Periodic classification and trends for NEET

Periodic classification looks like a chapter of facts, but it is really a chapter of reasons. Every trend follows from two things: how strongly the nucleus pulls on the outer electrons (effective nuclear charge), and how far away those electrons are (the shell number). Once you can argue from those two ideas, you can work out most orders instead of memorising them, and the exceptions become easy to remember because each has a clear reason.

NEET questions here are mostly "arrange in order" and "which is the exception" questions, plus a few on locating an element and naming new ones.

From Mendeleev to the modern periodic law

• Döbereiner's triads and Newlands' law of octaves were early attempts that worked only for some elements.
• Mendeleev arranged elements by atomic mass, but he let chemical properties override mass where needed, left gaps for undiscovered elements (such as eka-aluminium, later gallium) and predicted their properties.
• Moseley showed that the atomic number is the more basic property. The modern periodic law says that the properties of elements are a periodic function of their atomic numbers.

The table has 7 periods and 18 groups. A period starts a new shell, and its length is the number of electrons that shell's filling order can take: 2, 8, 8, 18, 18, 32, 32.

Locating an element from its configuration

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

IUPAC names for elements above 100

Until a permanent name is approved, an element is named from the digits of its atomic number, using these roots, followed by "ium":

The symbol takes the first letter of each root. So Z = 120 is unbinilium (Ubn) and Z = 109 is unnilennium (Une). One spelling rule: the final "i" of "bi" and "tri" is dropped before "ium", so Z = 102 is unnilbium (Unb), not unnilbiium.

Atomic and ionic radius

• Across a period, the radius decreases: the nuclear charge rises, while the electrons enter the same shell and shield each other poorly.
• Down a group, it increases: each step adds a new shell.
• A cation is smaller than its atom (fewer electrons, same pull, often a shell lost). An anion is larger (more electrons, more repulsion, same pull).
• Isoelectronic species have the same number of electrons. The one with more protons pulls harder and is smaller.

Ionisation enthalpy

Ionisation enthalpy is the energy needed to remove the most loosely held electron from an isolated gaseous atom in its ground state. It is always positive, and each successive ionisation enthalpy is larger than the one before.

It increases across a period and decreases down a group, with two exceptions in each of periods 2 and 3:

So the period 2 order is Li < B < Be < C < O < N < F < Ne. Approximate values in kJ mol⁻¹: Li 520, B 801, Be 899, C 1086, O 1314, N 1402, F 1681, Ne 2080.

A sudden large jump in successive ionisation enthalpies shows that the next electron must come from an inner, noble-gas shell. The number of electrons removed before the jump is the number of valence electrons.

Electron gain enthalpy

Electron gain enthalpy (ΔegH) is the enthalpy change when a gaseous atom gains an electron. For most elements it is negative (energy is released).

• It generally becomes more negative across a period and less negative down a group.
• Noble gases have large positive values: the added electron must enter a new shell.
• Cl is more negative than F (about −349 against −328 kJ mol⁻¹). F is so small that the incoming electron is strongly repelled by the electrons already crowded in its 2p subshell.
• For the same reason, S is more negative than O.
• Group 2 elements and nitrogen have values that are close to zero or positive, because their subshells are filled or half-filled.

Electronegativity, metallic character and oxides

Electronegativity is the ability of an atom in a molecule to attract the shared electrons towards itself. It is not a measurable energy like the two above. On the Pauling scale, fluorine is highest (4.0), followed by oxygen. It increases across a period and decreases down a group.

Metallic character follows the opposite trend: it decreases across a period and increases down a group.

This is reflected in the oxides. Across period 3, they change from strongly basic to strongly acidic: Na₂O and MgO are basic, Al₂O₃ is amphoteric, and SiO₂, P₄O₁₀, SO₃ and Cl₂O₇ are acidic, increasingly so.

Anomalous first members and the diagonal relationship. The first element of each group (Li, Be, B, …) differs from the rest of its group because of its small size, high charge density and lack of d orbitals. Some of them resemble the element diagonally below and to the right: Li and Mg, Be and Al, B and Si.

Worked NEET-style questions

Example 1: locating an element

Find the period, block and group of the element with Z = 35.

• Configuration: [Ar] 3d¹⁰ 4s² 4p⁵.
• Highest n = 4: period 4. Last electron enters 4p: p-block.
• Group = 10 + (2 + 5) = 17. The element is bromine.

Example 2: an order with exceptions

Arrange Na, Mg, Al and Si in increasing first ionisation enthalpy.

• The general rise across the period gives Na < Mg < Al < Si.
• But Al's 3p electron is easier to remove than Mg's paired 3s electron, so Al < Mg.
• Order: Na < Al < Mg < Si (about 496, 577, 738 and 786 kJ mol⁻¹).

Example 3: isoelectronic ions

Arrange N³⁻, O²⁻, F⁻, Na⁺, Mg²⁺ and Al³⁺ by decreasing radius.

• All have 10 electrons. Their nuclear charges are 7, 8, 9, 11, 12 and 13.
• More protons, smaller ion: N³⁻ > O²⁻ > F⁻ > Na⁺ > Mg²⁺ > Al³⁺.

Example 4: reading successive ionisation enthalpies

An element has successive ionisation enthalpies of about 738, 1451, 7733 and 10,540 kJ mol⁻¹. Which group is it in?

• The big jump comes between the second (1451) and third (7733) values.
• So it has 2 valence electrons: a group 2 element. These are the values for magnesium.

Practice MCQs

1. Which has the higher first ionisation enthalpy? (a) N (b) O (c) both equal (d) cannot be said
2. The element with the most negative electron gain enthalpy is: (a) F (b) Cl (c) Br (d) I
3. The largest of Na⁺, Mg²⁺, F⁻ and O²⁻ is: (a) Na⁺ (b) Mg²⁺ (c) F⁻ (d) O²⁻
4. An amphoteric oxide is: (a) Na₂O (b) MgO (c) Al₂O₃ (d) SO₃
5. An element with configuration [Ne] 3s² 3p³ is in: (a) period 3, group 13 (b) period 3, group 5 (c) period 2, group 15 (d) period 3, group 15
6. The IUPAC temporary symbol for Z = 109 is: (a) Une (b) Unn (c) Uno (d) Uue
7. A pair showing a diagonal relationship is: (a) Li and Mg (b) Li and Na (c) Be and Mg (d) B and Al
8. The correct order of first ionisation enthalpy is: (a) Be < B < C (b) B < Be < C (c) C < B < Be (d) B < C < Be

Answers

1. (a) N has a stable half-filled 2p subshell.
2. (b) F is small, so electron–electron repulsion lowers its value below Cl's.
3. (d) All have 10 electrons; O²⁻ has the fewest protons (8).
4. (c) Al₂O₃ reacts with both acids and bases.
5. (d) n = 3; group = 10 + (2 + 3) = 15.
6. (a) 1, 0, 9 gives un-nil-enn-ium: Une.
7. (a) Li and Mg sit diagonally and have similar charge-to-size ratios.
8. (b) B (801) < Be (899) < C (1086) kJ mol⁻¹.

What to do next

• Draw the skeleton table and fill in blocks, group numbers and the first 36 elements from memory.
• Write the four ionisation enthalpy exceptions and the two electron gain exceptions, each with its reason.
• Practise ten "arrange in order" questions, arguing each from nuclear charge and shell number before you answer.
• Revisit structure of the atom if the configurations feel slow, then move on to chemical bonding.