VCE Chemistry (Year 11) · Unit 1 — Atomic Structure · 11m read

The Periodic Table — Trends

Ever noticed that some elements are super reactive while others barely react at all? The periodic table isn't random — it's organized so that elements in the same column share similar properties, and properties change in predictable patterns as you move across rows. Understanding these trends (like how easily atoms lose electrons or how big they are) helps you predict how chemicals will behave together, which is essential for VCE chemistry and understanding real-world reactions from rusting to battery chemistry.

Learning objective: Explain periodic trends in atomic properties.

By the end of this lesson you can:

• I describe atomic radius, ionisation, electronegativity trends.
• I link trends to electron configuration.
• I predict properties from position.

Periodic trends are tug-of-war strength. Across a period, the nucleus pulls harder on electrons; down a group, extra shells weaken the pull.

Trends across and down

ORGANISATION:
• GROUPS (columns, 1-18) — similar valence, similar chemistry.
• PERIODS (rows, 1-7) — same highest occupied shell.

KEY GROUPS:
• Group 1 — Alkali metals (very reactive).
• Group 2 — Alkaline earth metals.
• Group 17 — Halogens (very reactive non-metals).
• Group 18 — Noble gases (unreactive).

ATOMIC RADIUS:
• DECREASES across a period (higher Z, stronger nuclear pull).
• INCREASES down a group (more shells).

IONISATION ENERGY (energy to remove an electron):
• INCREASES across a period.
• DECREASES down a group.
• Low for metals; high for non-metals.

SUCCESSIVE IONISATION ENERGIES:
• Big jumps indicate transitions between shells.
• Useful for identifying group number.

ELECTRONEGATIVITY (attraction for shared electrons):
• INCREASES across a period.
• DECREASES down a group.
• Fluorine = highest (3.98 on Pauling scale).

METALLIC CHARACTER:
• Increases down a group.
• Decreases across a period.

REACTIVITY:
• Metals: reactivity increases DOWN the group.
• Non-metals: reactivity increases UP the group (halogens).

EXPLANATIONS:
• Atomic radius: more shells OR less effective nuclear charge.
• Ionisation: pull on outermost electron.
• Electronegativity: ability of nucleus to attract bonding electrons.

EFFECTIVE NUCLEAR CHARGE:
• Felt by valence electrons.
• Inner electrons SHIELD outer electrons.
• Higher effective Z → more pull.

METALS vs NON-METALS vs METALLOIDS:
• Metals — left side. Shiny, conductive, malleable.
• Non-metals — right side. Brittle, poor conductors.
• Metalloids (B, Si, Ge, As, Sb, Te) — borderline properties.

TRANSITION METALS:
• Variable oxidation states.
• Coloured compounds.
• Catalytic activity.

SIZES OF IONS:
• Cations SMALLER than parent atom (lost shell).
• Anions LARGER than parent atom (added electrons, repulsion).

EXAM TIP:
• Sketch trend arrows on a mini periodic table.
• Explanations often go deeper than memory — link to shells.

EXCEPTIONS:
• Half-filled stability (N vs O first ionisation).
• Some irregularities in transition metals.

Trend directions

• Size: ↓ down, ← across.
• IE / EN: ↑ across, ↓ down.
• Metal: ↓ down, ← across.

Quick reference

| Trend | Across period → | Down group ↓ |
|---|---|---|
| Atomic radius | Decreases | Increases |
| Electronegativity | Increases | Decreases |
| First ionisation energy | Increases | Decreases |
| Metallic character | Decreases | Increases |

Worked example (Medium)

Rank Na, K, Li by atomic radius.

1. Same group (1). Radius increases DOWN the group.
2. Li (period 2), Na (3), K (4). K > Na > Li.

Answer: K > Na > Li (K largest because lowest in group 1).

Worked example (Hard)

The first ionisation energies of N (1402 kJ/mol) and O (1314 kJ/mol) seem to break the trend — O is higher in Z but has LOWER IE than N. Explain this anomaly.

1. Expected trend. IE generally increases across a period due to increasing nuclear charge and decreasing atomic radius.
2. Electron configuration. N: 1s² 2s² 2p³ (half-filled 2p). O: 1s² 2s² 2p⁴ (one 2p orbital has a PAIRED electron).
3. Explanation. The extra electron in oxygen occupies an already half-filled 2p orbital, creating electron-electron repulsion within that orbital. This makes it EASIER to remove than expected. Nitrogen's half-filled 2p subshell has extra stability (one electron per orbital, no pairing repulsion).

Answer: Oxygen's lower IE is due to electron-electron repulsion in the doubly-occupied 2p orbital, which destabilises the extra electron despite higher nuclear charge.

Common mistakes

• Saying "atoms want to have a full shell" Atoms don't "want" anything — this is anthropomorphism. Atoms achieve lower energy states when they have stable electron configurations. Fix: Say "atoms achieve a more stable electron configuration" rather than "atoms want full shells".
• Ignoring shielding when explaining trends Across a period, nuclear charge increases but shielding from INNER electrons stays roughly constant. Students often mention only nuclear charge without discussing shielding. Fix: Always mention BOTH effective nuclear charge and electron shielding when explaining periodic trends.
• Confusing ionisation energy with electronegativity Ionisation energy = energy to REMOVE an electron from a gaseous atom. Electronegativity = ability to ATTRACT shared electrons in a bond. Different concepts with similar trends. Fix: IE = removing from isolated atom; EN = attracting in a bond. Both increase across a period, but for different reasons.

Effective nuclear charge — the hidden driver of periodic trends

Nearly every periodic trend can be explained by a single concept: effective nuclear charge ($Z_{\text{eff}}$). This is the net positive charge experienced by valence electrons after accounting for the shielding effect of inner electrons.

The formula is approximately: $Z_{\text{eff}} = Z - S$, where $Z$ is the total nuclear charge (number of protons) and $S$ is the shielding constant (roughly equal to the number of inner-shell electrons).

Across a period (e.g., Li to Ne): $Z$ increases by 1 with each element, but electrons are added to the SAME shell — they don't shield each other very effectively. So $Z_{\text{eff}}$ increases steadily. This explains:
- Decreasing atomic radius: stronger pull draws electrons closer.
- Increasing ionisation energy: harder to remove tightly-held electrons.
- Increasing electronegativity: nucleus attracts bonding electrons more strongly.

Down a group (e.g., Li, Na, K): each element adds a new shell of inner electrons, dramatically increasing shielding. Even though $Z$ increases, $Z_{\text{eff}}$ experienced by valence electrons stays roughly constant or decreases slightly. Meanwhile, valence electrons are FURTHER from the nucleus. This explains:
- Increasing atomic radius: more shells, similar pull.
- Decreasing ionisation energy: easier to remove distant, shielded electrons.
- Increasing metallic character: valence electrons lost more easily.

Understanding $Z_{\text{eff}}$ transforms periodic trends from memorised facts into logical consequences of atomic structure — exactly the kind of reasoning VCAA rewards.

Practice questions

1. Going ACROSS a period, electronegativity:

   1. Decreases
   2. Increases
   3. Stays same
   4. Fluctuates

   (1 marks)

2. Why is fluorine's electronegativity higher than chlorine's?

   (3 marks)

Summary

• Describe atomic radius, ionisation, electronegativity trends.
• Link trends to electron configuration.
• Predict properties from position.