Periodic Properties
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Direct answer
The modern periodic table arranges elements by atomic number into 18 groups and 7 periods, and properties repeat because electron configurations repeat. Across a period, rising nuclear charge with a constant shell count pulls electrons in: radius falls while ionisation enthalpy, electronegativity and (generally) electron gain enthalpy rise; down a group new shells dominate and reverse each trend. The high-yield skill is knowing both the trends and the handful of exceptions JEE tests.
What you must remember
- Cations are smaller, anions larger than the parent atom; in isoelectronic series radius falls with rising charge — N3- > O2- > F- > Na+ > Mg2+ > Al3+. Noble-gas radii are van der Waals radii, larger than neighbouring covalent radii.
- Ionisation enthalpy rises across a period, falls down a group; sodium's second ionisation enthalpy is huge (electron torn from the neon core).
- Classic anomalies: Be > B (filled 2s2 more penetrating than 2p1) and N > O (half-filled p3 stability); the pattern repeats Mg > Al and P > S.
- Electron gain enthalpy: chlorine is more negative than fluorine, whose compact 2p shell repels the incoming electron; order Cl > F > Br > I; noble gases positive.
- Electronegativity (Pauling) rises across a period, falls down a group; fluorine tops the scale. Unlike ionisation enthalpy, it carries no energy units.
- Period 3 oxides: Na2O strongly basic, MgO basic, Al2O3 amphoteric, SiO2 weakly acidic, P4O10 and SO3 strongly acidic.
- Lanthanoid contraction — poor 4f shielding — shrinks the 5d series, making Zr and Hf nearly identical.
- Diagonal relationships: Li-Mg, Be-Al, B-Si, because opposing trends partly cancel along a diagonal.
Common confusion
The recurring confusion is treating the three "pull" quantities as one: ionisation enthalpy removes an electron from a gaseous atom, electron gain enthalpy is the energy change on adding one, and electronegativity is a bonded-atom tendency, not an energy. Students expect monotonic ionisation enthalpy across a period and are caught by Be-B and N-O — penetration and half-filled stability outrank plain nuclear charge. The F versus Cl case trips many: fluorine's pull loses to inter-electronic repulsion in its cramped 2p subshell.
Exam-focused takeaway
JEE Main tests trends as statements: ordering elements by radius or ionisation enthalpy, identifying the amphoteric oxide, matching property to definition. JEE Advanced builds comparison chains — second versus third ionisation enthalpies, isoelectronic size ordering, successive electron gains (the second always endothermic) and consequence questions like why thallium prefers Tl+. Justify each ordering in one line from nuclear charge, shielding or subshell stability; the reasoning survives unfamiliar cases that memory does not.
Frequently asked questions
Why is beryllium's ionisation enthalpy greater than boron's?
Beryllium loses an electron from the penetrating, filled 2s orbital; boron's comes from the higher-energy 2p, so boron ionises more easily.
Why is chlorine's electron gain enthalpy more negative than fluorine's?
Fluorine's tiny 2p shell creates electron-electron repulsion that cancels part of the energy released; chlorine's roomier 3p accepts the electron more comfortably.
What is the lanthanoid contraction?
Steady radius decrease across the 4f series from poor f-orbital shielding; it makes Zr and Hf almost identical and hard to separate.
Which Period 3 oxide is amphoteric?
Al2O3 — reacting as base with HCl and as acid with NaOH, between basic Na2O/MgO and acidic SiO2 and beyond.
Why is a cation smaller than its parent atom?
Electron loss shrinks the cloud at constant nuclear charge, often removing an entire shell, as in Na+ versus Na.
What is a diagonal relationship?
Similarity of the first element of a group with the second of the next — Li with Mg, Be with Al — because charge density and electronegativity changes cancel along the diagonal.