Periodic Table — Trends, Properties & Memory Tricks

The periodic table is not just a chart you memorise for the sake of it. Every trend in it has a reason — and once you understand the "why", the "what" falls into place automatically. This is one of the highest-scoring topics in CBSE Class 11, JEE Main, and NEET, with consistent weightage every year.

The Modern Periodic Table — Structure First

Mendeleev's periodic table arranged elements by atomic mass. The modern periodic table, proposed by Moseley (1913), arranges them by atomic number — and that small shift fixed almost every anomaly Mendeleev's table had.

The modern periodic table has 7 periods (horizontal rows), 18 groups (vertical columns), and 118 elements currently.

The Four Blocks

Elements are classified into four blocks based on which subshell receives the last (differentiating) electron.

Memory Tricks for s-Block Groups

Group 1 (Alkali metals): H Li Na K Rb Cs Fr Mnemonic: "Hi, Little Naughty Kids Run Crazy Fast"

Group 2 (Alkaline earth metals): Be Mg Ca Sr Ba Ra Mnemonic: "Beta Mange Car Scooter Baap Raazi" — classic Indian classroom trick that never fails.

Periodic Trends — The Core of This Chapter

1. Atomic Radius

Atomic radius is the measure of the size of an atom. We use covalent radius (half the distance between two bonded same-element atoms) as the standard for non-metals, and metallic radius for metals.

Across a period (left to right): atomic radius decreases.

Nuclear charge increases as we move across a period, but electrons are added to the same shell — so shielding stays roughly constant. The stronger nuclear pull contracts the electron cloud inward.

Down a group (top to bottom): atomic radius increases.

Each new period adds a new electron shell. More shells means more shielding and greater distance from the nucleus — the expanded size wins over the increased nuclear charge.

Key sequence for Period 3 (decreasing radius): Na > Mg > Al > Si > P > S > Cl > Ar

Key sequence for Group 1 (increasing radius): Li < Na < K < Rb < Cs

2. Ionization Energy (IE)

Ionization energy is the minimum energy required to remove the outermost electron from a gaseous atom in its ground state. Higher IE = more tightly held electron = harder to remove.

Across a period: IE generally increases.

Nuclear charge increases and atomic radius decreases across a period. The outer electron is held more tightly by the nucleus, so more energy is needed to remove it.

Down a group: IE decreases.

Atomic radius increases and inner electrons provide more shielding. The outer electron is farther from the nucleus and more screened — it requires less energy to remove.

Critical IE Anomalies — These Come in Every Exam

IE(Be) > IE(B): Beryllium has configuration [He] 2s² — a fully filled 2s subshell with extra stability. Boron's differentiating electron is in 2p¹, which is higher in energy and easier to remove. Despite having more nuclear charge, B has lower IE than Be.

IE(N) > IE(O): Nitrogen has [He] 2s² 2p³ — a half-filled 2p subshell with extra stability (Hund's rule symmetry). Oxygen has [He] 2s² 2p⁴ — one pair of electrons is forced into the same orbital, causing electron–electron repulsion. This repulsion makes one electron easier to remove from oxygen.

Second Ionization Energy (IE₂): Always greater than IE₁ — you're removing an electron from a positively charged ion. The sharp jump in IE₂ of Na is famous: Na⁺ has the noble gas configuration of Ne (2,8), making the second removal extraordinarily difficult. This pattern helps identify group membership from IE data.

3. Electron Affinity (EA)

Electron affinity is the energy change when a gaseous atom gains one electron to form an anion. A more negative EA means the atom more readily accepts an electron.

Across a period: EA generally becomes more negative (increases in magnitude). Down a group: EA generally becomes less negative (decreases in magnitude).

Key EA Anomalies

EA(F) less negative than EA(Cl): This is the most-tested EA fact in exams. Fluorine should theoretically have the highest EA (most electronegative element), but chlorine's EA is more negative. Fluorine's 2p orbitals are extremely compact — adding an electron into the already dense 2p subshell creates significant electron–electron repulsion. Chlorine's 3p orbitals are larger and can accommodate the extra electron with less repulsion.

EA of noble gases ≈ 0: Fully filled valence shells — no room or desire for another electron.

EA of alkaline earth metals ≈ 0 or positive: Their s-subshells are full; the next electron must go into a higher-energy p subshell — energetically unfavourable.

4. Electronegativity (EN)

Electronegativity is the ability of an atom in a covalent bond to attract shared electrons toward itself. The Pauling scale is standard.

Across a period: EN increases. Down a group: EN decreases.

Fluorine (EN = 4.0) has the highest electronegativity of all elements. Caesium (or Francium) has the lowest.

Bond character from EN difference:

• EN difference > 1.7 → Ionic bond (e.g., NaCl: 3.0 − 0.9 = 2.1)
• EN difference 0.4–1.7 → Polar covalent (e.g., HCl: 3.0 − 2.1 = 0.9)
• EN difference < 0.4 → Non-polar covalent (e.g., H₂: 0)

5. Metallic and Non-Metallic Character

Metallic character (tendency to lose electrons) decreases across a period and increases down a group — directly mirroring IE trends.

Non-metallic character (tendency to gain electrons) increases across a period and decreases down a group — mirroring EA and EN trends.

The staircase line on the periodic table separates metals (left) from non-metals (right). Metalloids (Si, Ge, As, Sb, Te) lie along this boundary and have intermediate properties — crucial for NEET.

6. Valency and Oxidation States

For main group elements:

• Group 1: Valency = 1
• Group 2: Valency = 2
• Groups 13–17: Common valency = Group number − 10 (or 18 − Group number, the "alternative" valency)

For transition metals, variable valency is standard because both 4s and 3d electrons can participate in bonding.

The d-Block — Transition Metals (Groups 3–12)

Transition metals have incompletely filled d-subshells in their elemental state or in at least one stable oxidation state.

Key properties:

• Variable oxidation states: Fe shows +2 and +3; Mn shows +2, +3, +4, +6, +7
• Coloured ions: Due to d–d electron transitions (e.g., Cu²⁺ is blue, Cr³⁺ is green, MnO₄⁻ is purple)
• Catalytic activity: Fe in Haber process, V₂O₅ in Contact process, MnO₂ in thermal decomposition of KClO₃
• Complex formation: Empty d orbitals accept lone pairs from ligands

Atomic radius across d-block: Does not decrease smoothly. As d electrons are added, they shield each other reasonably well — the radius stays nearly constant from Sc to Cu, then increases slightly at Zn.

Lanthanide Contraction

In lanthanides (Ce to Lu), electrons fill the 4f subshell. The 4f orbitals have complex angular distributions and shield outer electrons very poorly from nuclear charge. So as nuclear charge increases across the lanthanides, the electron cloud contracts progressively.

Consequence: Elements immediately after the lanthanides (Period 6 d-block: Hf, Ta, W, Re…) have almost the same atomic radii as their Period 5 counterparts (Zr, Nb, Mo…). Zr and Hf are virtually identical in size (both ~160 pm) — making them the hardest pair of elements to separate in industry.

Full Trend Comparison Table

5 Common Mistakes Students Make

Mistake 1: Confusing period trend with group trend for atomic radius. Atomic radius decreases going across a period (left to right). It increases going down a group. Students frequently reverse one of these.

Mistake 2: Thinking fluorine has the highest electron affinity. Chlorine has a more negative EA than fluorine. F's compact 2p orbitals cause electron–electron repulsion when a new electron tries to enter. This is one of the most reliable trick questions in NEET MCQs.

Mistake 3: Ignoring IE anomalies (Be > B, N > O, P > S). These anomalies arise from extra stability of fully-filled or half-filled subshells. Missing these means losing marks on questions that appear almost every year.

Mistake 4: Applying s/p block trends directly to d-block. Transition metal radii don't decrease smoothly across a period. Their IE and EN trends are also less pronounced. Don't extrapolate main-group rules blindly to transition metals.

Mistake 5: Treating electronegativity and electron affinity as the same thing. EN is about shared electrons in a bond (relative concept). EA is about an isolated gaseous atom gaining an electron (measurable energy). They often trend the same way, but the F vs Cl anomaly shows they can diverge.

Practice Questions

Q1. Which has the larger atomic radius — Na or Mg? Explain why.

Q2. Why does the ionization energy of oxygen fall below that of nitrogen?

Q3. Arrange in decreasing order of atomic radius: Na, K, Cl, Mg.

Q4. The second ionization energy of Na is about 9 times its first. Why?

Q5. Why is the ionic radius of Cl⁻ greater than the atomic radius of Cl?

Q6. Why is EA(Cl) more negative than EA(F) even though F is more electronegative?

Q7. What is lanthanide contraction? Give one important consequence.

Q8. Which has higher first ionization energy — P or S? Explain.

Frequently Asked Questions

How many elements are in Period 3?

Eight: Na, Mg, Al, Si, P, S, Cl, Ar. Period 3 contains only s-block and p-block elements — the d-block starts in Period 4 with Sc (Z = 21).

Why does Period 6 have 32 elements?

Period 6 fills 6s (2 electrons), 4f (14 electrons), 5d (10 electrons), and 6p (6 electrons) subshells. 2 + 14 + 10 + 6 = 32 elements. The f-block makes its first appearance in Period 6 (lanthanides, Ce–Lu) and Period 7 (actinides, Th–Lr).

What is effective nuclear charge (Z_eff)?

Z_eff = Z − σ, where σ is the shielding constant from inner electrons. It represents the net nuclear charge felt by an outer electron after accounting for repulsion by inner electrons. Slater's rules give an approximate σ. Higher Z_eff → smaller radius, higher IE, higher EN. This is the single concept that explains almost all periodic trends.

Why are Group 2 elements harder to ionise than Group 1 in the same period?

Group 2 elements have one more proton than Group 1 elements in the same period, with both sets of outermost electrons in the same s-subshell. For example, Mg (Z = 12) vs Na (Z = 11) — both have 3s outer electrons, but Mg's nucleus pulls them more strongly. IE(Mg) = 738 kJ/mol > IE(Na) = 496 kJ/mol.

Is the periodic table memorisation-heavy for JEE?

Not if you understand trends. JEE doesn't ask you to recite atomic radii in picometres — it asks you to compare elements or explain anomalies. Know the direction of each trend, know the four key anomalies (Be > B, N > O, P > S for IE; F < Cl for EA), and you can handle any periodic table question without rote learning.

What is the difference between ionization energy and electron affinity?

IE is the energy you must supply to remove an electron from a neutral gaseous atom (always endothermic, always positive). EA is the energy released when a neutral gaseous atom gains an electron (usually exothermic). IE involves losing electrons; EA involves gaining electrons. Both generally increase across a period and decrease down a group — but their anomalies differ. The most important difference in exam context: for IE, N > O; for EA, F < Cl.