Periodic properties for JEE Main

Periodicity is a short chapter that pays off twice. It brings direct questions of its own, usually "arrange in order" or "which has the highest", and it supplies the reasoning for p-block, d-block and bonding questions later. The trends themselves take an hour to learn. The marks are in the exceptions, and every exception has a reason you can state in one line.

The official 2026 syllabus lists the modern periodic law and the present form of the table, the s, p, d and f blocks, and periodic trends in atomic and ionic radii, ionisation enthalpy, electron gain enthalpy, valence, oxidation states and chemical reactivity. Electronegativity is listed in the bonding unit and is covered here as well.

One idea behind every trend

Almost every trend follows from the effective nuclear charge felt by the outer electrons and the distance of those electrons from the nucleus.

• Across a period, the nuclear charge rises by one each step while the added electron goes into the same shell and shields poorly. Outer electrons are pulled in more tightly.
• Down a group, a new shell is added each step. The outer electrons are farther away and better shielded, so they are held less tightly.

If you can explain a trend in these terms, you can usually explain its exceptions too.

The modern periodic table

The modern periodic law says that properties are a periodic function of atomic number (Moseley's work showed atomic number, not atomic mass, is the basis). The table has 18 groups and 7 periods.

The period number is the highest principal quantum number in the configuration.

IUPAC names for Z above 100 use numerical roots: 0 nil, 1 un, 2 bi, 3 tri, 4 quad, 5 pent, 6 hex, 7 sept, 8 oct, 9 enn, followed by -ium. So Z = 104 is unnilquadium (Unq) and Z = 120 is unbinilium (Ubn). These are temporary names until a permanent name is approved.

Atomic and ionic radii

Two points JEE likes:

• Ga is slightly smaller than Al. The ten 3d electrons shield the nucleus poorly, so gallium's outer electrons feel a larger effective charge.
• Noble gas radii are van der Waals radii, which are larger than covalent radii. Do not include them in a covalent-radius trend across a period.

Ionisation enthalpy

Ionisation enthalpy is the energy needed to remove the most loosely held electron from an isolated gaseous atom. It increases across a period and decreases down a group, and successive ionisation enthalpies always increase: IE₁ < IE₂ < IE₃.

The exceptions come from subshell structure:

• Be > B and Mg > Al. B and Al lose a p electron, which is higher in energy and shielded by the filled s subshell.
• N > O and P > S. N and P have half-filled p subshells. O and S lose a paired electron, which is easier to remove because of repulsion within the pair.

So the order in period 2 is Li < B < Be < C < O < N < F < Ne.

A large jump between successive ionisation enthalpies shows that a noble gas core has been reached, which identifies the group.

Electron gain enthalpy

Electron gain enthalpy is the enthalpy change when a gaseous atom gains an electron. A more negative value means the atom accepts an electron more readily.

• It generally becomes more negative across a period and less negative down a group.
• Cl is more negative than F (about −349 and −328 kJ/mol). The fluorine 2p subshell is so compact that the added electron feels strong repulsion. The same applies to S being more negative than O.
• Noble gases and elements with stable configurations (Be, Mg, N) have positive or near-zero values.
• The second electron gain enthalpy of oxygen is positive: adding an electron to O⁻ means pushing it against a negative ion.

Electronegativity, valence and reactivity

Electronegativity is the tendency of an atom in a bond to attract the shared pair. On the Pauling scale F is highest (4.0), followed by O (3.5), then N and Cl (3.0). It rises across a period and falls down a group.

Valence of the main-group elements equals the number of valence electrons, or 8 minus that number. In the heavier p-block elements, the inert pair effect makes the lower oxidation state more stable: Tl⁺ over Tl³⁺, Pb²⁺ over Pb⁴⁺, Bi³⁺ over Bi⁵⁺.

Metallic character and oxides. Metallic character falls across a period and rises down a group. Oxides follow: across period 3, Na₂O and MgO are basic, Al₂O₃ is amphoteric, and SiO₂, P₄O₁₀, SO₃ and Cl₂O₇ are acidic, with acidity increasing to the right.

Worked problems

Problem 1: ordering ionisation enthalpies. Arrange B, C, N and O in order of increasing first ionisation enthalpy.

The general trend is B < C < N < O, but O is lower than N because N has a half-filled 2p subshell.
Order: B < C < O < N.

Problem 2: reading successive ionisation enthalpies. An element has IE₁ to IE₄ of 738, 1451, 7733 and 10,540 kJ/mol. Which group is it in?

The big jump is between IE₂ and IE₃, so the third electron comes from a noble gas core. The element has two valence electrons: group 2 (these values are magnesium's).

Problem 3: isoelectronic radii. Arrange S²⁻, Cl⁻, K⁺ and Ca²⁺ by size.

All have 18 electrons. Nuclear charges are 16, 17, 19 and 20.
Order: S²⁻ > Cl⁻ > K⁺ > Ca²⁺.

Problem 4 (numerical answer): energy from ionisation enthalpy. The first ionisation enthalpy of sodium is 496 kJ/mol. How much energy, in kJ, is needed to convert 2.3 g of gaseous sodium atoms into Na⁺ ions? Round to the nearest integer.

Moles of Na = 2.3/23 = 0.1.
Energy = 0.1 × 496 = 49.6 kJ, which rounds to 50.

Traps to watch

• Treating trends as absolute. Learn the exceptions with their reasons: Be/B, N/O, Mg/Al, P/S and Cl/F.
• Electron gain enthalpy signs. "Most negative" is the same as "releases most energy".
• Confusing electronegativity with electron gain enthalpy. The first is about an atom in a bond; the second is a measured energy for an isolated atom.
• Placing f-block elements. Lanthanoids and actinoids belong to group 3.

Practice set

1. Which has the largest radius: Na⁺, F⁻ or O²⁻?
2. Which has the higher first ionisation enthalpy, Mg or Al?
3. Which of F, Cl, Br and I has the most negative electron gain enthalpy?
4. Give the IUPAC temporary name and symbol for Z = 115.
5. Which is the most acidic oxide: Na₂O, MgO, Al₂O₃ or Cl₂O₇?
6. State the block, period and group of the element with Z = 29.
7. Electron gain enthalpy of Cl is −349 kJ/mol. How much energy is released when 3.55 g of gaseous Cl atoms become Cl⁻ ions?
8. Why is the first ionisation enthalpy of oxygen lower than that of nitrogen?

Answers

1. O²⁻: all three have 10 electrons, and O has the lowest nuclear charge.
2. Mg: Al loses a 3p electron, which is easier to remove than Mg's 3s electron.
3. Cl.
4. Un-un-pent-ium: ununpentium, Uup (now named moscovium, Mc).
5. Cl₂O₇.
6. Copper, [Ar]3d¹⁰4s¹: d-block, period 4, group 11.
7. 0.1 mol × 349 = 34.9 kJ released.
8. N has a half-filled 2p³ subshell. In O (2p⁴), one 2p orbital holds a pair, and repulsion within the pair makes one electron easier to remove.

What to do next

• Write the period 2 and period 3 ionisation enthalpy orders from memory, with the reason for each exception.
• Practise ten "arrange in order" questions on radii, including isoelectronic sets.
• Revise atomic structure if configurations still feel slow.
• Move on to chemical bonding, where electronegativity and size decide bond character.