The Periodic Table is one of the most important topics in IGCSE Chemistry (Cambridge 0620/0971). It organises all known elements by atomic number and reveals patterns in their chemical behaviour. Whether you are sitting Paper 2, 4, or 6, a solid understanding of the Periodic Table will help you answer questions on element properties, trends, and reactions with confidence.

This comprehensive revision guide covers everything you need to know for the IGCSE exam: how the table is arranged, electron configurations, group properties, trends across periods and down groups, metals versus non-metals, and the key exam techniques that earn full marks.

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1. How the Periodic Table Is Arranged

1.1 A Brief History

Dmitri Mendeleev published the first widely recognised Periodic Table in 1869. He arranged elements by atomic mass and left gaps for elements not yet discovered. The modern table arranges elements by atomic number (number of protons), which resolved inconsistencies in Mendeleev’s original ordering.

1.2 Groups and Periods

Term	Definition	What it tells you
Group	A vertical column	Elements in the same group have the same number of outer-shell (valence) electrons, so they share similar chemical properties.
Period	A horizontal row	Elements in the same period have the same number of electron shells. Moving left to right across a period, the atomic number increases by one each time.

Exam tip: If a question gives you an element’s electron configuration — for example 2,8,3 — the group number equals the number of outer electrons (Group III) and the period number equals the number of shells (Period 3).

1.3 Key Regions of the Table

• Groups I and II (left side) — Reactive metals (alkali metals and alkaline earth metals).
• Transition elements (centre block) — Metals with variable oxidation states, coloured compounds, and catalytic properties.
• Groups IV to VII (right side) — Mix of non-metals and metalloids.
• Group 0 (VIII) — Noble gases with full outer shells (stable, unreactive).

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2. Electron Configuration and the Periodic Table

Understanding electron configuration is the key to unlocking the Periodic Table. The way electrons fill shells determines an element’s position and its chemical properties.

2.1 Rules for Filling Electron Shells

1. The first shell holds a maximum of 2 electrons.
2. The second shell holds a maximum of 8 electrons.
3. The third shell holds a maximum of 8 electrons (at IGCSE level).
4. Electrons fill the lowest available shell first before moving to the next.

2.2 Worked Examples

Element	Atomic number	Electron configuration	Group	Period
Lithium (Li)	3	2, 1	I	2
Carbon (C)	6	2, 4	IV	2
Sodium (Na)	11	2, 8, 1	I	3
Chlorine (Cl)	17	2, 8, 7	VII	3
Calcium (Ca)	20	2, 8, 8, 2	II	4

2.3 Why Electron Configuration Matters

Elements in the same group react similarly because they have the same number of valence electrons. For instance, lithium (2,1), sodium (2,8,1), and potassium (2,8,8,1) all have one outer electron, which they readily lose to form a +1 ion. This is why all Group I metals react with water to produce hydrogen gas and an alkaline solution.

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3. Metals, Non-Metals, and Metalloids

The Periodic Table has a clear dividing line between metals and non-metals, running roughly in a diagonal staircase from boron to astatine.

3.1 Properties of Metals

• Located on the left and centre of the table.
• Lose electrons in reactions to form positive ions (cations).
• Good conductors of heat and electricity.
• Malleable (can be hammered into shape) and ductile (can be drawn into wire).
• Generally have high melting and boiling points (except Group I metals).
• Form basic or amphoteric oxides.

3.2 Properties of Non-Metals

• Located on the right side of the table.
• Gain or share electrons in reactions to form negative ions (anions) or covalent bonds.
• Poor conductors of heat and electricity (except graphite).
• Often brittle as solids and may exist as gases or liquids at room temperature.
• Generally have lower melting and boiling points.
• Form acidic or neutral oxides.

3.3 The Transition Elements

Transition metals sit in the central block between Groups II and III. At IGCSE level, you should know that they:

• Have high melting points and high densities compared to Group I metals.
• Often form coloured compounds (e.g. copper(II) sulfate is blue, iron(III) oxide is brown).
• Can show variable oxidation states (e.g. iron can be Fe²⁺ or Fe³⁺).
• Act as catalysts in industrial processes (e.g. iron in the Haber process, vanadium(V) oxide in the Contact process).

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4. Group I — The Alkali Metals

Group I contains lithium (Li), sodium (Na), potassium (K), rubidium (Rb), caesium (Cs), and francium (Fr). These are the most reactive metals in the Periodic Table.

4.1 Physical Properties

• Soft — can be cut with a knife.
• Low density — lithium, sodium, and potassium float on water.
• Low melting points for metals — they decrease going down the group.
• Shiny when freshly cut but tarnish rapidly in air (they react with oxygen).

4.2 Chemical Properties

All Group I metals have one outer electron which they lose to form a +1 ion. Their reactions include:

Reaction with water: $$2\text{Na}(s) + 2\text{H}_2\text{O}(l) \rightarrow 2\text{NaOH}(aq) + \text{H}_2(g)$$

• The metal moves on the surface, fizzes, and may melt into a ball.
• The solution formed is alkaline (hence “alkali metals”) — it turns universal indicator blue/purple.
• Potassium reacts more vigorously than sodium, often catching fire with a lilac flame.

Reaction with oxygen: $$4\text{Li}(s) + \text{O}_2(g) \rightarrow 2\text{Li}_2\text{O}(s)$$

Reaction with chlorine: $$2\text{Na}(s) + \text{Cl}_2(g) \rightarrow 2\text{NaCl}(s)$$

4.3 Trend in Reactivity Down Group I

Reactivity increases going down the group. This is because:

1. Each successive element has more electron shells, so the outer electron is further from the nucleus.
2. There is greater shielding from inner electrons.
3. The electrostatic attraction between the nucleus and the outer electron is weaker.
4. Therefore, the outer electron is lost more easily, and the element reacts more vigorously.

Exam tip: When asked to explain a trend, always link the number of shells → distance → shielding → ease of losing the electron. A bare statement like “it reacts faster” will not earn the explanation mark.

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5. Group VII — The Halogens

Group VII contains fluorine (F), chlorine (Cl), bromine (Br), iodine (I), and astatine (At). These are the most reactive non-metals.

5.1 Physical Properties

Halogen	State at room temperature	Colour
Fluorine	Gas	Pale yellow
Chlorine	Gas	Yellow-green
Bromine	Liquid	Red-brown
Iodine	Solid	Dark grey/purple vapour

• Melting and boiling points increase going down the group (the molecules become larger, and intermolecular forces become stronger).
• They exist as diatomic molecules (F₂, Cl₂, Br₂, I₂).

5.2 Chemical Properties

All halogens have seven outer electrons and need to gain one electron to achieve a stable octet, forming a −1 ion.

Reaction with metals (e.g. iron): $$2\text{Fe}(s) + 3\text{Cl}_2(g) \rightarrow 2\text{FeCl}_3(s)$$

Reaction with hydrogen: $$\text{H}_2(g) + \text{Cl}_2(g) \rightarrow 2\text{HCl}(g)$$

5.3 Trend in Reactivity Down Group VII

Reactivity decreases going down the group. This is because:

1. Each successive element has more electron shells, so the outer shell is further from the nucleus.
2. There is greater shielding from inner electrons.
3. The electrostatic attraction between the nucleus and an incoming electron is weaker.
4. It is therefore harder to gain the extra electron needed to form a −1 ion.

Key contrast: Group I reactivity increases down the group (easier to lose an electron), while Group VII reactivity decreases down the group (harder to gain an electron). Examiners love asking you to compare these two trends.

5.4 Displacement Reactions of Halogens

A more reactive halogen will displace a less reactive halogen from its salt solution:

• Chlorine + potassium bromide → potassium chloride + bromine $$\text{Cl}_2(aq) + 2\text{KBr}(aq) \rightarrow 2\text{KCl}(aq) + \text{Br}_2(aq)$$

• Chlorine + potassium iodide → potassium chloride + iodine $$\text{Cl}_2(aq) + 2\text{KI}(aq) \rightarrow 2\text{KCl}(aq) + \text{I}_2(aq)$$

• Bromine + potassium iodide → potassium bromide + iodine $$\text{Br}_2(aq) + 2\text{KI}(aq) \rightarrow 2\text{KBr}(aq) + \text{I}_2(aq)$$

A less reactive halogen cannot displace a more reactive one. For example, bromine will not displace chlorine from potassium chloride.

Exam tip: In Paper 6 (Alternative to Practical), you may be asked to predict observations when halogens are added to halide solutions. Remember: bromine turns the solution orange/brown, and iodine turns it brown/yellow (or dark brown with starch indicator turning blue-black).

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6. Group 0 — The Noble Gases

Group 0 (sometimes called Group VIII or Group 18) contains helium (He), neon (Ne), argon (Ar), krypton (Kr), xenon (Xe), and radon (Rn).

6.1 Why Are They Unreactive?

Noble gases have full outer electron shells (helium has 2; the others have 8). Because their electron configuration is already stable, they have no tendency to gain, lose, or share electrons. This makes them chemically inert under normal conditions.

6.2 Properties and Uses

• They are colourless, odourless gases.
• They exist as monatomic molecules (single atoms, not diatomic).
• Boiling points increase going down the group (larger atoms have stronger intermolecular forces).
• Uses: Helium — balloons and airships (low density, non-flammable); Neon — advertising signs (glows red-orange); Argon — welding and filling light bulbs (provides an inert atmosphere).

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7. Trends Across a Period

While groups show trends vertically, important patterns also exist horizontally across a period.

7.1 Across Period 3 (Na → Ar)

Property	Trend left to right
Atomic number	Increases (more protons)
Metallic character	Decreases — Na, Mg, Al are metals; Si is a metalloid; P, S, Cl, Ar are non-metals
Type of bonding	Changes from metallic → giant covalent (Si) → simple molecular → monatomic (Ar)
Melting point	Increases across metals (Na < Mg < Al), peaks at silicon (giant covalent), then drops sharply for simple molecular substances
Electrical conductivity	High for metals, low/zero for non-metals (except graphite in Period 2)
Type of oxide	Basic (Na₂O, MgO) → Amphoteric (Al₂O₃) → Acidic (SiO₂, P₄O₁₀, SO₃)

7.2 Atomic Size Across a Period

Atoms get smaller across a period because the nuclear charge increases (more protons) while electrons are added to the same shell. The stronger attraction pulls electrons closer to the nucleus.

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8. Common Exam Questions and How to Answer Them

8.1 “State” Questions (1 mark)

State the group and period of an element with electron configuration 2, 8, 6.

Answer: Group VI, Period 3.

Technique: Outer electrons = group number. Number of shells = period number.

8.2 “Describe” Questions (2–3 marks)

Describe the trend in reactivity of Group I metals going down the group.

Answer: Reactivity increases going down the group. The metals react more vigorously with water — lithium fizzes gently, sodium fizzes more vigorously and may melt, potassium reacts violently and catches fire.

8.3 “Explain” Questions (3–4 marks)

Explain why potassium is more reactive than sodium.

Answer: Potassium has more electron shells than sodium (4 vs 3). The outer electron is further from the nucleus and is more shielded by inner electrons. The attraction between the nucleus and the outer electron is therefore weaker, so the outer electron is lost more easily. This makes potassium more reactive.

8.4 “Predict” Questions

Predict whether bromine would displace chlorine from sodium chloride solution. Explain your answer.

Answer: No, bromine would not displace chlorine. Bromine is less reactive than chlorine because it has more electron shells, making it harder for bromine to attract and gain an electron. A less reactive halogen cannot displace a more reactive halogen from its compound.

8.5 “Compare” Questions

Compare the properties of Group I metals with transition metals.

Answer: Group I metals have lower melting points, lower densities, and are softer than transition metals. Group I metals have only one oxidation state (+1), whereas transition metals often show variable oxidation states. Transition metals form coloured compounds; Group I compounds are typically white or colourless. Transition metals are often used as catalysts; Group I metals are not.

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9. Common Mistakes to Avoid

1. Confusing group and period. Groups are vertical (same outer electrons); periods are horizontal (same number of shells).

2. Saying Group VII reactivity increases down the group. It decreases — larger atoms find it harder to attract an incoming electron.

3. Forgetting to explain trends fully. Simply saying “the atom gets bigger” is not enough. Link the number of shells → distance from nucleus → shielding → strength of attraction → ease of gaining or losing electrons.

4. Mixing up displacement rules. A more reactive halogen displaces a less reactive one, never the other way round.

5. Ignoring noble gas stability. Noble gases are unreactive because they already have full outer shells — not because they are “not interested” in reacting.

6. Writing unbalanced equations. Always check that the number of atoms of each element is the same on both sides.

7. Confusing electron configuration with atomic number. The electron configuration 2,8,1 tells you the element has 11 electrons (and therefore 11 protons and an atomic number of 11 — sodium), not that its atomic number is 2, 8, or 1.

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10. Revision Checklist

Use this checklist to make sure you have covered all key areas before your exam:

• I can define group and period and state what they tell us about electron configuration.
• I can write the electron configuration of the first 20 elements and use it to identify group and period.
• I can list the properties of metals and non-metals and state where they are found in the table.
• I can describe and explain the trend in reactivity down Group I (alkali metals).
• I can describe and explain the trend in reactivity down Group VII (halogens).
• I can predict the outcome of halogen displacement reactions.
• I can explain why noble gases are unreactive.
• I can describe trends across Period 3 (metallic character, melting point, type of oxide).
• I can compare Group I metals with transition elements.
• I can write balanced chemical equations for reactions of Group I metals and halogens.

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