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Periodic trends seem simple until the exceptions appear: boron below beryllium, oxygen below nitrogen, chlorine above fluorine. Students often memorise these as a list, which works until the exam phrases them differently. This post explains each trend and exception from the same few ideas, with measured values, so you can reason out a question you have not seen.
What drives every periodic trend?
Three things: how strongly the nucleus pulls on the outer electrons (effective nuclear charge), how far those electrons are from it (which shell they are in), and how much other electrons get in the way (shielding and repulsion). The NCERT Class 11 chapter on Classification of Elements and Periodicity in Properties builds every trend from these. Both the JEE (Main) 2026 syllabus and the NEET (UG)-2026 syllabus list atomic and ionic radii, ionisation enthalpy and electron gain enthalpy as periodic trends.
How does atomic radius change?
Across period 2, NCERT's table shows radius falling from 152 pm for lithium to 64 pm for fluorine. Electrons are added to the same shell while nuclear charge rises, so the shell is pulled in. Down group 1, radius rises from 152 pm for lithium to 262 pm for caesium, because each step adds a shell.
Why is ionisation enthalpy lower for boron and oxygen?
The general trend is upward across a period. The NIST Atomic Spectra Database gives first ionisation energies of 9.32 eV for beryllium and 8.30 eV for boron, and 14.53 eV for nitrogen and 13.62 eV for oxygen. Multiplying by 96.485 to convert eV per atom to kJ per mole gives about 899, 801, 1402 and 1314 kJ mol⁻¹.
Boron below beryllium. Beryllium loses a 2s electron; boron loses a 2p electron. NCERT explains that a 2s electron penetrates closer to the nucleus than a 2p electron, so boron's 2p electron is better shielded and easier to remove.
Oxygen below nitrogen. In nitrogen, the three 2p electrons sit in different orbitals (Hund's rule). In oxygen, two of the four 2p electrons must share an orbital, and the extra repulsion makes one of them easier to remove.
Gallium and thallium. NCERT gives first ionisation enthalpies for group 13 of 801 (B), 577 (Al), 579 (Ga), 558 (In) and 589 (Tl) kJ mol⁻¹. Gallium is not below aluminium, and thallium is above indium. Gallium has a filled 3d subshell and thallium filled 4f and 5d subshells; d and f electrons shield the nucleus poorly, so the outer electrons feel a stronger pull than the simple trend predicts.
Why does chlorine gain an electron more readily than fluorine?
NCERT's table of electron gain enthalpies gives −328 kJ mol⁻¹ for fluorine and −349 kJ mol⁻¹ for chlorine, and −141 for oxygen against −200 for sulphur. The general trend says the smaller atom should attract an added electron more strongly. But fluorine's 2p shell is so compact that an incoming electron is strongly repelled by the electrons already there. Chlorine's 3p shell is larger, so the repulsion is smaller and more energy is released overall. The same reasoning explains oxygen and sulphur.
The same table lists positive values for the noble gases (+116 for neon, +96 for argon): an added electron would have to enter a new, higher shell.
What ideas get in the way?
Taber (2003) gave 30 true-false statements about ionisation energy to over 300 college chemistry students in 17 UK institutions. Many agreed with two flawed ideas: that the nucleus has a fixed amount of "pull" shared out among its electrons, and that full shells are stable no matter what. A follow-up study of A-level students in Singapore by Tan, Taber, Goh and Chia (2005) found the same octet-rule thinking.
Both ideas give wrong reasons for right answers, and then fail on harder questions. Removing an electron always costs energy, even from a species with an "incomplete" shell. Reason in terms of attraction to the nucleus and repulsion between electrons, not what an atom "wants".
Which mistakes should you watch for?
Explaining with "wants a full shell" instead of attraction and repulsion.
Mixing up electron gain enthalpy and electronegativity. The first is a measured energy for an isolated atom; the second describes an atom in a bond.
Ionic radii. Cations are smaller than their atoms; anions are larger. For isoelectronic ions, more protons means smaller.
Assuming the trend holds down group 13. It does not, for the reason above.
A practice routine
Make a three-column table: trend, reason, exception with its reason.
Each day, predict the order for five sets, such as "first ionisation enthalpy of Na, Mg, Al", then check against NCERT or NIST values.
Explain one exception aloud in under thirty seconds.
Space your reviews; our guide to inorganic chemistry with retrieval practice shows how.
The trends rest on orbitals and their energies; our post on quantum numbers and electron configuration covers that ground. If a reason still does not make sense, the Fermi Concept Library has every concept in chemistry, each with its own lesson and its place in the map.
Tonight, explain in two lines why sulphur's electron gain enthalpy is more negative than oxygen's.
