Periodic Trends: Atomic Radius, Ionisation Energy, and Electronegativity

Why Do Periodic Trends Exist?

The periodic table isn't just organised by atomic number — it reveals predictable patterns in the properties of elements. These trends are driven by three fundamental factors:

1. Nuclear charge ($Z$): More protons = stronger pull on electrons.
2. Shielding: Inner electron shells block outer electrons from the full nuclear charge.
3. Atomic radius: Distance between the nucleus and the outermost electron.

Learning Goals: By the end of this guide, you should be able to:

1. Describe and explain the trends in atomic radius, ionisation energy, electron affinity, and electronegativity.
2. Use shielding and effective nuclear charge to explain each trend.
3. Identify anomalies (e.g., Group 13 IE drop) and explain them.
4. Apply trends to predict chemical behaviour.

1. Atomic Radius

Across a Period (Left → Right): Decreases

Why: Nuclear charge increases ($+1$ proton per element), but electrons are added to the same shell (same shielding). Greater effective nuclear charge pulls electrons closer.

Down a Group (Top → Bottom): Increases

Why: Each new period adds a new electron shell, increasing the distance between the nucleus and outermost electrons. Increased shielding also reduces the effective nuclear charge felt by outer electrons.

2. Ionisation Energy

First ionisation energy ($IE_1$): The energy required to remove one electron from a gaseous atom.

Across a Period: Generally increases

More protons → more nuclear charge → harder to remove an electron.

Two anomalies to know:

• Group 13 dip (e.g., $Al < Mg$): The electron being removed from Al is in a $3p$ orbital, which is higher in energy and easier to remove than Mg's $3s$ electron.
• Group 16 dip (e.g., $O < N$): In oxygen, two electrons are paired in the same $2p$ orbital, creating electron-electron repulsion that makes one easier to remove.

Down a Group: Decreases

Outermost electron is further from the nucleus and more shielded → easier to remove.

3. Electronegativity

Electronegativity: A measure of an atom's ability to attract the shared pair of electrons in a covalent bond (Pauling scale).

Across a Period: Increases

Higher nuclear charge pulls bonding electrons more strongly.

Down a Group: Decreases

Greater atomic radius and shielding reduce the pull on bonding electrons.

Most electronegative: Fluorine (4.0). Least electronegative (of main group elements): Francium.

4. Electron Affinity

First electron affinity: The enthalpy change when a gaseous atom gains one electron.

Across a Period: Generally becomes more exothermic (more negative)

Atoms with higher nuclear charge attract additional electrons more strongly.

Down a Group: Generally becomes less exothermic

The added electron is further from the nucleus → weaker attraction.

Master Comparison Table

Worked Examples

Example 1: Explain Why Na Has a Larger Radius Than Mg

Both are in Period 3. Mg has one more proton than Na ($Z = 12$ vs $11$), but both add electrons to the same $3s$ subshell. The extra proton in Mg creates a higher effective nuclear charge, pulling all electrons closer → smaller radius for Mg.

Example 2: Explain the IE₁ Dip from N to O

Nitrogen: $1s^2\ 2s^2\ 2p^3$ — all three $2p$ orbitals are singly occupied (half-filled, extra stable).

Oxygen: $1s^2\ 2s^2\ 2p^4$ — one $2p$ orbital has a paired electron, creating extra electron-electron repulsion.

This repulsion makes it easier to remove the paired electron from oxygen, so $IE_1(O) < IE_1(N)$.

Example 3: Why Is Fluorine More Electronegative Than Chlorine?

Both are in Group 17, but F is in Period 2 and Cl is in Period 3. Fluorine has a smaller atomic radius and less shielding, so its nucleus pulls on bonding electrons more strongly → higher electronegativity.

Common Mistakes

1. Saying "more electrons = more shielding" — Shielding comes from inner shells, not outer electrons. Adding electrons to the same shell doesn't increase shielding significantly.

2. Ignoring sub-shell effects on IE — The Group 13 and Group 16 anomalies are frequently tested. Don't assume IE increases monotonically across a period.

3. Confusing electronegativity with electron affinity — Electronegativity is a bond property (pulling shared electrons). Electron affinity is an atomic property (gaining a free electron).

4. Forgetting noble gases have no electronegativity — Electronegativity is defined for covalent bonds. Noble gases don't normally bond, so they have no Pauling electronegativity value.

Exam Tips (A-Level / AP / IB)

• For "explain the trend" questions, always mention nuclear charge AND shielding AND distance.
• Know the two IE anomalies (Groups 13 and 16) — they appear in almost every exam.
• When comparing elements in different groups AND periods, analyse each factor separately (nuclear charge, shielding, distance) before concluding.
• Use the electron configuration to justify anomalies: "The electron is removed from a $3p$ orbital which is higher in energy than $3s$."

Frequently Asked Questions

What causes periodic trends?

All periodic trends are driven by the interplay of nuclear charge (number of protons), electron shielding (inner shells blocking nuclear attraction), and atomic radius (distance from nucleus to outer electrons).

Why do noble gases have high ionisation energies?

Noble gases have full outer shells, making their electron configurations extremely stable. Removing an electron disrupts this stability, requiring very high energy.

Why does metallic character increase down a group?

Metals lose electrons easily. Down a group, the outer electrons are further from the nucleus and more shielded, making them easier to remove → stronger metallic character.

Related Topics

• Electron Configuration — The arrangement of electrons that drives all periodic trends.
• Atomic Models — How quantum mechanics explains shell structure and shielding.
• Chemical Bonds — Electronegativity differences determine bond type.