Schooling · Class 11 · Chemistry · Classification of Elements and Periodicity in Properties · Chapter on a page
CBSE Class 11 · Chemistry · NCERT Chemistry Part-I
Classification of Elements and Periodicity in Properties: chapter on a page
A cut of Shishya's notes for this chapter: the first paragraph of the introduction and the key ideas, formulas and facts, common mistakes and quick revision lists, word for word. Nothing is added or reworded; the worked examples are left out (they are on the chapter page).
About the notes: Written by Shishya's AI on 27 Sept 2026 from the chapter's title and class level, in Shishya's own words — not a copy or summary of the textbook. Read the official chapter for the book's own text, activities and exercises.
Fact-checked on 21 Oct 2026: an AI-assisted check of this note against cited sources; the errors it found were corrected that day. Some errors may remain.
Mind map
Drawn from the same notes by a fixed rule (no AI): each label is the item's own words — its bold term, its label before a colon, a short formula as written, or its first words.
The mind map as a list
- Key ideas
- Modern periodic law
- Electronic configuration determines…
- Atomic radius decreases across a…
- Atomic radius increases down a group
- Ionisation enthalpy increases across…
- Electron gain enthalpy generally…
- Electronegativity increases across a…
- Formulas and facts
- Effective nuclear charge (Z_eff)
- Ionisation enthalpy (Δ_i H)
- Electron gain enthalpy (Δ_eg H)
- Electronegativity
- Covalent radius
- Periods and blocks
- Diagonal relationship
- Worked examples
- Locating an element in the periodic…
- Comparing ionisation enthalpies
- Predicting electron gain enthalpy…
- Watch out for
- Assuming atomic radius always…
- Thinking higher electronegativity…
- Confusing ionisation enthalpy with…
- Believing Mendeleev arranged…
- Forgetting anomalies due to…
What this chapter is about
This chapter introduces how elements are organised in the periodic table based on their atomic number and electronic configuration. You will learn about the historical attempts to classify elements, starting from early groupings to Mendeleev's periodic table and finally the modern periodic table based on atomic number.
Key ideas
- Modern periodic law: Properties of elements are a periodic function of their atomic number, not atomic mass as Mendeleev originally proposed.
- Electronic configuration determines position: The period number equals the highest principal quantum number (n) of occupied orbitals; the block (s, p, d, f) depends on which subshell receives the last electron.
- Atomic radius decreases across a period because increasing nuclear charge pulls electrons closer, even though electrons enter the same shell.
- Atomic radius increases down a group because each successive element has an additional electron shell, placing valence electrons farther from the nucleus.
- Ionisation enthalpy increases across a period (more energy needed to remove an electron) due to stronger nuclear attraction; it decreases down a group as the outermost electron is farther from the nucleus.
- Electron gain enthalpy generally becomes more negative across a period (energy released on gaining an electron), reflecting greater tendency to accept electrons; halogens have the most negative values.
- Electronegativity increases across a period and decreases down a group; fluorine is the most electronegative element.
- Metallic character decreases across a period (elements lose electrons less readily) and increases down a group.
Formulas and facts to remember
- Effective nuclear charge (Z_eff): Z_eff = Z − σ, where Z is atomic number and σ is shielding constant. It measures the net positive charge felt by a valence electron.
- Ionisation enthalpy (Δ_i H): Minimum energy required to remove the most loosely bound electron from an isolated gaseous atom in its ground state. Unit: kJ mol⁻¹.
- Electron gain enthalpy (Δ_eg H): Enthalpy change when an electron is added to an isolated gaseous atom. A negative value means energy is released.
- Electronegativity: A measure of an atom's ability to attract shared electrons in a covalent bond. Pauling scale is commonly used; fluorine = 4.0.
- Covalent radius: Half the distance between nuclei of two identical atoms bonded covalently.
- Periods and blocks: Period 1 has 2 elements (1s); periods 2 and 3 have 8 elements each (s and p blocks); period 4 has 18 elements (s, p and d blocks).
- Diagonal relationship: Certain pairs (Li–Mg, Be–Al, B–Si) show similar properties because the increase in nuclear charge across a period offsets the increase in size down a group.
Common mistakes
- Assuming atomic radius always increases with atomic number → it decreases across a period because effective nuclear charge rises while the shell number stays the same.
- Thinking higher electronegativity means more negative electron gain enthalpy in every case → noble gases have positive or near-zero electron gain enthalpy despite being after halogens.
- Confusing ionisation enthalpy with electron gain enthalpy → ionisation removes an electron (energy absorbed); electron gain adds an electron (usually energy released for non-metals).
- Believing Mendeleev arranged elements by atomic number → he used atomic mass; Moseley later established atomic number as the basis.
- Forgetting anomalies due to half-filled and fully filled subshell stability → Mg has higher ionisation enthalpy than Al; N has higher ionisation enthalpy than O.
Quick revision
- Period = highest principal quantum number; group = number of valence electrons for s-block elements and 10 + number of valence electrons for p-block elements.
- Across a period: size decreases, ionisation enthalpy increases, electronegativity increases, metallic character decreases.
- Down a group: size increases, ionisation enthalpy decreases, electronegativity decreases, metallic character increases.
- s-block: groups 1–2; p-block: groups 13–18; d-block: groups 3–12; f-block: lanthanoids and actinoids.
- Electron gain enthalpy is most negative for halogens; Cl > F in magnitude due to size effect.