Classification of Elements and Periodicity in Properties: common doubts, answered
The questions students ask most often about Classification of Elements and Periodicity in Properties, each with a short answer. For the full chapter, read the Classification of Elements and Periodicity in Properties notes.
About the chapter
How is Classification of Elements and Periodicity tested in NEET?
Questions usually ask you to arrange elements or ions by radius, ionization enthalpy, electron gain enthalpy or electronegativity, with the exceptions built in: boron below beryllium, oxygen below nitrogen, and chlorine ahead of fluorine. Other regular items cover IUPAC names for elements above 100, placing an element in its period and group from the configuration, and whether a given oxide is acidic, basic or amphoteric.
Mendeleev's periodic table
Read this section in the notes →Why did Mendeleev leave gaps in his periodic table?
He left gaps wherever no known element had the properties the pattern called for, and he predicted that undiscovered elements would fill them. He even forecast their properties: eka-aluminium and eka-silicon were later found as gallium and germanium, with values close to his predictions. These successful forecasts, more than the arrangement itself, persuaded chemists that his table reflected a real natural law.
Why did Mendeleev place iodine after tellurium?
Because iodine's properties match the halogens of Group VII, while tellurium resembles sulphur and selenium, even though tellurium has the higher atomic weight. Mendeleev chose chemical similarity over strict weight order. The modern periodic law later proved him right: arranged by atomic number, tellurium (52) properly comes before iodine (53), so the apparent anomaly disappears.
Modern periodic law and the long form
Read this section in the notes →What is the difference between Mendeleev's periodic law and the modern periodic law?
Mendeleev made properties a periodic function of atomic weight, while the modern law makes them a periodic function of atomic number. Moseley's X-ray studies showed that atomic number, the charge on the nucleus, is the more basic property. Using it removes the old misfits, such as argon before potassium, and gives all isotopes of an element one place because they share the same Z.
Naming elements with Z above 100
Read this section in the notes →How do you write the IUPAC name of an element with atomic number above 100?
Join the numerical roots for each digit and finish with -ium: nil 0, un 1, bi 2, tri 3, quad 4, pent 5, hex 6, sept 7, oct 8, enn 9. So element 120 is unbinilium, symbol Ubn, and element 101 is unnilunium, Unu. These provisional names are used only until IUPAC approves a permanent name for a newly made element.
Electronic configuration, periods and groups
Read this section in the notes →How do you find the period and group of an element from its electronic configuration?
The period equals the principal quantum number of the outermost shell. For the group, count valence electrons in the s-block, add 10 to the outer s plus p electrons in the p-block, and add the (n − 1)d and ns electrons in the d-block. So [Ne] 3s² 3p⁴ is in period 3 and group 10 + 6 = 16, which is sulphur.
Why do the periods have 2, 8, 8, 18, 18 and 32 elements?
Each period fills a new shell, and its length is twice the number of orbitals being filled in it. Period 1 fills only 1s, giving 2 elements; periods 2 and 3 fill one s and three p orbitals, giving 8; periods 4 and 5 add five d orbitals, giving 18; period 6 also fills seven f orbitals, giving 32. Each orbital takes two electrons.
The s, p, d and f blocks
Read this section in the notes →Why is helium placed in group 18 if it is an s-block element?
Helium's configuration, 1s², makes it an s-block element, but that single shell is completely full. A full shell makes it chemically inert in the same way as neon and argon, so it is grouped with the noble gases in group 18. Here the table follows chemical behaviour rather than simply the type of orbital that received the last electron.
Why are the lanthanoids and actinoids placed separately below the periodic table?
They are set apart to keep the table a practical width and to keep elements of similar behaviour together. The lanthanoids, cerium to lutetium, and the actinoids, thorium to lawrencium, fill 4f and 5f orbitals that lie deep inside the atom and hardly change outer-shell chemistry, so members of each series are much alike. Placing them inline would widen the table to 32 columns.
Metals, non-metals and metalloids
Read this section in the notes →What are metalloids, with examples?
Metalloids, also called semi-metals, are elements whose properties fall between those of metals and non-metals; silicon, germanium, arsenic, antimony and tellurium are examples. They lie along the zigzag line that divides metals from non-metals in the p-block. Silicon and germanium look shiny but are brittle and conduct electricity only modestly, which makes them useful as semiconductors.
Atomic and ionic radii
Read this section in the notes →Why does atomic radius decrease across a period but increase down a group?
Across a period, electrons go into the same shell while the nuclear charge keeps rising, and electrons in one shell screen each other poorly, so the outer shell is drawn inwards. Down a group, every new period adds a shell farther from the nucleus, and the extra inner electrons shield the outer ones, so the atom grows even though the nuclear charge also increases.
Why is a cation smaller and an anion larger than its parent atom?
A cation has lost electrons, so the unchanged nuclear charge acts on fewer electrons and holds them closer, and often a whole outer shell is gone. An anion has gained electrons that repel one another while the nuclear charge stays the same, so its electron cloud expands. That is why Na⁺ (95 pm) is much smaller than Na (186 pm) and F⁻ (136 pm) larger than F (64 pm).
How do you arrange isoelectronic ions in order of size?
Isoelectronic species carry the same number of electrons, so their size depends on nuclear charge alone: more protons pull the same electron cloud in more tightly. For the 10-electron set the order is O²⁻ > F⁻ > Na⁺ > Mg²⁺, with nuclear charges of 8, 9, 11 and 12. A quick check is that the most negatively charged ion is the largest.
Ionization enthalpy
Read this section in the notes →Why is the ionization enthalpy of boron lower than that of beryllium?
Boron's outermost electron sits in a 2p orbital, whereas beryllium's sits in 2s. A 2s electron penetrates closer to the nucleus and is held more firmly, and boron's 2p electron is also screened by its filled 2s pair. It therefore feels a smaller effective nuclear charge and is removed more easily, even though boron has one more proton than beryllium.
Why is the ionization enthalpy of oxygen lower than that of nitrogen?
Nitrogen's three 2p electrons occupy separate orbitals, giving a stable half-filled set. Oxygen's fourth 2p electron has to pair up in an orbital that already holds one electron, and the repulsion between the two makes one of them easier to pull away. So oxygen dips below nitrogen, and the period 2 order is Li < B < Be < C < O < N < F < Ne.
Why does ionization enthalpy decrease down a group?
Going down a group, the outermost electron lies in a shell farther from the nucleus and is screened by more filled inner shells. The rise in nuclear charge is outweighed by this extra distance and shielding, so less energy is needed to remove the electron. That is why caesium gives up its outer electron far more readily than lithium and is a much more reactive metal.
Electron gain enthalpy and electronegativity
Read this section in the notes →Why does chlorine have a more negative electron gain enthalpy than fluorine?
Chlorine releases more energy on accepting an electron, −349 kJ mol⁻¹ against −328 kJ mol⁻¹ for fluorine. In fluorine the extra electron must enter the compact n = 2 shell, already crowded with seven electrons, so it meets strong repulsion. In chlorine it enters the roomier n = 3 shell, where repulsion is smaller, so the addition is more favourable even though fluorine is more electronegative.
Why do noble gases have positive electron gain enthalpy?
A noble gas already has a completely filled valence shell, so an extra electron would have to begin a new, higher shell, where it is held only weakly. Energy must be supplied to force it on, making the electron gain enthalpy positive; for neon it is about +116 kJ mol⁻¹. This is also why noble gases do not normally form negative ions.
What is the difference between electron gain enthalpy and electronegativity?
Electron gain enthalpy is the measured energy change when an isolated gaseous atom takes up an electron, whereas electronegativity describes how strongly an atom in a bond pulls the shared electron pair towards itself. The first is an experimental value for free atoms; the second is a relative scale, such as Pauling's, on which fluorine tops the list at 4.0. Chlorine leads the first, fluorine the second.
Valence and the anomalous second period
Read this section in the notes →Why do second period elements behave differently from the rest of their group?
Lithium to fluorine are unusually small, have a high charge-to-radius ratio and high electronegativity, and have no d orbitals in their valence shell. Lacking d orbitals limits their covalency to four, so boron stops at BF₄⁻ while aluminium forms AlF₆³⁻. The same size effects explain diagonal relationships, such as lithium resembling magnesium and beryllium resembling aluminium.
Periodic trends and chemical reactivity
Read this section in the notes →Why do oxides become more acidic across a period?
Across a period, elements shift from strongly metallic to non-metallic, and their oxides shift from basic to acidic. On the left, Na₂O dissolves to give a strong base; in the middle, Al₂O₃ is amphoteric and reacts with both acids and bases; on the right, Cl₂O₇ gives a strong acid. Down a group the opposite happens, and oxides grow more basic as metallic character rises.
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