels or shells. These shells have specific capacities:

• 1st shell: up to 2 electrons
• 2nd shell: up to 8 electrons
• 3rd shell: up to 18 electrons
• 4th shell: up to 32 electrons

Electrons fill the lowest available energy level first, starting closest to the nucleus. This arrangement determines the atom’s chemical behaviour.

For example:

• Hydrogen (1 electron): 1
• Helium (2 electrons): 2
• Lithium (3 electrons): 2.1
• Neon (10 electrons): 2.8
• Sodium (11 electrons): 2.8.1

Explore electron configurations interactively at ChemCollective’s periodic table.

6. Orbitals and subshells

Each shell contains orbital regions where electrons are likely to be found. Orbitals are grouped into subshells:

• s orbital: 1 type, holds 2 electrons
• p orbitals: 3 types, hold 6 electrons
• d orbitals: 5 types, hold 10 electrons
• f orbitals: 7 types, hold 14 electrons

Orbitals fill in a specific order based on energy:

This is known as the Aufbau principle. Electrons also obey the Pauli exclusion principle and Hund’s rule.

Learn orbital shapes and filling order at ChemLibreTexts.

7. Spectroscopic notation

Electron configurations can be written using spectroscopic notation, which shows how electrons fill orbitals:

• Hydrogen: 1s¹
• Helium: 1s²
• Lithium: 1s² 2s¹
• Neon: 1s² 2s² 2p⁶
• Sodium: 1s² 2s² 2p⁶ 3s¹
• Sulphur: 1s² 2s² 2p⁶ 3s² 3p⁴

This notation helps predict chemical reactivity and bonding.

Practice writing configurations at Chemistry LibreTexts’ configuration tutorial.

8. The periodic table

The periodic table organises elements by:

• Atomic number (left to right)
• Electron configuration
• Recurring chemical properties

Periods (rows)

• Indicate the number of electron shells
• Atomic number increases left to right
• Atomic size decreases across a period

Groups (columns)

• Elements in the same group have similar outer electron configurations
• Group number = number of valence electrons
• Atomic size increases down a group

Explore trends interactively at Ptable’s trend visualizer.

9. Metals, non-metals, and metalloids

The periodic table divides elements into:

• Metals are good conductors, malleable, and form positive ions
• Non-metals are poor conductors, brittle, form negative ions or share electrons
• Metalloids intermediate properties (e.g. silicon, arsenic)

Metals dominate the left and centre of the table; non-metals are on the right.

Compare properties at the Royal Society of Chemistry’s element explorer.

10. Atomic size trends

Across a period

• Atomic number increases → more protons
• Electrons pulled closer → atomic radius decreases

Down a group

• Energy levels added → electrons farther from the nucleus
• Atomic radius increases

Visualise atomic radii at WebElements’ atomic size chart.

Available in PDF format with an associated quiz

Closing: What atoms unlock

Atoms are more than particles; they’re the grammar of matter. By understanding atomic structure, electron arrangement, and periodic trends, we gain the ability to predict chemical behaviour, explain bonding, and design new materials. This lecture lays the groundwork for everything that follows in chemistry, from reactions and compounds to spectroscopy and sensors.

Whether you’re building outreach kits, revising for exams, or curating legacy-driven resources, this atomic foundation is your launchpad.

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