Periodic Trends: Atomic Radius, Ionisation Energy and Electronegativity

The two forces behind everything

• •Across a period (left → right): protons are added but electrons go into the same shell, so shielding barely changes and effective nuclear charge rises. The nucleus pulls the outer shell in tighter.
• •Down a group (top → bottom): a whole new shell is added and shielding increases sharply, so the outer electrons are further out and held more loosely.

The trends themselves

The exceptions that carry marks

• •Ionisation energy dips from group 2 to group 13 (Be → B): the electron removed comes from a higher-energy p orbital.
• •It dips again from group 15 to 16 (N → O): oxygen's paired p electrons repel each other, so one is easier to remove.
• •Noble gases have the highest ionisation energies in their period — full shells are stable.
• •Fluorine, not chlorine, has the highest electronegativity (4.0 on the Pauling scale) — but chlorine has the more negative electron affinity, because fluorine's small size crowds the incoming electron.
• •Transition metal radii change very little across the d block, since added electrons go into an inner d subshell that shields well.

Ionic radius

• •Cations are always smaller than the parent atom: electrons are removed, often emptying a whole shell, and the remaining electrons feel a higher charge per electron.
• •Anions are always larger: added electrons increase repulsion while nuclear charge is unchanged.
• •For an isoelectronic series (same electron count, e.g. N³⁻, O²⁻, F⁻, Na⁺, Mg²⁺), radius decreases as nuclear charge increases.

How trends show up in reactivity

• •Metals lose electrons, so reactivity rises with lower ionisation energy — down group 1, caesium is more reactive than lithium.
• •Non-metals gain electrons, so reactivity rises with higher electronegativity — up group 17, fluorine is more reactive than iodine.
• •Electronegativity difference predicts bond type: roughly under 0.4 non-polar covalent, 0.4–1.7 polar covalent, above 1.7 largely ionic.

Flashcards

Tap a card to reveal the answer.

• Define effective nuclear charge.⌄

  The net positive charge an outer electron actually experiences after inner electrons shield it.

• Atomic radius across a period?⌄

  Decreases — effective nuclear charge rises while shielding stays roughly constant.

• Atomic radius down a group?⌄

  Increases — each row adds a shell and more shielding.

• Define first ionisation energy.⌄

  The energy needed to remove one mole of electrons from one mole of gaseous atoms.

• Why does ionisation energy dip from N to O?⌄

  Oxygen's first paired 2p electrons repel each other, making one easier to remove.

• Which element is the most electronegative?⌄

  Fluorine, 4.0 on the Pauling scale.

• Which has the most negative electron affinity?⌄

  Chlorine — fluorine's small size makes the added electron more crowded.

• Are cations larger or smaller than their parent atoms?⌄

  Smaller, usually because an entire outer shell is lost.

• Order an isoelectronic series by radius.⌄

  Radius falls as nuclear charge rises: N³⁻ > O²⁻ > F⁻ > Na⁺ > Mg²⁺.

• Why is caesium more reactive than lithium?⌄

  Its outer electron is further out and better shielded, so ionisation energy is lower.

• What does an electronegativity difference above about 1.7 suggest?⌄

  Largely ionic bonding.

• Why do transition metal radii change so little across the d block?⌄

  Added electrons enter an inner d subshell that shields the outer s electrons well.

Practice quiz

Answer first, then open the explanation.

1. 1. Atomic radius decreases across a period because:

   • A. Electrons are removed
   • B. Effective nuclear charge increases while shielding stays similar
   • C. Shielding increases sharply
   • D. A new shell is added
   Show answer

   B. Effective nuclear charge increases while shielding stays similar
   Protons are added to the same shell, so the nucleus pulls the outer electrons in more tightly.

2. 2. Which element has the highest first ionisation energy?

   • A. Sodium
   • B. Chlorine
   • C. Helium
   • D. Caesium
   Show answer

   C. Helium
   Helium: a tiny atom with a full shell and almost no shielding.

3. 3. The dip in ionisation energy from beryllium to boron is caused by:

   • A. Electron pairing repulsion
   • B. Removal of an electron from a higher-energy p orbital
   • C. Increased shielding by a new shell
   • D. A half-filled d subshell
   Show answer

   B. Removal of an electron from a higher-energy p orbital
   Boron's outermost electron is in 2p, which is higher in energy than beryllium's 2s.

4. 4. Which ion is the smallest?

   • A. N³⁻
   • B. F⁻
   • C. Na⁺
   • D. Mg²⁺
   Show answer

   D. Mg²⁺
   All four are isoelectronic; magnesium has the most protons pulling on the same ten electrons.

5. 5. Reactivity of the halogens:

   • A. Increases down the group
   • B. Increases up the group
   • C. Is the same throughout
   • D. Peaks in the middle
   Show answer

   B. Increases up the group
   Halogens react by gaining electrons, and electronegativity increases up the group — fluorine is the most reactive.

6. 6. An electronegativity difference of 0.9 between two bonded atoms suggests:

   • A. Non-polar covalent
   • B. Polar covalent
   • C. Ionic
   • D. Metallic
   Show answer

   B. Polar covalent
   Roughly 0.4–1.7 is the polar covalent range: unequal sharing, but no full transfer.

7. 7. Anions are larger than their parent atoms because:

   • A. They gain a shell
   • B. Nuclear charge decreases
   • C. Added electrons increase repulsion with unchanged nuclear charge
   • D. They lose shielding electrons
   Show answer

   C. Added electrons increase repulsion with unchanged nuclear charge
   The proton count is unchanged, so the extra electrons spread the electron cloud out.