Periodic Table and Periodic Properties

Periodic Table and Periodic Properties

🟢 Lite — Quick Review (1h–1d)

Rapid summary for last-minute revision before your exam.

The modern periodic table arranges all elements by increasing atomic number (Z), with elements of similar electron configuration stacked in vertical groups (1–18) and placed across horizontal periods (1–7). The number of valence electrons equals the group digit for main-group (s- and p-block) elements, and the period number reveals the highest occupied electron shell.

Three trends to memorise cold: atomic radius decreases left → right across a period and increases top → bottom down a group; first ionization energy (IE₁) and electronegativity rise left → right and fall top → bottom; metallic character falls left → right and rises top → bottom. Successive ionisations obey IE₁ < IE₂ < IE₃ < IE₄ because removing an electron from a positive ion costs more energy. NECO SSCE Paper II Objective loves trend-comparison MCQs on Periods 2 and 3 — practise those first.

---

🟡 Standard — Regular Study (2d–2mo)

Standard content for students with a few days to months.

Atomic Number, Mass Number and Electron Shells

For a neutral atom, Z = number of protons = number of electrons, and mass number A = protons + neutrons, giving neutrons = A − Z. Electrons fill shells K, L, M, N, O, P, Q (periods 1–7), with maximum occupancy 2n² (2, 8, 18, 32 …). The outermost shell determines valency; the period number tells you which shell is filling.

Groups and Blocks

Groups 1 and 2 are the s-block (ns¹, ns²); groups 13–18 are the p-block (ns²np¹⁻⁶); the d-block sits in the middle as transition elements; the f-block (lanthanides, actinides) is placed below. Group 1 elements are alkali metals, Group 17 are halogens, Group 18 are noble gases with stable octets.

Periodic Trends

Across a period (e.g. Na → Ar) the nuclear charge rises while the shielding stays roughly constant, so Z_eff felt by outer electrons increases → atomic and ionic radii shrink, while ionization energy, electron affinity and electronegativity generally climb. Down a group (e.g. Li → Cs) each new shell is added and inner electrons shield outer ones, so Z_eff falls → radii grow and IE, EA and electronegativity drop. Metallic character shows the opposite pattern: stronger on the lower-left, weaker on the upper-right.

Worked Comparison (NECO-favourite)

Element	Na	Mg	Al	Si	P	S	Cl	Ar
Atomic radius (pm)	186	160	143	118	110	104	99	71*
1st IE (kJ mol⁻¹)	496	738	578	786	1012	1000	1251	1521
Electronegativity	0.93	1.31	1.61	1.90	2.19	2.58	3.16	—

*Note the dip from Mg → Al (new 3p sub-shell is higher in energy) and from P → S (paired 3p electron repels the incoming electron).

Successive Ionization Energies

A sudden large jump in successive IE values signals the removal of an electron from a shell closer to the nucleus. The element whose IEₙ/IEₙ₊₁ ratio is largest is the one with (n−1) outer electrons — the classic NECO Paper II question “the element with IE₁ = 786, IE₂ = 1577, IE₃ = 3232, IE₄ = 4356 kJ mol⁻¹” is silicon (group 14, 4 valence electrons).

---

🔴 Extended — Deep Study (3mo+)

Comprehensive coverage for students on a longer study timeline.

Explaining the Irregularities

The steady increase in IE across a period is broken by three recurring dips: between groups 2 and 13 (3p electron is further from nucleus and lightly shielded), between groups 15 and 16 (the paired electron in the new p-orbital experiences strong inter-electron repulsion, easing its removal), and at the start of every new period (a new, higher-energy shell is entered). Down a group, IE does not fall perfectly smoothly: fluorine has a higher IE than chlorine because its 2p electrons are so close to the nucleus that inter-electron repulsion partially offsets the pull.

Ionisation Energy, Electron Affinity and Electronegativity — Subtle Distinctions

• First ionization energy is the energy required to remove one mole of electrons from one mole of gaseous atoms (M(g) → M⁺(g) + e⁻); always endothermic, so values are positive in kJ mol⁻¹.
• Electron affinity is the energy change when one mole of gaseous atoms gains one mole of electrons (X(g) + e⁻ → X⁻(g)); energy is released, so modern IUPAC tables report a negative enthalpy change. A more negative EA = greater tendency to form an anion.
• Electronegativity (Pauling scale) is a relative, dimensionless measure of an atom’s ability to attract the bonding pair in a covalent bond. Fluorine tops the scale at 4.0; caesium and francium sit at ~0.7.

Cations vs Anions and Isoelectronic Series

A cation is smaller than its parent atom (lost electron, less shielding, more Z_eff per remaining electron); an anion is larger (added electron, more repulsion). An isoelectronic series — e.g. Na⁺, Mg²⁺, Al³⁺, Ne — all carry 10 electrons, so size falls as nuclear charge rises (Na⁺ > Mg²⁺ > Al³⁺). This same argument explains why O²⁻ > F⁻ > Ne > Na⁺ > Mg²⁺ > Al³⁺ across the 10-electron series.

Common NECO Traps

1. Confusing “atomic radius” with “ionic radius” when the species is charged.
2. Forgetting that metallic character decreases across a period but Na, Mg and Al are still metals while Si, P, S, Cl and Ar are not all non-metals (Si is a metalloid).
3. Predicting IE of N > O from the trend alone; the correct order is N > O, but the reason is the half-filled 2p³ stability of nitrogen.
4. Treating electron affinity values as if larger-positive means “more negative enthalpy”; modern sign convention makes more-negative the more exothermic.

Practice Prompts

1. Trend question: Arrange Na⁺, Mg²⁺, Al³⁺ and Ne in order of increasing ionic radius. (Answer: Al³⁺ < Mg²⁺ < Na⁺ < Ne.)
2. IE jump question: Successive IEs of element X (in kJ mol⁻¹) are 738, 1451, 7733, 10543, 13630 … identify the group of X. (Big jump between 2nd and 3rd → 2 valence electrons → Group 2; element is magnesium.)

---

Content adapted based on your selected roadmap duration. Switch tiers using the selector above.