🧪 Periodic Table – Full Breakdown

ICSE Class 10 Chemistry June 2026

        This is the chapter that scares people because of so many trends. But honestly? It's the most logical chapter in Chemistry. Everything follows a pattern. Once you get the "why" behind each trend, you don't need to memorise anything. I'll explain the reasons, the exceptions, and even some out-of-syllabus stuff that helps you actually understand what's happening. Written by someone who scored 93.4% in ICSE 2026 — so trust me, this works.
    

1) Why Do We Even Need a Periodic Table?

Back in the day, scientists kept discovering elements randomly. By 1860s, they had around 60+ elements with no system. Imagine trying to study 118 elements without any order — nightmare. So people started looking for patterns.

2) Dobereiner's Triads (1817)

Johann Dobereiner noticed that when you take three elements with similar properties, the atomic mass of the middle element is roughly the average of the other two. Example:

Li (7), Na (23), K (39) → (7+39)/2 = 23 ✅
Ca (40), Sr (88), Ba (137) → (40+137)/2 = 88.5 ✅
Cl (35.5), Br (80), I (127) → (35.5+127)/2 = 81.25 ✅ (close enough)

Why it failed: He could only make a few triads. Most elements didn't fit into groups of three. Also, as more elements were discovered, the pattern broke.

3) Newlands' Law of Octaves (1864)

Newlands arranged elements in order of increasing atomic mass and noticed that every 8th element had similar properties — like musical notes (sa-re-ga-ma…). He called it the Law of Octaves.

Why it failed:

Only worked till Calcium (up to 56 elements). After that, the pattern broke completely.
Noble gases weren't discovered yet — if they were, the whole "every 8th" thing wouldn't fit.
He tried to force elements into the same slot even when they didn't match (like Co and Ni).

4) Mendeleev's Periodic Table (1869) — The Real OG

Mendeleev is the GOAT of periodic table history. His table was based on atomic mass, but here's what made him brilliant:

Left gaps for undiscovered elements and predicted their properties with scary accuracy. He called one "Eka-aluminium" (below Al) — later discovered as Gallium (Ga), and its properties matched his prediction almost perfectly.
Swapped elements when atomic mass order didn't make sense chemically. For example, he placed Tellurium (128) before Iodine (127) even though Te is heavier — because Iodine's properties matched the Halogen group better. That was chemically correct even though the mass order was wrong. Mendeleev trusted chemistry over mass.

Limitations:

No place for isotopes (same element, different mass — where do they go?)
Couldn't explain why atomic mass order had to be broken in some cases (the answer came later with Moseley)
Position of Hydrogen was unclear (still debated today tbh)

📖 Out of syllabus but cool: Mendeleev predicted Eka-boron (Scandium), Eka-aluminium (Gallium), and Eka-silicon (Germanium). His predictions included density, melting point, and formula of oxides — and when these elements were discovered, his numbers were off by less than 5%. Absolutely insane for a guy working without modern instruments.

5) Modern Periodic Table — What We Study Today

Henry Moseley (1913) fixed everything. He used X-ray experiments to show that atomic number (Z) is the fundamental property, not atomic mass. This is why the modern table is called the Periodic Law: "Properties of elements are a periodic function of their atomic numbers."

Structure:

7 Periods (horizontal rows) — period number = number of shells
18 Groups (vertical columns) — same group = same valence electrons
Blocks: s-block (Gr 1-2), p-block (Gr 13-18), d-block (Gr 3-12, transition metals), f-block (lanthanides & actinides, placed below)

6) Periodic Properties — The Heart of the Chapter

This section has the most marks. Let's go property by property with reasons, exceptions, and my personal tricks.

💡 My personal memory trick (scored 93.4% with this):
Atomic Radius and Metallic Character — these TWO decrease across a period and increase down a group.
Everything else (Ionization Energy, Electron Affinity, Electronegativity, Non-metallic Character) — do the opposite.

So just remember: "AR and MC are the rebels. Rest follow the opposite pattern."

a) Atomic Radius (Size)

Definition: Distance from the centre of the nucleus to the outermost shell.

Across a period → Decreases
Reason: Nuclear charge (number of protons) increases. Same shell, so same distance. More protons pull the electrons closer to the nucleus. Net effect: size shrinks.

Down a group → Increases
Reason: New shell is added each time you go down. The outermost electron is farther from the nucleus. Also, inner shells shield the outer electron from the nuclear pull (shielding effect).

Exception to know: Noble gases have slightly larger atomic radii than the halogen before them in the same period. Reason? Atomic radius for noble gases is measured as van der Waals radius (distance between nuclei of non-bonded atoms) while for others we use covalent or metallic radius. Different measurements, so direct comparison is unfair. ICSE usually doesn't ask this directly but it comes in MCQs sometimes.

📖 Out of syllabus: Francium (Fr) has the largest atomic radius among known elements. Helium (He) has the smallest atomic radius as a measured element (but technically Hydrogen has the smallest nucleus). Also, atomic radius isn't a fixed value — it changes depending on whether the atom is bonded or not. That's why we have different types: covalent radius, metallic radius, van der Waals radius.

b) Ionization Energy (IE)

Definition: Energy required to remove the most loosely bound electron from an isolated gaseous atom.

Across a period → Increases
Reason: Nuclear charge increases, atomic size decreases → electrons are held more tightly. Harder to remove an electron, so more energy needed.

Down a group → Decreases
Reason: Size increases, more shielding → outermost electron is less tightly held. Easier to remove.

Exceptions (very important for boards):

IE of Mg > Al even though Al is to the right of Mg. Why? Mg has 3s² configuration (fully filled s-orbital). Removing an electron from a filled orbital is harder. Al has 3s²3p¹ — that 3p¹ electron is in a higher energy orbital and is easier to remove.
IE of N > O even though O is to the right of N. Why? N has 2s²2p³ — half-filled p-orbital = extra stability. O has 2s²2p⁴ — removing one electron gives half-filled 2p³ which is actually MORE stable, so O's IE is lower than expected.
IE of Be > B — Be has 2s² (filled s), B has 2s²2p¹ (p electron easier to remove).

📖 Out of syllabus: The first IE of Helium is the highest of any element (2372 kJ/mol). Francium has the lowest (about 380 kJ/mol). Also, successive ionization energies always increase (IE₁ < IE₂ < IE₃…) because you're removing an electron from an increasingly positive ion. A huge jump in IE values tells you that you've gone past the valence shell into the core electrons — this is how scientists determine valency experimentally.

c) Electron Affinity (EA)

Definition: Energy released when an electron is added to a neutral gaseous atom.

Across a period → Increases (generally)
Reason: Nuclear charge increases, size decreases → incoming electron feels stronger attraction → more energy released.

Down a group → Decreases (generally)
Reason: Size increases, shielding increases → incoming electron is less attracted → less energy released.

Exceptions (board loves these):

EA of Cl > F. This is the most famous exception. Fluorine is the most electronegative element, but Chlorine has the highest electron affinity. Why? F is tiny — when you add an electron to its 2p orbital, the electron experiences strong repulsion from the already existing electrons (electron-electron repulsion). Cl has a larger 3p orbital with more room, so less repulsion.
EA of N is nearly zero or positive. N has half-filled 2p³. Adding an electron means disrupting a stable half-filled configuration. Energy is actually required (endothermic) rather than released.
Noble gases have negative EA (they don't want an extra electron, energy is absorbed, not released).

📖 Out of syllabus: Some textbooks call Electron Affinity with a negative sign convention (energy released = negative value). ICSE usually just talks about "more" or "less" energy released, so don't stress about signs. But in higher classes (11th, JEE), EA₁ is usually negative (exothermic) and EA₂ is positive (endothermic, because you're adding electron to an already negative ion).

d) Electronegativity (EN)

Definition: The tendency of an atom to attract a shared pair of electrons in a covalent bond.

Across a period → Increases
Reason: Smaller size + higher nuclear charge = stronger pull on bonding electrons.

Down a group → Decreases
Reason: Larger size + more shielding = weaker pull.

Exceptions:

There aren't really "exceptions" in the same way as IE and EA, but Fluorine is the most electronegative element (4.0 on Pauling scale) and Francium is the least (about 0.7).
Oxygen and Nitrogen are close (3.44 vs 3.04) but Oxygen wins by a tiny margin.
Noble gases don't have electronegativity values (they don't form bonds under normal conditions). Xenon is an exception — it can form compounds (XeF₂, XeF₄) and has an EN around 2.6.

📖 Out of syllabus: The Pauling scale (0-4) is what's used most. There's also the Mulliken scale (average of IE and EA) and the Allred-Rochow scale. F is the most EN on all scales. Second place varies: O on Pauling, Cl on some others. For ICSE, just know F > O > Cl > N.

e) Metallic Character

Definition: Tendency of an element to lose electrons and form positive ions.

Across a period → Decreases
Reason: IE increases → harder to lose electrons → less metallic. Elements go from metals → metalloids → non-metals.

Down a group → Increases
Reason: IE decreases → easier to lose electrons → more metallic.

Exceptions: None major in ICSE. But interestingly, Hydrogen sits at the top of Group 1 but is NOT a metal (it's a non-metal). Its position is debated — it also has properties of Group 17 (halogens).

f) Non-Metallic Character

Opposite of metallic character. Increases across a period, decreases down a group.

Reason: Higher EN and EA → tendency to gain electrons → non-metallic behaviour.

7) Quick Reference Table (Print This)

Property	Across Period →	Down Group ↓
Atomic Radius	⬇ Decreases	⬆ Increases
Ionization Energy	⬆ Increases	⬇ Decreases
Electron Affinity	⬆ Increases*	⬇ Decreases*
Electronegativity	⬆ Increases	⬇ Decreases
Metallic Character	⬇ Decreases	⬆ Increases
Non-Metallic Character	⬆ Increases	⬇ Decreases
Valence Electrons	⬆ Increases by 1	Same (constant)
Number of Shells	Same (constant)	⬆ Increases by 1

* with exceptions (Cl > F for EA, etc.)

⚠️ My go-to trick for board exam: If a question asks "Arrange these elements in increasing order of X property" — first write down their periods and groups. Same period? Compare from left to right. Same group? Compare from top to bottom. Different period AND group? The one in the lower period is generally bigger (for atomic radius). ICSE loves these comparison questions. I've seen at least one every year.

8) Group-wise Personality (What Each Group is Like)

Group	Name	Personality	Valence
1	Alkali Metals	Super reactive, explode in water, +1 charge	1
2	Alkaline Earth Metals	Reactive but less than Gr 1, +2, harder	2
13	Boron Family	+3, mixed bag (B is metalloid, rest metals)	3
14	Carbon Family	The middle ground — C is non-metal, Si/Ge metalloids, Sn/Pb metals	4
15	Nitrogen Family	N and P non-metals, As/Sb metalloids, Bi metal	5
16	Oxygen Family (Chalcogens)	O, S non-metals, Se/Te metalloids, Po metal	6
17	Halogens	Most reactive non-metals, -1, diatomic, colour intensity increases down	7
18	Noble Gases	"Royal" — don't react, full outer shell, stable	8 (except He = 2)
3–12	Transition Metals	Hard, high melting points, form coloured compounds, variable oxidation states	Variable

9) Top Questions ICSE Loves to Ask

Q1: Arrange in order of increasing atomic radius: Na, Mg, Al, Si, P, S, Cl

Answer: Cl < S < P < Si < Al < Mg < Na
Reason: Same period (3), nuclear charge increases from left to right, so atomic radius decreases.

Q2: Why is the ionization energy of Mg greater than Al?

Answer: Mg has electronic configuration 2,8,2 — its outermost electrons are in 3s² (filled s-orbital). Al has 2,8,3 — outermost electron is in 3p¹. The 3p electron is in a higher energy orbital and is shielded by s-electrons, making it easier to remove. So IE of Mg > Al despite Al being to the right.

Q3: Why does Cl have higher electron affinity than F?

Answer: F is very small (2nd period). Adding an electron to its 2p orbital causes significant electron-electron repulsion because the 2p orbitals are compact. Cl (3rd period) has larger 3p orbitals with more space, so less repulsion when adding an electron. Thus, Cl releases more energy (higher EA) despite F being more electronegative.

Q4: Which is more metallic — Na or Mg? Why?

Answer: Na. Both in period 3. Na has lower IE (easier to lose 1 electron) than Mg (needs more energy to lose 2 electrons). Na readily forms Na⁺, while Mg is less willing. Metallic character decreases across a period.

Q5: Why is Nitrogen less reactive than Phosphorus?

Answer (slightly out of syllabus but useful): N₂ has a triple bond (extremely strong, 941 kJ/mol). P₄ has single bonds (much weaker, ~200 kJ/mol). P is more reactive because its bonds break easily. But in terms of periodicity, N has higher EN and smaller size.

⚠️ From my exam experience: The most common mistake students make is mixing up the trends. Remember my trick — "AR and MC decrease across a period, increase down a group. Everything else is opposite." Write this at the top of your rough sheet before starting the paper. I did this and it saved me in the exam hall.

🧠 One more tip from a 93.4% scorer: For reasoning questions, always mention TWO things: (1) the change in nuclear charge, and (2) the change in atomic size / shielding. Most students only mention one. If you mention both, you get full marks even if your language is messy. Example: "IE increases across a period because nuclear charge increases AND atomic size decreases, so electrons are held more tightly."