Deep-Dive Guides · Periodic Table & Periodicity
Periodic Table & Periodicity — the complete deep dive
Periodicity is not a list of trends to memorise — it is one cause (effective nuclear charge felt by the outermost electrons) expressed through several observables (size, ionisation energy, electron affinity, electronegativity). Fix the cause and every trend, every anomaly, and every exam curveball becomes derivable.
Free guide — the complete guide, open to everyone.
- How to use this guide — the one-cause map
- Electronic structure engine
- Atomic & ionic size
- Ionisation energy (and its anomalies)
- Electron affinity / electron-gain enthalpy
- Electronegativity (scales) & polarising power
- Contractions & the inert-pair effect
- Diagonal relationships & anomalous 2nd period
- Worked problems (easy → hard)
- CSIR Part-C style problems
- Look-alike traps & anomalies
- Quick-recall master tables
0. How to use this guide — the one-cause map
Almost every periodic property is a downstream consequence of a single quantity: the effective nuclear chargeZeff = Z − σ experienced by the valence electrons, together with the principal quantum number n of the valence shell. Across a period Z rises but the added electrons enter the same shell and shield each other poorly, so Zeff climbs → the shell contracts and holds electrons tighter. Down a group n increases (a whole new shell, larger & well-shielded) so despite larger Z the valence electrons are farther and looser.
| Property | Reads off which factor? | Across period (L→R) | Down group |
|---|---|---|---|
| Atomic radius | n and Zeff | decreases (Zeff↑) | increases (n↑) |
| Ionisation energy | Zeff, n, subshell stability | increases (with dips) | decreases |
| Electron affinity | Zeff vs added e⁻–e⁻ repulsion | more exothermic (with dips) | generally less exothermic — except period 2 → period 3, which is MORE exothermic (2nd-row anomaly) |
| Electronegativity | IE + EA (ability to attract bonding e⁻) | increases | decreases |
| Metallic / basic character | ease of losing e⁻ (low IE) | decreases | increases |
1. The electronic structure engine ▲ top
1.1 Aufbau, the (n+l) rule, and the real exceptions
Orbitals fill in order of increasing (n + l); for equal (n+l), the lower n fills first (Madelung / Klechkowski rule). Thus 4s (n+l=4) fills before 3d (n+l=5); 4f (n+l=7) before 5d (also n+l=7, but 4f has the lower n — 4 against 5 — so 4f fills first).
The two classic exceptions (Cr and Cu) and why. The predicted [Ar]4s²3d⁴ (Cr) and [Ar]4s²3d⁹ (Cu) are beaten by [Ar]4s¹3d⁵ and [Ar]4s¹3d¹⁰. The stabilisation comes from the exchange energy of a half-filled/filled d set plus the near-degeneracy of 4s and 3d: promoting one 4s electron to complete a d⁵ or d¹⁰ set pays off. This is an exchange-energy argument, not merely "half-filled is pretty". Analogous: Mo, Ag, Au; Pd is [Kr]4d¹⁰5s⁰ (a stronger anomaly).
1.2 Shielding, penetration, and Zeff
Inner electrons screen the nuclear charge. But screening is imperfect and depends on the radial penetration of the orbital: for a given n, s > p > d > f in penetration toward the nucleus. A penetrating s electron spends time inside the inner shells, feels a larger Zeff, and is bound more tightly than a p or d electron of the same shell. This penetration ordering is the reason ns lies below np below nd in a multi-electron atom (in H, all subshells of a given n are degenerate).
1.3 Slater's rules — computing Zeff
Slater's rules give an empirical screening constant σ. Group the orbitals as: (1s)(2s,2p)(3s,3p)(3d)(4s,4p)(4d)(4f)(5s,5p)...
Slater's screening contributions (for the electron of interest)
- Electrons in groups to the right (higher) of the electron: contribute 0.
- For an ns or np electron: each other electron in the same group contributes 0.35 (but only 0.30 if the electron is in 1s); each electron in the n−1 shell contributes 0.85; each electron in n−2 or lower contributes 1.00.
- For an nd or nf electron: each other electron in the same (nd)/(nf) group contributes 0.35; every electron in any group to the left (lower) contributes 1.00.
Worked: Zeff on a 3p electron of chlorine (Z = 17)
Config: (1s²)(2s²2p⁶)(3s²3p⁵). The electron of interest is one 3p.
Same group (3s,3p): 7 electrons total, minus the one of interest = 6 × 0.35 = 2.10.
n−1 shell (2s,2p): 8 electrons × 0.85 = 6.80. Inner (1s): 2 × 1.00 = 2.00.
σ = 2.10 + 6.80 + 2.00 = 10.90. Zeff = 17 − 10.90 = 6.10.
Worked: why a 4s vs 3d electron differ (K/Ca vs Sc)
For a 3d electron, all inner electrons screen at 1.00, so a 3d electron is strongly screened by the 3s,3p core; for a 4s electron the 3s, 3p and 3d electrons all count at only 0.85 (they all lie in the n−1 = 3 shell) and 4s penetrates more. In K and Ca the 4s is lower; once Z grows (Sc onward) the 3d drops below 4s. Slater captures the crossover qualitatively — and rationalises 4s-first ionisation.
2. Atomic & ionic size ▲ top
Definitions matter. "Atomic radius" is operational: covalent radius (half the homonuclear bond length), metallic radius (half the nearest-neighbour distance in the metal), van der Waals radius (half the closest non-bonded contact — largest of the three), and ionic radius (partitioned from crystal interionic distances, Shannon radii). Always state which you mean; a common exam trap compares a vdW radius of a noble gas to covalent radii of neighbours.
Ions. A cation is always smaller than its parent atom (often a whole shell is lost, and remaining electrons feel higher Zeff). An anion is larger (extra electron–electron repulsion, lower Zeff per electron). For an isoelectronic series (same electron count), radius decreases as Z increases: e.g. N³⁻ > O²⁻ > F⁻ > Ne > Na⁺ > Mg²⁺ > Al³⁺ (all 10 electrons). This is a guaranteed exam item.
Worked: order the isoelectronic set S²⁻, Cl⁻, K⁺, Ca²⁺ by size
All are isoelectronic with Ar (18 e⁻). Radius falls as nuclear charge rises: S²⁻ (Z=16) > Cl⁻ (17) > K⁺ (19) > Ca²⁺ (20). The 20 protons of Ca²⁺ pull the same 18 electrons in tightest.
3. Ionisation energy and its anomalies ▲ top
The first ionisation energy IE₁ is the enthalpy to remove the most loosely bound electron from a gaseous atom: X(g) → X⁺(g) + e⁻. It rises across a period and falls down a group, tracking Zeff/n — but with two famous dips in every short period.
The two dips explained (the exam's favourite)
- Group 13 dip (B < Be, Al < Mg): the electron removed from B is a 2p electron. 2p is higher in energy than the filled 2s and is shielded by it, so it is easier to remove than one of Be's 2s² electrons. The filled 2s² of Be is extra-stable.
- Group 16 dip (O < N, S < P): N is 2p³ (half-filled, each p orbital singly occupied, maximal exchange stabilisation, no pairing repulsion). O is 2p⁴ — one orbital now doubly occupied; the inter-electron pairing repulsion makes that fourth 2p electron easier to remove, dropping IE below N's.
Successive IEs and the "big jump". Each successive IE rises (removing from an increasingly positive ion). A large jump appears when you break into a noble-gas core: Na has a huge IE₂/IE₁ ratio (removing from [Ne] core); Mg has the jump at IE₃; Al at IE₄. Locating the jump reveals the group (number of valence electrons).
Worked: identify the group from successive IEs
Element X has IEs (kJ/mol, schematic): 738, 1451, 7733, 10540, 13630. The huge jump is between IE₂ and IE₃.
Two electrons come off easily, the third is enormously harder → 2 valence electrons → Group 2. The figures quoted are in fact magnesium's experimental values (IE₁–IE₅ = 737.7, 1450.7, 7732.7, 10542.5, 13630 kJ/mol), so they may be quoted as they stand. Read the ratios rather than the absolute sizes: IE₂/IE₁ ≈ 2.0 (both electrons leave the 3s valence shell) but IE₃/IE₂ ≈ 5.3 (the third must be torn out of the [Ne] core).
4. Electron affinity / electron-gain enthalpy ▲ top
Electron affinity (EA) or electron-gain enthalpy ΔegH is the energy change when a gaseous atom gains an electron: X(g) + e⁻ → X⁻(g). A more negative ΔegH (larger positive EA) means the atom "wants" the electron more. The trend broadly parallels Zeff (more exothermic to the right), but carries the topic's most-loved anomaly.
Three EA features to master
- 2nd-row anomaly (F < Cl, O < S): the compact 2p subshell forces the incoming electron into a small, already crowded region → repulsion lowers the release of energy. So the maximum EA in a group is at the 3rd-period membe, not the 2nd (Cl > F, S > O, Si > C, Al > B). The one group where the maximum slips a further row down is group 15: N (−7) ≪ P (72) < As (78 kJ/mol).
- Near-zero or endothermic electron capture by N, and by Be, Mg, noble gases (ΔegH ≈ 0 or positive — equivalently EA ≈ 0 or negative; mind the sign, since this guide takes positive EA to mean exothermic): half-filled 2p³ (N) and filled ns² (Be, Mg) or full octet (Ne) resist adding an electron — the added electron must enter a higher-energy or already-stable arrangement.
- Second electron affinity is always endothermic: O⁻ + e⁻ → O²⁻ is positive (adding to an anion against repulsion). Oxide/sulfide dianions are stabilised only in lattices (lattice energy pays the cost).
5. Electronegativity & polarising power ▲ top
Electronegativity (χ) is the tendency of an atom in a molecule to attract the shared bonding electrons. Unlike IE and EA it is not a single measurable quantity but a derived scale; several definitions exist.
| Scale | Basis | Formula / idea |
|---|---|---|
| Pauling | bond energies | χA−χB ∝ √Δ, where Δ = D(A–B) − ½[D(A–A)+D(B–B)] (extra ionic resonance energy). Dimensionless; F set to 3.98. |
| Mulliken | IE and EA | χM = ½(IE + EA). Direct atomic-property definition; rescaled onto Pauling by χP ≈ χM/252.4 with χM in kJ/mol (equivalently χM(eV)/2.62); the divisor is fixed by forcing F to 3.98. |
| Allred–Rochow | electrostatic force | χAR = 0.359 Zeff/r² + 0.744 (r in Å). Force the nucleus exerts on a bonding electron at the covalent radius. |
| Allen | spectroscopic | configuration energy = average valence-electron energy from spectroscopy. |
Trends: χ increases across a period (Zeff↑) and decreases down a group (n↑). F is the most electronegative element (~4.0 Pauling); Cs/Fr the least (~0.7). Electronegativity governs bond polarity, dipole moments, acid strength of oxoacids and E–H bonds, and (via Δχ) the ionic/covalent character of a bond.
Worked: Mulliken EN of Cl from IE and EA
Cl: IE ≈ 1251 kJ/mol, EA ≈ 349 kJ/mol. χM = ½(1251 + 349) = 800 kJ/mol ≈ 8.3 eV/atom. Converting to the Pauling scale with the standard normalisation χP ≈ χM(kJ/mol)/252.4 gives 800/252.4 ≈ 3.2 — essentially the accepted Pauling value for chlorine, 3.16. The teaching point: Mulliken EN is literally the average of the atom's "grab" (EA) and "hold" (IE).
6. Contractions & the inert-pair effect ▲ top
6.1 d-block contraction (scandide contraction)
After the first d-series is filled, the poorly shielding 3d electrons cause elements of the 4th period p-block (Ga onward) to be smaller/more electronegative than a naive extrapolation predicts. Consequence: Ga is no larger than Al — by metallic/atomic radius Ga (135 pm) is in fact slightly smaller than Al (143 pm), while their covalent radii (122 vs 121 pm) are near-identical. Hence the examined group-13 atomic-radius order B < Ga < Al < In < Tl. Ge and As show the same compression.
6.2 Lanthanide contraction
Across the 4f series (Ce→Lu) the 4f electrons shield very poorly, so Zeff rises steadily and the radius shrinks by ~ the size of one period's normal expansion. The cumulative contraction (~0.17 Å: r(La³⁺) = 1.03 Å falls to r(Lu³⁺) = 0.861 Å, six-coordinate) has a huge downstream effect: the 4d and 5d transition metals of the same group become nearly identical in size (Zr≈Hf, Nb≈Ta, Mo≈W). This is why Zr/Hf and Nb/Ta are notoriously hard to separate and why 5d metals are dense and have high IEs.
6.3 Inert-pair effect
Down groups 13–15, the heavier elements increasingly favour an oxidation state two lower than the group valence because the valence ns² pair becomes reluctant to participate in bonding (poor shielding by intervening d/f electrons raises Zeff and the relativistic contraction/stabilisation of the 6s orbital deepens it). Hence Tl⁺ > Tl³⁺ stability, Pb²⁺ > Pb⁴⁺, Bi³⁺ > Bi⁵⁺. PbCl₄ and TlCl₃ are oxidising/unstable; PbCl₂ and TlCl are stable. This underlies the oxidising power of Pb(IV), Bi(V), Tl(III).
7. Diagonal relationships & the anomalous second period ▲ top
Diagonal relationships: the first element of a group resembles the second element of the next group (Li↔Mg, Be↔Al, B↔Si). Reason: moving right increases χ/polarising power while moving down decreases it — the two changes cancel diagonally, giving similar ionic potential (charge/size) and hence similar chemistry.
| Pair | Shared behaviour |
|---|---|
| Li ↔ Mg | Both form nitrides (Li₃N, Mg₃N₂), carbonates decompose to oxides, Li/Mg salts partly covalent; Li burns to normal oxide (not peroxide) like Mg. |
| Be ↔ Al | Both amphoteric oxides/hydroxides; both form covalent, polymeric/bridged halides (BeCl₂, Al₂Cl₆); both passivated; both carbides (Be₂C, Al₄C₃) are methanides, hydrolysing to CH₄ — unlike CaC₂, which gives acetylene. |
| B ↔ Si | Both metalloids; form acidic oxides; hydrides are volatile, flammable, hydrolysable; form network/covalent structures. |
Why the 2nd period is anomalous: Li–Ne have no available d orbitals and are very small. Consequences: (i) maximum covalency of 4 (N, O, F cannot expand the octet — NF₅, OF₆ don't exist, whereas PF₅, SF₆ do); (ii) strong pπ–pπ multiple bonding (N≡N, O=O, C=O) unlike heavier congeners that prefer single-bonded catenated/network forms (P₄, S₈, SiO₂ network); (iii) high electronegativity and small size drive H-bonding for N, O, F.
8. Worked problems (graded easy → hard) ▲ top
P1. Ground-state configurations & ions.
Write ground-state configurations for Cr, Cu, Fe³⁺, and identify the two 3d anomalies.
Show solution
Cr = [Ar]3d⁵4s¹; Cu = [Ar]3d¹⁰4s¹ (exchange-energy stabilised half/full d). Fe = [Ar]3d⁶4s²; Fe³⁺ = [Ar]3d⁵ (remove 4s² then one 3d — 4s ionises first). The anomalies are Cr and Cu adopting 4s¹.
P2. Isoelectronic ordering.
Arrange O²⁻, F⁻, Na⁺, Mg²⁺, Al³⁺ by increasing radius.
Show solution
All 10 electrons; radius falls as Z rises: Al³⁺ < Mg²⁺ < Na⁺ < F⁻ < O²⁻ — read left to right, that list is already the order of increasing radius, as the question asks. (Reversed, it is the order of decreasing radius.)
P3. Explain the IE dips.
Why is IE₁(Be) > IE₁(B) and IE₁(N) > IE₁(O)?
Show solution
Be→B: B loses a higher-energy, 2s-shielded 2p electron, easier than Be's filled 2s². N→O: N is stable half-filled 2p³; O's 2p⁴ has an electron-pair repulsion that lowers its IE. Both are subshell/pairing effects overriding the Zeff rise.
P4. Slater Zeff.
Compute Zeff for a 4s electron of potassium (Z=19) and for a 3d electron of a hypothetical K with the electron in 3d instead. Comment.
Show solution
4s: config (1s²)(2s²2p⁶→8)(3s²3p⁶→8)(4s¹). σ = [0 same-group others] + (n−1=3 shell: 8×0.85=6.80) + (inner 1s,2s,2p: 10×1.00=10.0) = 16.8. Zeff=19−16.8=2.2. 3d: a 3d electron sees all 18 inner electrons at 1.00 → σ=18.0, Zeff=1.0. The 4s electron feels a larger Zeff (more penetration) → lower energy → K adopts 4s¹, not 3d¹. Slater rationalises the 4s-before-3d filling.
P5. Halogen electron affinity.
Explain why EA(Cl) > EA(F) in magnitude, but F is still the more electronegative and more reactive.
Show solution
EA(Cl) is more exothermic because F's tiny 2p shell raises electron–electron repulsion on the incoming electron (2nd-row anomaly). But electronegativity reflects attraction for electrons already in a bond plus small size/high Zeff, so F is most electronegative; and F₂'s weak F–F bond (lone-pair repulsion) plus strong bonds it forms make it most reactive. Different quantities — a classic trap.
P6. Inert-pair reasoning.
Predict the more stable chloride in each pair and justify: TlCl vs TlCl₃; PbCl₂ vs PbCl₄; SnCl₂ vs SnCl₄.
Show solution
Inert-pair effect strengthens down 13–15. Tl: TlCl (Tl⁺) > TlCl₃ (strongly oxidising). Pb: PbCl₂ (Pb²⁺) > PbCl₄ (oxidising, decomposes). Sn: here Sn(IV) is still favoured (inert-pair not yet dominant so high in group 14) — SnCl₄ is a stable covalent liquid; SnCl₂ is a reducing agent. So the effect is strong for Tl/Pb (6p elements) but weak for Sn (5p). The trend intensifies down the group.
P7. Lanthanide contraction consequences.
Give three measurable consequences of the lanthanide contraction.
Show solution
(1) Zr≈Hf, Nb≈Ta, Mo≈W radii → chemically inseparable pairs; (2) the 5d elements have higher densities and IEs than expected (Os/Ir/Pt very dense; Au/Hg high IE); (3) a gradual decrease in basicity of Ln(OH)₃ across the series and the difficulty of separating adjacent lanthanides (used historically to justify ion-exchange separation).
P8. Successive IE fingerprint.
An element's successive IEs (kJ/mol) jump sharply after the 3rd electron. Which block/group, and give an example configuration.
Show solution
Three easy removals then a huge jump → 3 valence electrons → Group 13 (or a +3-favouring metal). Example Al = [Ne]3s²3p¹: IE₃ removes the last valence electron; IE₄ breaks the [Ne] core → the big jump. Locating the jump = counting valence electrons = identifying the group.
9. CSIR Part-C style problems ▲ top
Part-C rewards mechanistic reasoning about Zeff, multi-factor comparison, and quantitative estimates — not bare trend recall.
C1. Multi-factor size comparison.
Explain, using Zeff and n, why atomic radius order is Na > Mg but Mg²⁺ < Na⁺, and why K > Na even though K has more protons.
Show solution
Na > Mg: same shell (n=3), Mg has higher Zeff (one more proton, poor same-shell shielding) → smaller. Mg²⁺ < Na⁺: both are [Ne] isoelectronic; Mg²⁺ has Z=12 pulling 10 e⁻ vs Na⁺ Z=11 → Mg²⁺ smaller. K > Na: K's valence electron is in n=4 (a whole new, well-shielded shell) — the increase in n dominates the increase in Z (inner electrons cancel most of the extra protons). n change beats Z change.
C2. Anomalous EA of nitrogen.
Nitrogen has a near-zero (slightly endothermic) electron affinity while carbon and oxygen are exothermic. Explain, and predict the trend N < P for EA magnitude.
Show solution
N is 2p³ (half-filled, each 2p singly occupied). Adding an electron forces pairing in a small 2p orbital → large repulsion, so ΔegH is ≈0/positive. C (2p²) and O (2p⁴) do not sit at the special half-filled point (O already pays the pairing cost in the atom, and gains Zeff). For N vs P: P's larger 3p orbitals reduce the added-electron repulsion, so P has a more exothermic EA than N — the 2nd-row small-orbital repulsion again.
C3. Quantitative estimate.
Use Slater's Zeff and E ≈ −13.6(Zeff/n*)² eV to estimate the energy of the 2p electron in fluorine (Z=9), and compare its sign/magnitude expectation to IE(F).
Show solution
F 2p electron: same group (2s,2p) has 7 e⁻ → 6×0.35=2.10; 1s: 2×0.85=1.70. σ=3.80, Zeff=9−3.80=5.20. n=2 → n*=2. E ≈ −13.6×(5.20/2)² = −13.6×6.76 ≈ −92 eV (orbital energy; large negative → tightly bound). The observed IE(F)≈17.4 eV is far smaller in magnitude — Slater orbital energies overestimate binding because they neglect relaxation/correlation; the exercise shows the trend, not the exact IE. (The arithmetic to hold onto: σ = 6×0.35 + 2×0.85 = 2.10 + 1.70 = 3.80, so Zeff = 5.20 and E = −13.6 × 2.60² = −91.9 eV, against a measured IE(F) = 1681 kJ/mol = 17.42 eV. A factor of about five is the normal size of the Slater overestimate for a valence orbital energy, so use Slater Zeff for comparisons and ratios and never as a predicted IE.)
C4. Diagonal relationship prediction.
Predict two ways Be resembles Al rather than Mg, and give the electronic reason.
Show solution
Be(OH)₂ and Al(OH)₃ are amphoteric (Mg(OH)₂ is basic); BeCl₂ and AlCl₃ are covalent/polymeric with bridging Cl (MgCl₂ is ionic). Reason: Be²⁺ has an ionic potential (φ = q/r) close to Al³⁺ (both small, high charge density → high polarising power → covalent, amphoteric behaviour), because the rightward increase in charge and the downward increase in size cancel diagonally.
C5. Relativistic/inert-pair synthesis.
Rationalise: (a) Au is coloured and a good oxidant (Au³⁺, even Au⁻ in aurides); (b) Hg is liquid; (c) Pb(IV) oxidises but Sn(IV) does not.
Show solution
(a)/(b): relativistic contraction/stabilisation of the 6s orbital in period-6 heavy elements lowers 6s and simultaneously expands and destabilises 5d, narrows the 5d→6s gap out of the UV (where Ag's 4d→5s gap still sits) and down into the visible so that Au absorbs blue/violet and reflects yellow (Au colour), stabilises the filled 6s² (Hg's weak metallic bonding → low m.p., liquid) and makes Au electron-attracting (auride Au⁻). (c) Inert-pair effect: the 6s² pair of Pb is reluctant to bond → Pb(II) favoured, so Pb(IV) is oxidising; Sn's 5s² pair is not yet inert, so Sn(IV) is stable. All trace to poor shielding + relativistic stabilisation of the outer s pair down the group.
10. Look-alike traps & anomalies ▲ top
⚠ The classic confusions — memorise the discriminators
- IE dips are at groups 13 and 16 (Be>B, N>O), not random. Group 13 dip = s→p; group 16 dip = pairing in p⁴.
- EA maximum is period 3, not period 2: Cl > F, S > O in EA magnitude (small-orbital repulsion). But electronegativity maximum is F.
- Electronegativity ≠ electron affinity ≠ ionisation energy. EN is a bond property (derived); EA and IE are atomic (measurable). Mulliken links them: χ=½(IE+EA).
- Ionise 4s before 3d. Transition-metal cations lose ns electrons first (Fe²⁺=3d⁶, not 3d⁴4s²).
- Cation always smaller, anion always larger than the parent atom; in an isoelectronic series, higher Z → smaller.
- d-block contraction (Ga≈Al) vs lanthanide contraction (Zr≈Hf) — different causes (poor 3d vs poor 4f shielding), both shrink later elements.
- Inert-pair effect grows DOWN groups 13–15 (Tl⁺, Pb²⁺, Bi³⁺ favoured), not up. Weak for Sn, strong for Pb.
- Second electron affinity is endothermic (O⁻+e⁻→O²⁻ costs energy); dianions exist only in lattices.
- vdW radius > covalent radius — don't compare a noble gas vdW radius directly to a covalent radius and call the noble gas "huge".
- Pauling is relative (bond energies); Mulliken absolute (IE,EA); Allred–Rochow electrostatic. Know which basis each uses.
- Diagonal pairs (Li–Mg, Be–Al, B–Si) arise from matched ionic potential, not from being in the same group.
- Anomalous 2nd period: no d orbitals → max covalency 4, strong pπ–pπ bonding (N₂, O₂), unlike heavier congeners.
11. Quick-recall master tables ▲ top
11.1 Trend summary (valence-shell logic)
| Property | Across period → | Down group ↓ | Extremum |
|---|---|---|---|
| Atomic radius | decreases | increases | largest Cs/Fr; smallest He |
| IE₁ | increases (dips at 13, 16) | decreases | highest He/Ne/F; lowest Cs |
| EA / −ΔegH | more exothermic (dips) | less exothermic; max at period 3 | Cl most exothermic |
| Electronegativity | increases | decreases | F highest (~4.0) |
| Metallic/basic char. | decreases | increases | — |
| Non-metallic/acidic | increases | decreases | — |
11.2 Anomaly index
| Anomaly | Cause (one line) |
|---|---|
| Be > B, Mg > Al (IE) | removing s (filled ns²) harder than p |
| N > O, P > S (IE) | half-filled p³ stability; p⁴ pairing repulsion |
| Cl > F, S > O (EA) | compact 2p → high added-e⁻ repulsion |
| N, Be, Mg, noble gases (low EA) | half/fully-filled subshell resists extra e⁻ |
| Cr 4s¹3d⁵, Cu 4s¹3d¹⁰ | exchange-energy of half/full d + 4s/3d near-degeneracy |
| Ga ≈ Al size | d-block (scandide) contraction |
| Zr ≈ Hf, Nb ≈ Ta | lanthanide contraction (poor 4f shielding) |
| Tl⁺, Pb²⁺, Bi³⁺ stable | inert-pair effect (+ relativistic 6s stabilisation) |
| Li–Mg, Be–Al, B–Si similar | diagonal relationship (matched ionic potential φ = q/r) |
11.3 Ionic potential (φ = q/r) applications
| High-φ (q/r) cation ⇒ | Consequence |
|---|---|
| strong polarising power | covalent character (Fajans), thermal instability of oxosalts (Li₂CO₃, MgCO₃ decompose easily) |
| high charge density | strong hydration (small ions more hydrated: Li⁺ largest hydrated radius among alkali) |
| amphoterism | Be, Al oxides amphoteric |
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