What you'll learn
This topic forms the foundation of GCSE Chemistry. You'll explore how atoms are constructed, how elements are organised in the periodic table, and why elements in the same group share similar properties. Understanding atomic structure is essential for explaining chemical reactions, bonding, and the behaviour of materials.
Key terms and definitions
Atom — the smallest particle of an element that retains its chemical properties, consisting of a nucleus containing protons and neutrons, surrounded by electrons in shells.
Element — a pure substance made of only one type of atom, which cannot be broken down into simpler substances by chemical means.
Isotope — atoms of the same element with the same number of protons but different numbers of neutrons, resulting in different mass numbers.
Mass number — the total number of protons and neutrons in the nucleus of an atom.
Atomic number — the number of protons in the nucleus of an atom, which defines the element and its position in the periodic table.
Electron configuration — the arrangement of electrons in shells (energy levels) around the nucleus of an atom.
Group — a vertical column in the periodic table containing elements with the same number of electrons in their outer shell, resulting in similar chemical properties.
Period — a horizontal row in the periodic table, with elements arranged in order of increasing atomic number.
Core concepts
Structure of the atom
Atoms contain three types of subatomic particles:
Protons:
- Found in the nucleus
- Relative mass of 1
- Relative charge of +1
- Number of protons defines the element
Neutrons:
- Found in the nucleus
- Relative mass of 1
- No charge (neutral)
- Number varies in isotopes
Electrons:
- Orbit the nucleus in shells (energy levels)
- Relative mass of 1/2000 (negligible)
- Relative charge of -1
- Number equals protons in a neutral atom
The nucleus is extremely small compared to the overall atom size, but contains almost all the atom's mass. Atoms have no overall charge because the number of protons equals the number of electrons.
Atomic number and mass number
Every element has a unique atomic number (Z), which equals the number of protons. This determines the element's identity and its position in the periodic table.
The mass number (A) is the sum of protons and neutrons. These are shown in standard notation:
Mass number → ᴬX Atomic number → ᶻ
For example, carbon-12 is written as ¹²₆C, meaning 6 protons and 6 neutrons (12 - 6 = 6).
Calculating subatomic particles:
- Number of protons = atomic number
- Number of electrons = atomic number (in neutral atoms)
- Number of neutrons = mass number - atomic number
Isotopes
Isotopes are different forms of the same element. They have:
- The same atomic number (same number of protons)
- Different mass numbers (different numbers of neutrons)
- Identical chemical properties (same electron configuration)
- Different physical properties (different masses)
Common examples include:
- Carbon-12 (¹²₆C) and carbon-14 (¹⁴₆C)
- Chlorine-35 (³⁵₁₇Cl) and chlorine-37 (³⁷₁₇Cl)
Relative atomic mass (Aᵣ) is the weighted average mass of all isotopes of an element, taking into account their relative abundances. This explains why Aᵣ values are rarely whole numbers (e.g., chlorine = 35.5).
Electron configuration and shells
Electrons occupy shells (energy levels) around the nucleus. Each shell can hold a maximum number of electrons:
- 1st shell: maximum 2 electrons
- 2nd shell: maximum 8 electrons
- 3rd shell: maximum 8 electrons (for the first 20 elements)
Electrons fill shells from the innermost outward. The electron configuration is written showing the number of electrons in each shell, separated by commas or full stops.
Examples:
- Sodium (11 electrons): 2,8,1
- Chlorine (17 electrons): 2,8,7
- Calcium (20 electrons): 2,8,8,2
The number of electrons in the outer shell determines an element's chemical properties and reactivity.
Development of the periodic table
Early attempts:
- Before 1800: Elements listed by properties but no clear pattern
- Johann Döbereiner (1817): Grouped elements in triads with similar properties
- John Newlands (1864): Law of Octaves — every 8th element had similar properties
- Newlands' pattern broke down after calcium; his work was initially rejected
Dmitri Mendeleev (1869):
- Arranged elements by atomic weight (mass)
- Left gaps for undiscovered elements
- Predicted properties of missing elements (e.g., gallium, germanium)
- Sometimes reversed order to keep elements with similar properties in groups
- His predictions proved accurate when elements were discovered
Modern periodic table:
- Arranged by atomic number (not atomic mass)
- Discovery of protons explained why Mendeleev's reversals were correct
- Elements in the same group have the same number of outer electrons
- Elements in the same period have the same number of electron shells
Modern periodic table structure
The periodic table organises all known elements systematically:
Groups (vertical columns):
- Numbered 1-7 and 0 (or 1-18 in extended notation)
- Elements in the same group have similar chemical properties
- All group members have the same number of outer shell electrons
- Group number = number of outer electrons (for Groups 1-7)
Important groups:
- Group 1: Alkali metals (lithium, sodium, potassium, etc.)
- Group 7: Halogens (fluorine, chlorine, bromine, iodine)
- Group 0/8: Noble gases (helium, neon, argon, etc.)
Periods (horizontal rows):
- Numbered 1-7
- Period number = number of electron shells
- Properties change across a period from metallic to non-metallic
Metals and non-metals:
- Metals: left side and centre of periodic table
- Non-metals: right side of periodic table
- Stepped line separates metals from non-metals
Periodic trends in Groups 1, 7, and 0
Group 1 — Alkali metals:
- One electron in outer shell
- Very reactive metals
- React vigorously with water, producing hydrogen and metal hydroxide
- Reactivity increases down the group (outer electron more easily lost)
- Melting points decrease down the group
- Examples: 2Na + 2H₂O → 2NaOH + H₂
Group 7 — Halogens:
- Seven electrons in outer shell
- Reactive non-metals
- Exist as diatomic molecules (F₂, Cl₂, Br₂, I₂)
- Reactivity decreases down the group (harder to gain an electron)
- Melting and boiling points increase down the group
- More reactive halogens displace less reactive ones from solutions
Group 0 — Noble gases:
- Eight electrons in outer shell (except helium with 2)
- Full outer shells make them unreactive (stable)
- Do not readily form compounds
- Exist as single atoms (monatomic)
- Boiling points increase down the group
- Used where unreactive atmospheres needed (e.g., argon in light bulbs)
Worked examples
Example 1: Calculating subatomic particles
Question: An atom of fluorine can be represented as ¹⁹₉F. State the number of protons, neutrons and electrons in this atom. [3 marks]
Answer:
- Number of protons = 9 (atomic number) ✓
- Number of neutrons = 19 - 9 = 10 (mass number - atomic number) ✓
- Number of electrons = 9 (same as protons in a neutral atom) ✓
Examiner tip: Always show your working for neutrons. Many students forget that mass number includes both protons and neutrons.
Example 2: Electron configuration
Question: Draw the electron configuration for a magnesium atom (atomic number 12) and explain why magnesium is in Group 2 of the periodic table. [3 marks]
Answer: Electron configuration: 2,8,2 ✓
[Drawing would show: nucleus with 12+ inside, first shell with 2 electrons, second shell with 8 electrons, third shell with 2 electrons]
Magnesium is in Group 2 because it has 2 electrons in its outer shell ✓, and the group number equals the number of outer electrons ✓
Examiner tip: For drawing questions, ensure electrons are clearly shown as crosses or dots in distinct shells, not overlapping.
Example 3: Isotopes and relative atomic mass
Question: Chlorine has two main isotopes: chlorine-35 and chlorine-37. Chlorine-35 has an abundance of 75% and chlorine-37 has an abundance of 25%. Calculate the relative atomic mass of chlorine. [3 marks]
Answer: Aᵣ = (35 × 75) + (37 × 25) / 100 ✓
Aᵣ = 2625 + 925 / 100 ✓
Aᵣ = 35.5 ✓
Examiner tip: Always divide by 100 when percentages are given. Show each step of your calculation clearly to secure method marks even if your final answer is incorrect.
Common mistakes and how to avoid them
Confusing mass number with atomic number. Remember: atomic number is always the smaller number and equals protons; mass number is larger and equals protons plus neutrons.
Incorrectly calculating neutrons. Students often add instead of subtract. Always use: neutrons = mass number - atomic number.
Mixing up group and period numbers. Groups run vertically (columns) and determine chemical properties; periods run horizontally (rows) and indicate number of shells.
Forgetting that isotopes have identical chemical properties. Isotopes differ only in neutron number, not electron arrangement, so they react the same way chemically.
Stating that electrons have no mass. Electrons do have mass (1/2000 of a proton), but it's negligible compared to protons and neutrons. Be precise in exams.
Reversing Group 1 and Group 7 reactivity trends. Group 1 reactivity increases down the group; Group 7 reactivity decreases down the group. Link this to electron loss/gain.
Exam technique for "Atomic Structure and the Periodic Table"
Command words matter. "State" requires a simple answer with no explanation (1 mark). "Explain" requires a reason using scientific knowledge (2+ marks). "Calculate" requires working shown for method marks.
Precision in definitions. When defining isotopes, include all key points: same element, same proton number, different neutron number. Missing any component loses marks.
Show calculations clearly. Even if your final answer is wrong, you can gain marks for correct method. Write formulae first, substitute values second, calculate third.
Link structure to properties. Questions often ask you to explain trends or reactivity. Always connect your answer to electron configuration or position in the periodic table — this is what examiners want to see.
Quick revision summary
Atoms contain protons and neutrons in the nucleus, with electrons in shells around it. Atomic number defines the element; mass number equals protons plus neutrons. Isotopes are atoms of the same element with different neutron numbers. Electrons fill shells in a 2,8,8 pattern for the first 20 elements. The periodic table arranges elements by atomic number into groups (same outer electrons, similar properties) and periods (same number of shells). Group 1 metals become more reactive down the group; Group 7 non-metals become less reactive.