The Periodic Table

This lesson covers the periodic table — its history, organisation and significance — as required by AQA GCSE Chemistry specification (5.1.2). The periodic table is one of the most important tools in chemistry. It organises all known elements according to their atomic structure and allows chemists to predict the properties and reactions of elements.

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Early Attempts to Classify Elements

Before the modern periodic table was developed, scientists attempted to organise elements in various ways. You need to know the key historical developments:

John Newlands (1864) — Law of Octaves

• Newlands arranged the known elements in order of atomic weight (relative atomic mass).
• He noticed that every eighth element had similar properties — he called this the Law of Octaves (by analogy with musical octaves).
• Problems: Newlands' table only worked for the first 16 elements. He forced some elements into groups where they did not fit. He did not leave gaps for undiscovered elements. His ideas were ridiculed by the scientific community at the time.

Dmitri Mendeleev (1869) — The First Successful Periodic Table

• Mendeleev also arranged elements by atomic weight but with a crucial difference.
• He left gaps in his table for elements that had not yet been discovered.
• He swapped the order of some elements so that elements with similar properties were in the same group (vertical column).
• He used the gaps to predict the properties of undiscovered elements — for example, he predicted the existence and properties of germanium (which he called "eka-silicon").
• When these elements were later discovered and their properties matched his predictions, Mendeleev's table was accepted as a major scientific achievement.

Exam Tip: The key reason Mendeleev's table was eventually accepted was that his predictions about undiscovered elements were proved correct. In the exam, always explain that he left gaps AND that his predictions were later confirmed. Both points are needed for full marks.

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The Modern Periodic Table

The modern periodic table arranges elements in order of atomic number (proton number), not atomic weight. This was made possible after the discovery of protons and the concept of atomic number.

Key Features

Feature	Description
Rows (Periods)	Horizontal rows. Elements in the same period have the same number of electron shells.
Columns (Groups)	Vertical columns. Elements in the same group have the same number of outer electrons and similar chemical properties.
Atomic number	Elements are arranged in order of increasing atomic number (left to right, top to bottom).
Metals and non-metals	Metals are on the left and centre; non-metals are on the right. A "staircase" line roughly divides them.

Groups and Periods

Term	Definition	Chemical Significance
Group	Vertical column	Elements have the same number of outer electrons and similar properties
Period	Horizontal row	Elements have the same number of electron shells
Group 1	Alkali metals	1 outer electron, very reactive metals
Group 7	Halogens	7 outer electrons, reactive non-metals
Group 0	Noble gases	Full outer shell, very unreactive
Transition metals	Central block (between Group 2 and Group 3)	Less reactive, form coloured compounds, variable oxidation states

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Organisation of the Periodic Table

graph TD
    A["Periodic Table"] --> B["Groups (vertical columns)"]
    A --> C["Periods (horizontal rows)"]
    B --> D["Same number of outer electrons"]
    B --> E["Similar chemical properties"]
    C --> F["Same number of electron shells"]
    A --> G["Metals (left/centre)"]
    A --> H["Non-metals (right)"]
    G --> I["Lose electrons to form positive ions"]
    H --> J["Gain/share electrons to form negative ions or molecules"]

    style A fill:#2c3e50,color:#fff
    style B fill:#e74c3c,color:#fff
    style C fill:#3498db,color:#fff
    style G fill:#e67e22,color:#fff
    style H fill:#27ae60,color:#fff

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Trends Across a Period

As you move across a period (left to right), the properties of elements change in a predictable way:

Property	Trend Across a Period
Atomic number	Increases by 1 for each element
Number of outer electrons	Increases from 1 to 8 (or 2 for Period 1)
Metallic character	Decreases — elements change from metals to non-metals
Reactivity of metals	Decreases across the period
Reactivity of non-metals	Increases across the period (up to Group 7)
Type of oxide	Changes from basic (metal oxides) to acidic (non-metal oxides)

Example: Period 3

Element	Na	Mg	Al	Si	P	S	Cl	Ar
Type	Metal	Metal	Metal	Metalloid	Non-metal	Non-metal	Non-metal	Noble gas
Outer electrons	1	2	3	4	5	6	7	8

Exam Tip: When asked about trends across a period, always link your answer to electronic structure. For example, metallic character decreases because elements have more outer electrons and are more likely to gain electrons (non-metal behaviour) rather than lose them (metal behaviour).

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Trends Down a Group

As you move down a group, the properties change in a predictable way:

Property	Trend Down a Group
Number of electron shells	Increases (one more shell per period)
Atomic radius	Increases (more shells, electrons further from nucleus)
Reactivity of metals	Increases (outer electron easier to lose — further from nucleus, more shielding)
Reactivity of non-metals	Decreases (harder to attract an extra electron — further from nucleus, more shielding)

Exam Tip: The reasons for reactivity trends down a group are always linked to atomic size, distance from the nucleus, and electron shielding. Metals become more reactive going down because the outer electron is further from the nucleus and easier to remove. Non-metals become less reactive going down because the nucleus has less attraction for incoming electrons due to increased distance and shielding.

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Metals vs Non-Metals

Property	Metals	Non-Metals
Position in periodic table	Left and centre	Right
Electronic behaviour	Lose electrons to form positive ions (cations)	Gain electrons to form negative ions (anions) or share electrons
Bonding	Metallic bonding (between metals) or ionic bonding (with non-metals)	Covalent bonding (with other non-metals) or ionic bonding (with metals)
Conductivity	Good conductors of electricity and heat	Poor conductors (insulators), except graphite and silicon
Melting/boiling points	Generally high (except Group 1 metals)	Generally low (except diamond, silicon dioxide)
State at room temperature	Solid (except mercury, which is liquid)	Solid, liquid (bromine) or gas
Appearance	Shiny, lustrous	Dull (except iodine crystals)
Malleability	Malleable and ductile	Brittle (when solid)

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Why the Periodic Table Is So Useful

The periodic table allows chemists to: