Atomic structure and the periodic table: atomic structure and subatomic particles

What you'll learn

This revision guide covers the fundamental building blocks of chemistry: atoms and their constituent particles. You'll learn about the structure of atoms, the properties of subatomic particles, and how to calculate atomic mass and use atomic notation. This topic forms the foundation for understanding chemical bonding, the periodic table, and reactivity patterns tested throughout your AQA GCSE Chemistry course.

Key terms and definitions

Atom — the smallest particle of an element that retains the chemical properties of that element; atoms cannot be broken down by chemical means.

Subatomic particles — particles smaller than an atom (protons, neutrons and electrons) that make up the structure of all atoms.

Mass number — the total number of protons and neutrons in the nucleus of an atom (also called nucleon number).

Atomic number — the number of protons in the nucleus of an atom; this determines which element the atom is and its position in the periodic table.

Isotopes — atoms of the same element (same number of protons) that have different numbers of neutrons, resulting in different mass numbers.

Relative atomic mass (Ar) — the average mass of all the isotopes of an element, taking into account their abundance, compared to 1/12th the mass of a carbon-12 atom.

Ion — a charged particle formed when an atom gains or loses electrons; positive ions (cations) have lost electrons, negative ions (anions) have gained electrons.

Electron configuration — the arrangement of electrons in shells (energy levels) around the nucleus of an atom.

Core concepts

Structure of the atom

Atoms consist of a central nucleus containing protons and neutrons, surrounded by electrons arranged in shells (energy levels).

The nucleus:

• Contains protons (positively charged) and neutrons (no charge)
• Is extremely small compared to the overall size of the atom
• Contains almost all the atom's mass
• Has a radius of approximately 1 × 10⁻¹⁴ metres

The electron shells:

• Electrons orbit the nucleus in fixed energy levels called shells
• The overall radius of an atom is approximately 1 × 10⁻¹⁰ metres
• This means the radius of an atom is about 10,000 times larger than the nucleus
• Most of an atom is empty space

Properties of subatomic particles

You must know the relative mass and relative charge of the three subatomic particles:

Key points:

• Protons and neutrons have approximately the same mass
• Electrons have negligible mass compared to protons and neutrons
• Atoms are neutral overall because they contain equal numbers of protons and electrons
• The positive charges (protons) balance the negative charges (electrons)

Atomic number and mass number

Every element is defined by its atomic number (number of protons). The mass number tells you the total number of particles in the nucleus.

Atomic number (Z):

• Number of protons in the nucleus
• Defines which element it is
• In a neutral atom, also equals the number of electrons
• Never changes for a given element

Mass number (A):

• Total number of protons + neutrons
• Can vary for the same element (isotopes)
• Always a whole number

Calculating the number of neutrons: Number of neutrons = mass number - atomic number

Standard atomic notation

Atoms are represented using standard notation:

   A
   X
   Z

Where:

• X = chemical symbol
• A = mass number (top)
• Z = atomic number (bottom)

Example: ²³₁₁Na represents a sodium atom with:

• Atomic number 11 (11 protons)
• Mass number 23 (23 particles in nucleus)
• Therefore 23 - 11 = 12 neutrons
• 11 electrons (in a neutral atom)

Isotopes

Isotopes are atoms of the same element with different mass numbers. They have the same number of protons but different numbers of neutrons.

Properties of isotopes:

• Same atomic number (same element)
• Different mass numbers
• Same chemical properties (because they have the same number of electrons and electron arrangement)
• Slightly different physical properties (such as density) due to different masses

Common examples you should know:

Carbon isotopes:

• Carbon-12 (¹²₆C): 6 protons, 6 neutrons, 6 electrons
• Carbon-13 (¹³₆C): 6 protons, 7 neutrons, 6 electrons
• Carbon-14 (¹⁴₆C): 6 protons, 8 neutrons, 6 electrons

Chlorine isotopes:

• Chlorine-35 (³⁵₁₇Cl): 75% abundance
• Chlorine-37 (³⁷₁₇Cl): 25% abundance

Relative atomic mass

Most elements exist as a mixture of isotopes. The relative atomic mass (Ar) is the weighted average mass of all isotopes, taking their abundance into account.

Formula for calculating relative atomic mass:

Ar = Σ (isotope mass × % abundance) / 100

Or:

Ar = [(mass₁ × abundance₁) + (mass₂ × abundance₂) + ...] / total abundance

Key points:

• Relative atomic mass values are rarely whole numbers (except for elements with only one naturally occurring isotope)
• Values are found on the periodic table
• Ar is relative to carbon-12, which is defined as exactly 12

Electron arrangement

Electrons occupy shells (energy levels) around the nucleus. Each shell can hold a maximum number of electrons.

Rules for electron arrangement:

• First shell (closest to nucleus): maximum 2 electrons
• Second shell: maximum 8 electrons
• Third shell: maximum 8 electrons (at GCSE level)
• Electrons fill the lowest energy shells first
• Shells fill in order: 1st → 2nd → 3rd

How to work out electron configuration:

1. Find the atomic number (number of electrons in a neutral atom)
2. Fill shells in order until all electrons are allocated
3. Write as numbers separated by commas

Examples:

• Sodium (Na, atomic number 11): 2, 8, 1
• Chlorine (Cl, atomic number 17): 2, 8, 7
• Calcium (Ca, atomic number 20): 2, 8, 8, 2

Electronic configuration and the periodic table:

• Elements in the same group have the same number of electrons in their outer shell
• The group number equals the number of outer shell electrons (for Groups 1-7)
• Group 0 (noble gases) have full outer shells
• The period number equals the number of shells occupied

Ions and electron loss/gain

Atoms form ions by losing or gaining electrons to achieve a stable electron configuration (usually a full outer shell).

Metal atoms (Groups 1, 2, 3):

• Lose electrons to form positive ions (cations)
• Example: Na → Na⁺ + e⁻ (sodium loses 1 electron)
• Electronic configuration of Na⁺: 2, 8 (same as neon)

Non-metal atoms (Groups 5, 6, 7):

• Gain electrons to form negative ions (anions)
• Example: Cl + e⁻ → Cl⁻ (chlorine gains 1 electron)
• Electronic configuration of Cl⁻: 2, 8, 8 (same as argon)

Important:

• The number of protons never changes when ions form
• Only electrons are lost or gained
• Ions have different electron configurations from their parent atoms

Worked examples

Example 1: Identifying subatomic particles

Question: An atom of aluminium can be represented as ²⁷₁₃Al. State the number of protons, neutrons and electrons in this atom. [3 marks]

Solution:

• Protons: 13 (the atomic number at the bottom) [1 mark]
• Neutrons: 27 - 13 = 14 (mass number minus atomic number) [1 mark]
• Electrons: 13 (equal to number of protons in a neutral atom) [1 mark]

Mark scheme guidance: Award 1 mark for each correct value. Students must show working for the neutron calculation or state the method.

Example 2: Calculating relative atomic mass

Question: Copper exists as two isotopes: copper-63 (⁶³Cu) with 69% abundance and copper-65 (⁶⁵Cu) with 31% abundance. Calculate the relative atomic mass of copper. Give your answer to 1 decimal place. [3 marks]

Solution:

Step 1: Write out the formula [1 mark] Ar = [(mass₁ × abundance₁) + (mass₂ × abundance₂)] / 100

Step 2: Substitute values [1 mark] Ar = [(63 × 69) + (65 × 31)] / 100 Ar = [4347 + 2015] / 100 Ar = 6362 / 100

Step 3: Calculate final answer [1 mark] Ar = 63.6 (to 1 d.p.)

Mark scheme guidance: Award marks for correct method, correct substitution, and correct final answer to the specified number of decimal places.

Example 3: Electron configuration and ions

Question: (a) Give the electronic configuration of a sulfur atom (atomic number 16). [1 mark] (b) Sulfur forms an ion with a 2- charge. Give the electronic configuration of this ion and explain how it is formed. [3 marks]

Solution:

(a) Electronic configuration of sulfur: 2, 8, 6 [1 mark]

(b) Electronic configuration of S²⁻: 2, 8, 8 [1 mark]

Explanation: The sulfur atom gains 2 electrons [1 mark] to achieve a full outer shell / stable electronic configuration [1 mark]

Mark scheme guidance: For part (b), students must state the correct configuration AND explain that electrons are gained (not "added" or vague language) with a reason.

Common mistakes and how to avoid them

• Confusing mass number and atomic number — Remember: atomic number is always smaller and goes at the bottom in standard notation. Mass number = protons + neutrons, so it's always larger (or equal for hydrogen-1).

• Thinking electrons have a mass of 0 — Electrons do have mass, it's just very small (approximately 1/1836 of a proton). For calculations, you can ignore electron mass, but don't say it's zero in descriptive answers.

• Calculating relative atomic mass incorrectly — Always divide by 100 (or total percentage) after multiplying each isotope mass by its abundance. Students often forget this final division step.

• Saying isotopes have different chemical properties — Isotopes have the same chemical properties because they have the same number of electrons and same electron arrangement. Only their physical properties (like density) differ slightly.

• Incorrect electron configurations for ions — When metals form positive ions, electrons are removed from the outer shell. When non-metals form negative ions, electrons are added to the outer shell. The number of protons never changes.

• Writing electron configurations that exceed shell capacity — Remember the maximum electrons per shell: 2 in the first, 8 in the second, 8 in the third (at GCSE). Never write something like 2, 9, 6 — this violates the rules.

Exam technique for "Atomic structure and the periodic table: atomic structure and subatomic particles"

• Command word "State" — Give a brief, specific answer without explanation. For example, if asked to "state the number of neutrons," just give the number and perhaps show your calculation (mass number - atomic number). No extra detail is needed or rewarded.

• Calculations must show working — For relative atomic mass calculations or determining numbers of subatomic particles, always show your method. Even if your final answer is wrong, you can gain marks for correct working. Write out formulas before substituting numbers.

• Use precise terminology — Say "electron shells" or "energy levels," not "orbits" or "rings." Use "atomic number" rather than "proton number" in formal answers. Say atoms "gain" or "lose" electrons, not "get" or "take."

• Standard form and units — When stating the size of atoms or nuclei, use standard form correctly (e.g., 1 × 10⁻¹⁰ m, not just "10⁻¹⁰"). Include units where required, particularly in measurement questions.

Quick revision summary

Atoms consist of a nucleus (containing protons and neutrons) surrounded by electrons in shells. The atomic number defines the element and equals the number of protons. Mass number equals protons plus neutrons. Isotopes are atoms of the same element with different numbers of neutrons. Relative atomic mass is the weighted average of isotope masses. Electrons fill shells (2, 8, 8) from the innermost outward. Atoms form ions by losing or gaining electrons to achieve stable electron configurations. Metal atoms lose electrons forming positive ions; non-metal atoms gain electrons forming negative ions.