Atomic structure and the periodic table

What you'll learn

Atomic structure and the periodic table forms the foundation of CIE IGCSE Chemistry, accounting for approximately 15% of Paper 2 and Paper 4 marks. This topic covers the structure of atoms, arrangement of elements in the periodic table, and how electron configuration determines chemical properties. Questions range from recalling subatomic particle properties to explaining trends across periods and groups.

Key terms and definitions

Atomic number (proton number) — the number of protons in the nucleus of an atom, which defines the element and equals the number of electrons in a neutral atom

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

Isotopes — atoms of the same element with the same number of protons but different numbers of neutrons, resulting in different mass numbers

Electron configuration — the arrangement of electrons in shells (energy levels) around the nucleus, written as a series of numbers (e.g. 2,8,1 for sodium)

Relative atomic mass (Aᵣ) — the weighted average mass of all isotopes of an element compared to 1/12 the mass of a carbon-12 atom

Period — a horizontal row in the periodic table, representing elements with the same number of electron shells

Group — a vertical column in the periodic table, representing elements with the same number of outer shell electrons and similar chemical properties

Valency — the combining power of an element, determined by the number of electrons in the outer shell

Core concepts

Structure of the atom

Atoms consist of a central nucleus containing protons and neutrons, surrounded by electrons arranged in shells. The nucleus contains over 99.9% of the atom's mass but occupies only a tiny fraction of its volume.

Subatomic particles:

• Protons: relative mass 1, relative charge +1, located in the nucleus
• Neutrons: relative mass 1, relative charge 0, located in the nucleus
• Electrons: relative mass 1/1840 (approximately 0), relative charge -1, orbit the nucleus in shells

The atomic number determines an element's identity. For example, all carbon atoms have 6 protons. In a neutral atom, the number of protons equals the number of electrons, resulting in no overall charge. Ions form when atoms lose or gain electrons: a sodium ion (Na⁺) has 11 protons but only 10 electrons.

Calculating neutrons: Number of neutrons = Mass number - Atomic number

For chlorine-37: 37 - 17 = 20 neutrons

Isotopes and relative atomic mass

Isotopes are variants of the same element with identical chemical properties but different physical properties (such as density). Chlorine has two main isotopes: chlorine-35 (75% abundance) and chlorine-37 (25% abundance). Both have 17 protons and 17 electrons, but chlorine-35 has 18 neutrons while chlorine-37 has 20 neutrons.

The relative atomic mass accounts for all naturally occurring isotopes:

Aᵣ of chlorine = (35 × 75) + (37 × 25) / 100 = 35.5

This explains why relative atomic masses in the periodic table are rarely whole numbers. CIE IGCSE Chemistry exam questions frequently test isotope calculations and understanding why chemical properties remain unchanged despite different mass numbers.

Electron configuration and shells

Electrons occupy shells at increasing distances from the nucleus, with each shell holding a maximum number of electrons:

• 1st shell: maximum 2 electrons
• 2nd shell: maximum 8 electrons
• 3rd shell: maximum 8 electrons (for elements 1-20)
• 4th shell: maximum 2 electrons (for calcium only at IGCSE level)

Rules for filling shells:

1. Electrons fill the lowest energy shell first
2. Each shell must be full before electrons enter the next shell
3. The outer shell determines chemical properties

Examples of electron configurations:

• Hydrogen (atomic number 1): 1
• Carbon (atomic number 6): 2,4
• Sodium (atomic number 11): 2,8,1
• Chlorine (atomic number 17): 2,8,7
• Calcium (atomic number 20): 2,8,8,2

The electron configuration can be represented using diagrams showing shells as circles with electrons as dots or crosses. For ions, adjust the electron count: Na⁺ is 2,8 (lost one electron), while Cl⁻ is 2,8,8 (gained one electron).

The modern periodic table

The periodic table arranges elements in order of increasing atomic number. Elements in the same group have the same number of outer shell electrons, giving them similar chemical properties. Elements in the same period have the same number of electron shells.

Key features:

• Group number = number of outer shell electrons (for Groups 1-7 and 0)
• Period number = number of electron shells occupied
• Metals are found on the left and centre
• Non-metals are found on the right
• The stepped line separates metals from non-metals (with metalloids like silicon adjacent to it)

Sodium (Na) is in Group 1, Period 3: it has 1 outer electron and 3 shells (2,8,1). Chlorine (Cl) is in Group 7, Period 3: it has 7 outer electrons and 3 shells (2,8,7).

Trends in the periodic table

Group 1 (alkali metals):

• All have one outer electron, configuration ending in 1
• React by losing one electron to form +1 ions
• Reactivity increases down the group (outer electron is further from nucleus, weaker attraction, easier to lose)
• Melting points decrease down the group

Group 7 (halogens):

• All have seven outer electrons, configuration ending in 7
• React by gaining one electron to form -1 ions
• Reactivity decreases down the group (incoming electron is further from nucleus, weaker attraction, harder to gain)
• Melting and boiling points increase down the group

Group 0 (noble gases):

• All have full outer shells (helium has 2, others have 8)
• Unreactive due to stable electron configuration
• Exist as single atoms (monatomic)
• Boiling points increase down the group

Across Period 3 (sodium to argon):

• Atomic number increases by one each time
• Number of shells stays constant (3 shells)
• Number of outer electrons increases from 1 to 8
• Character changes from metal to non-metal
• Elements change from basic oxides (Na₂O, MgO) through amphoteric (Al₂O₃) to acidic oxides (P₄O₁₀, SO₂)

Using the periodic table in chemical bonding

The position of an element determines its valency and bonding behaviour:

• Group 1 elements form +1 ions (Na⁺, K⁺)
• Group 2 elements form +2 ions (Mg²⁺, Ca²⁺)
• Group 3 elements form +3 ions (Al³⁺)
• Group 5 elements form -3 ions (N³⁻) or share electrons
• Group 6 elements form -2 ions (O²⁻, S²⁻)
• Group 7 elements form -1 ions (Cl⁻, Br⁻)
• Group 0 elements do not form ions

CIE IGCSE Chemistry examiners expect students to predict ion charges from group numbers and write correct formulae using valency rules. For example, aluminium oxide must be Al₂O₃ because aluminium has valency 3 and oxygen has valency 2.

Worked examples

Example 1: Isotope calculation

Question: Boron has two isotopes: boron-10 (20% abundance) and boron-11 (80% abundance). Calculate the relative atomic mass of boron. [2 marks]

Solution:

Step 1: Multiply each mass number by its percentage abundance (10 × 20) + (11 × 80) = 200 + 880 = 1080 ✓

Step 2: Divide by 100 1080 ÷ 100 = 10.8 ✓

Answer: 10.8

Mark scheme notes: 1 mark for correct calculation setup, 1 mark for correct answer

Example 2: Electron configuration

Question: (a) An atom of potassium has atomic number 19. Write the electron configuration of a potassium atom. [1 mark] (b) Potassium forms an ion with charge +1. Write the electron configuration of this ion. [1 mark] (c) Explain, in terms of electron configuration, why potassium is placed in Group 1 and Period 4 of the periodic table. [2 marks]

Solution:

(a) 2,8,8,1 ✓

(b) 2,8,8 ✓

(c) Potassium has 1 electron in its outer shell, so it is in Group 1 ✓ It has 4 occupied shells/energy levels, so it is in Period 4 ✓

Mark scheme notes: Precise wording required — "outer shell" not just "outside", "occupied shells" not "orbits"

Example 3: Trends in Group 7

Question: The table shows information about three halogens.

(a) Predict the state of bromine at room temperature (25°C). [1 mark] (b) Explain the trend in boiling points down Group 7. [2 marks]

Solution:

(a) Liquid ✓ (Room temperature is between the melting point and boiling point of bromine)

(b) Boiling point increases down the group ✓ Because the molecules become larger/have more electrons/have stronger intermolecular forces ✓

Mark scheme notes: Must state "increases" explicitly and provide a correct reason referencing molecular size or intermolecular forces

Common mistakes and how to avoid them

• Confusing mass number with atomic number — Students often write the mass number when asked for protons. Remember: atomic number = number of protons; mass number = protons + neutrons. The atomic number is always the smaller number in standard notation (e.g. ¹⁷Cl₃₅ or Cl-35).

• Incorrect electron configurations — Writing 2,8,10 for calcium instead of 2,8,8,2. At IGCSE level, the third shell holds a maximum of 8 electrons for elements 1-20. Always fill shells in order and check you haven't exceeded the maximum.

• Mixing up ion charges and group numbers — Assuming Group 6 forms +6 ions. Non-metals form negative ions: Group 6 forms -2 ions (they gain 2 electrons to achieve a full outer shell of 8). Only Groups 1, 2 and 3 form positive ions at IGCSE level.

• Stating isotopes have different numbers of electrons — Isotopes differ only in neutron number. They have identical numbers of protons and electrons, which is why their chemical properties are the same. Chlorine-35 and chlorine-37 both have 17 protons and 17 electrons.

• Reversing reactivity trends — Stating Group 1 becomes less reactive down the group or Group 7 becomes more reactive down the group. Group 1 reactivity increases going down (easier to lose the outer electron); Group 7 reactivity decreases going down (harder to gain an electron).

• Poor explanations of periodic trends — Writing "sodium is more reactive than lithium because it's bigger" without explaining why size matters. Full explanation: sodium is more reactive because the outer electron is further from the nucleus (due to more shells), so there is weaker attraction between the nucleus and outer electron, making it easier to lose.

Exam technique for Atomic structure and the periodic table

• Command word "State" requires factual recall with no explanation (e.g. "State the number of neutrons in carbon-14" = answer "8", worth 1 mark). Command word "Explain" requires a reason using scientific principles (typically worth 2-3 marks). Structure explanations as: statement of trend + because + scientific reason.

• Drawing electron configurations — Examiners award marks for correct total electrons and correct distribution across shells. Use clear circles for shells and dots/crosses for electrons. Label each shell with its electron count (2, 8, 8, 1). When drawing ions, clearly show the charge (e.g. Na⁺) and adjust the electron count accordingly.

• Isotope questions typically follow a pattern: identify number of protons, neutrons and electrons (3 marks), then explain why chemical properties are identical (1-2 marks). Always show working for relative atomic mass calculations — marks are awarded for method even if the final answer contains an arithmetic error.

• Periodic table pattern questions often ask you to predict properties of an unknown element based on its position. First identify the group and period, state the electron configuration, then predict properties by comparison with known elements in the same group. Use data from the question to support numerical predictions (e.g. melting points, reactivity).

Quick revision summary

Atoms contain protons and neutrons in the nucleus, with electrons in shells around it. Atomic number equals protons; mass number equals protons plus neutrons. Isotopes have the same protons but different neutrons. Electron configuration determines chemical properties and position in the periodic table: group number equals outer electrons, period number equals occupied shells. Group 1 reactivity increases down; Group 7 reactivity decreases down. Elements form ions to achieve stable full outer shells (8 electrons or 2 for helium). Master electron configurations for elements 1-20, subatomic particle properties, and the ability to predict trends from periodic table position.