Notes › Inorganic Chemistry

Periodic Classification & Trends

1. The Modern Periodic Law

Properties of elements are a periodic function of their atomic numbers (Moseley), not their atomic masses as Mendeleev proposed. The long form of the table has 18 groups and 7 periods, and is divided into s, p, d and f blocks according to the subshell that receives the last electron.

2. Atomic and Ionic Radius

  • Decreases across a period (nuclear charge increases while electrons enter the same shell).
  • Increases down a group (new shells are added).
  • Cations are smaller than their parent atoms; anions are larger.
  • Isoelectronic species: the higher the nuclear charge, the smaller the ion: $O^{2-} > F^- > Na^+ > Mg^{2+} > Al^{3+}$.

3. Ionisation Enthalpy

The energy needed to remove the outermost electron from an isolated gaseous atom. It generally increases across a period and decreases down a group.

⚠️ The Two Classic Anomalies

Be > B and N > O. Boron loses a 2p electron, which is shielded by 2s and easier to remove. Oxygen loses a paired 2p electron, and electron–electron repulsion makes that easier than breaking nitrogen's stable half-filled $2p^3$.

Successive ionisation enthalpies always increase ($IE_1 < IE_2 < IE_3$). A very large jump shows that a stable noble-gas core has been reached.

4. Electron Gain Enthalpy

  • Becomes more negative across a period and less negative down a group.
  • Chlorine has the most negative electron gain enthalpy, not fluorine: F is so small that the incoming electron feels strong repulsion from the compact 2p electrons.
  • Noble gases, Be, Mg and N have positive or near-zero values because of stable configurations.

5. Electronegativity

The tendency of a bonded atom to attract shared electrons. On the Pauling scale, F (4.0) > O (3.5) > N ≈ Cl (3.0). It increases across a period and decreases down a group.

Interactive diagram: open this page in a browser to use it.

6. Metallic Character and Oxides

  • Metallic character decreases across a period and increases down a group.
  • Oxides become more acidic across a period: $Na_2O$ (basic), $Al_2O_3$ (amphoteric), $Cl_2O_7$ (strongly acidic).

7. Diagonal Relationship

Li–Mg, Be–Al and B–Si show similar properties because their charge-to-radius ratios (polarising power) are similar.

💡 Valency Rule of Thumb

For s- and p-block elements, valency equals the number of valence electrons or 8 minus that number. This lets you predict formulae such as $AlCl_3$, $Mg_3N_2$ and $SiO_2$ straight from group numbers.

Practice questions

  1. Which element has the highest first ionisation enthalpy?

    1. B
    2. C
    3. N
    4. O
    Answer

    C. Nitrogen's half-filled 2p³ subshell is extra stable, so its IE₁ exceeds that of oxygen.

  2. The element with the most negative electron gain enthalpy is:

    1. F
    2. Cl
    3. Br
    4. O
    Answer

    B. Fluorine's small 2p subshell causes inter-electronic repulsion, so chlorine has the most negative value.

  3. The correct order of size among isoelectronic ions is:

    1. Na⁺ > Mg²⁺ > F⁻ > O²⁻
    2. O²⁻ > F⁻ > Na⁺ > Mg²⁺
    3. Mg²⁺ > Na⁺ > F⁻ > O²⁻
    4. F⁻ > O²⁻ > Mg²⁺ > Na⁺
    Answer

    B. All have 10 electrons; the ion with the smallest nuclear charge (O²⁻, Z = 8) is the largest.

  4. Which oxide is amphoteric?

    1. Na₂O
    2. MgO
    3. Al₂O₃
    4. SO₃
    Answer

    C. Al₂O₃ dissolves in both acids and bases.

  5. The diagonal relationship is seen between:

    1. Li and Na
    2. Be and Al
    3. B and C
    4. Mg and Ca
    Answer

    B. Be and Al have similar charge-to-radius ratios, giving similar chemistry.

  6. Across a period from left to right, atomic radius generally:

    1. Increases
    2. Decreases
    3. Stays constant
    4. First increases then decreases
    Answer

    B. Effective nuclear charge rises while electrons enter the same shell, pulling them closer.

Chemical Bonding →