Atomic Structure and Periodic Properties Questions for NDA

Structure of the Atom

An atom has a central nucleus containing protons (positive charge) and neutrons (no charge), surrounded by electrons (negative charge) moving in shells or orbitals. The atomic number Z equals the number of protons, which also equals the number of electrons in a neutral atom. Mass number A equals protons plus neutrons. Isotopes have the same atomic number but different mass numbers (same element, different neutron count).

• Atomic number Z = number of protons = number of electrons in neutral atom
• Mass number A = protons + neutrons, so neutrons = A minus Z
• Isotopes: same Z, different A. Example: Carbon-12 and Carbon-14
• Isobars: same A, different Z. Example: Argon-40 and Calcium-40
• Isotones: same number of neutrons. Example: Carbon-14 and Nitrogen-15

Electronic Configuration and Quantum Numbers

Electrons occupy orbitals arranged in subshells s, p, d, f. Each shell n has subshells up to n-1. The Aufbau principle says fill lower energy orbitals first. Hund's rule says each orbital in a subshell gets one electron before any gets two. Pauli's exclusion principle says no two electrons in an atom can have the same four quantum numbers. The four quantum numbers are: principal (n), azimuthal (l), magnetic (m), and spin (s).

• Aufbau order: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p... (fill 4s before 3d)
• s holds 2 electrons, p holds 6, d holds 10, f holds 14
• Hund's rule: half-filled and fully-filled subshells are extra stable (Cr and Cu are exceptions)
• Cr is [Ar] 3d5 4s1 NOT 3d4 4s2 — half-filled d is more stable
• Cu is [Ar] 3d10 4s1 NOT 3d9 4s2 — fully-filled d is more stable
• Principal quantum number n tells the shell; azimuthal l tells subshell shape (0=s, 1=p, 2=d, 3=f)

Periodic Table Organisation

The modern periodic table has 118 elements arranged in 7 periods (horizontal rows) and 18 groups (vertical columns). Elements are arranged in increasing atomic number. The period number tells you the number of electron shells. The group number (for main group elements) tells you the valence electrons. Blocks: s-block (groups 1-2), p-block (groups 13-18), d-block (groups 3-12, transition metals), f-block (lanthanides and actinides at bottom).

• Period 1: 2 elements. Period 2 and 3: 8 each. Period 4 and 5: 18 each. Period 6 and 7: 32 each
• s-block elements: alkali metals (group 1) and alkaline earth metals (group 2)
• p-block includes nonmetals, metalloids, halogens, noble gases
• d-block transition metals have partially filled d orbitals
• f-block: lanthanides (period 6) and actinides (period 7) — inner transition elements
• Valence electrons of p-block = group number minus 10

Periodic Trends — The Most Tested Section

Four key trends are tested repeatedly: atomic radius, ionisation energy, electron affinity, and electronegativity. Across a period (left to right), atomic radius decreases because more protons pull electrons closer with the same number of shells. Down a group (top to bottom), atomic radius increases because new shells are added. Ionisation energy is the energy needed to remove an electron from a gaseous atom. It is opposite to atomic radius — it increases across a period and decreases down a group.

• Atomic radius: decreases across period, increases down group
• Ionisation energy (IE): increases across period, decreases down group. IE of noble gases is highest in their period
• Electron affinity: energy released when an atom gains one electron. Generally increases across period (more negative), decreases down group. Fluorine is an exception — Cl has higher EA than F due to small size of F causing electron repulsion
• Electronegativity: increases across period, decreases down group. Fluorine is most electronegative element (value 4.0 on Pauling scale)
• Ionisation energy order exception: IE of group 15 > group 16 because half-filled p subshell (N, P) is extra stable
• Screening effect (shielding): inner electrons reduce effective nuclear charge felt by outer electrons, causing IE to decrease down a group

Key Properties of Periods and Groups

Group 1 (alkali metals) are highly reactive, soft, low density metals that react violently with water. Reactivity increases down the group. Group 17 (halogens) are highly reactive nonmetals; reactivity decreases down the group. Group 18 (noble gases) are inert with complete octets. Across period 3 (Na to Cl), oxides change from strongly basic (Na2O) to strongly acidic (Cl2O7) — this is an important trend for NDA.

• Metallic character decreases across a period, increases down a group
• Non-metallic character increases across a period, decreases down a group
• Oxide nature: metallic oxides are basic, non-metallic oxides are acidic, Al2O3 is amphoteric
• Diagonal relationship: Li resembles Mg, Be resembles Al, B resembles Si in properties
• Hydrogen has dual nature — placed in group 1 but also resembles group 17 halogens

• Confusing isotopes with isobars: isotopes = same atomic number (same element), isobars = same mass number (different elements). Getting this wrong in a one-liner MCQ costs you easily.
• Forgetting Cr and Cu exceptions in electronic configuration. NDA directly asks which element has configuration ending in 3d5 4s1 — answer is Cr not any other transition metal.
• Thinking fluorine has the highest electron affinity — it does NOT. Chlorine has higher electron affinity than fluorine. Fluorine is most electronegative but not highest EA.
• Confusing the trend for ionisation energy between group 14 and group 15. IE of N (group 15) is greater than O (group 16) because N has a stable half-filled 2p subshell. Many aspirants expect it to increase smoothly.
• Writing electronic configuration by filling 3d before 4s during Aufbau. Always fill 4s first (lower energy), but note that 4s electrons are removed first during ion formation.

• ARIE mnemonic for periodic trends going LEFT to RIGHT across a period: Atomic radius decreases, Reactivity of metals decreases, Ionisation energy increases, Electronegativity increases. Just remember ARIE walks left to right.
• For exceptions to ionisation energy trend, use the phrase NOBLE HALF LOVES STABILITY: noble gases peak at each period, half-filled subshells (N, P) have higher IE than the next element (O, S).
• To remember Cr and Cu exceptions: CrCu Cheat the Rule — both prefer half-filled or fully-filled d subshells over a normal Aufbau fill.
• Diagonal relationship: LiMg, BeAl, BSi — think of it as two neighbours sharing a fence diagonally on the periodic table grid.

• NDA papers regularly ask to identify the element with the highest ionisation energy in a period — it is always the noble gas. For highest electronegativity overall, always answer Fluorine.
• Electronic configuration questions in NDA often test exceptional configurations of Cr (Z=24) and Cu (Z=29). Memorise these two. A fill-in or identify question on these appears in almost every other NDA paper.
• Periodic trend comparison questions (which has larger radius, which has higher IE between two given elements) are very common and very quick to solve if you know the trends. Target these for guaranteed marks.
• Questions on isotopes, isobars, and isotones are short and factual — always appear as one-liners. Know the definitions cold and practice distinguishing them with one example each.
• NDA rarely asks deep quantum mechanics derivations. Focus on: which quantum number describes what, maximum electrons formula (2n^2), and how to write electronic configurations up to Z=30. That covers the realistic NDA scope.