Periodic classification arranges elements in a table so that elements with similar properties recur periodically when ordered by increasing atomic number.
Groups & Periods
Periodic Trends
Electronic Config
Mendeleev vs Modern
NCERT Class 10 — Periodic Classification
- Arranged by increasing atomic weight.
- Grouped elements with similar properties in columns.
- Left gaps for undiscovered elements and predicted their properties.
- Could not explain some anomalies (isotopes).
Modern periodic law: properties of elements are periodic functions of their atomic numbers (Z). This resolves anomalies caused by isotopes and places elements correctly by Z, not atomic weight.
A period is a horizontal row; elements in a period have the same number of electron shells. A group is a vertical column; elements in a group have similar valence electron configurations and chemical properties.
The modern periodic table has 18 groups and 7 periods (excluding the lanthanoids and actinoids which are placed separately as f-block).
Elements are placed into blocks (s, p, d, f) and groups according to their valence electron configuration. The valence electrons largely determine chemical reactivity and group properties.
- s-block: groups 1–2 (e.g., Na, Mg)
- p-block: groups 13–18 (e.g., C, O, Cl)
- d-block: transition metals (e.g., Fe, Cu)
- f-block: lanthanoids/actinoids (e.g., La, U)
Valency is the combining capacity of an element, generally equal to the number of electrons lost, gained or shared to achieve a noble gas configuration. From electronic configuration, valency is determined by the number of valence electrons (or 8 − valence electrons for main-group elements).
Na (11): 1s² 2s² 2p⁶ 3s¹. Valence electron is in 3s¹, placing Na in group 1 (s-block).
Atomic radius is half the distance between nuclei of two bonded atoms. Across a period: radius decreases (effective nuclear charge increases). Down a group: radius increases (more shells).
Ionization energy is the energy required to remove an electron from a gaseous atom. Across a period: increases (atoms hold electrons tightly). Down a group: decreases (outer electrons are farther and shielded).
Electron affinity is the energy change when an atom gains an electron (in gas phase). Electronegativity is a relative measure of an atom's ability to attract bonding electrons. Electronegativity is a periodic property used in bonds; electron affinity is a measurable energy change.
Electronegativity generally increases across a period (stronger nuclear attraction) and decreases down a group (increased distance and shielding).
Across a period: metallic character decreases, non-metallic increases. Down a group: metallic character increases (atoms more willing to lose electrons).
Hydrogen is placed in group 1 by many tables due to its 1s¹ configuration, but it shares properties with group 1 (forms H⁺) and group 17 (forms H⁻ in hydrides). Its behavior is unique, so its placement is often shown separately.
Noble gases (group 18) have full valence shells, are unreactive, monoatomic gases. Examples: He (balloons, cryogenics), Ne (neon signs), Ar (inert atmosphere for welding, bulbs).
Alkali metals are soft, highly reactive (react with water to form hydroxides and H₂), have low melting points, and form +1 ions. Uses: Na (NaCl, NaOH), K (fertilizers). Flame tests: Na — yellow; K — lilac.
Halogens are non-metals with seven valence electrons, highly reactive, form −1 ions, and react with metals to form ionic salts (e.g., NaCl). Chlorine used in water treatment and bleaching but toxic in gaseous form.
Transition elements are d-block elements showing variable oxidation states, form coloured compounds, and often act as catalysts (e.g., Fe, Cu, V, Mn).
Across a period, metallic reactivity decreases (atoms less willing to lose electrons). Down a group, metallic reactivity increases (outer electrons are easier to lose).
Cations (positive) are smaller than parent atoms; anions (negative) are larger. Across a period ionic radii vary with charge and type; general trend: cations decrease size across period; anions also show size changes but depend on exact element.
Diagonal relationship: certain pairs of diagonally adjacent elements in the s-and p-blocks show similar properties (e.g., Li and Mg). Similarity arises because of comparable ionic sizes and polarising power.
They have complete valence shells (stable electronic configuration), so they have little tendency to gain, lose or share electrons, making them largely unreactive.
Valence electrons = 6 → group 16 (oxygen family). The element is oxygen (atomic number 8).
Group 2 elements (alkaline earth metals) have 2 valence electrons (e.g., Be: 2, Mg: 2). They typically form +2 ions.
The period number equals the number of electron shells. Configuration 2,8,1 has electrons in three shells → period 3 (element is Na).
Periodicity is the recurring variation of properties of elements when arranged by atomic number. It is important because it helps predict behaviour and chemical properties of elements.
Because they have the same number of valence electrons and similar valence shell configurations, leading to similar chemical behavior (e.g., group 1 elements all form +1 ions).
Example: Alkali metals (Li → Na → K → Rb → Cs) become more metallic and reactive down the group because outer electrons are farther from nucleus and more easily lost.
Across period → shrink; Down group → enlarge.
Some elements did not fit when ordered strictly by atomic weight (e.g., iodine and tellurium) — ordering by atomic number resolved these anomalies.
They belong to the f-block and are placed separately to keep the table compact. Lanthanoids (4f) and actinoids (5f) have filling f-orbitals and similar properties within series.
Fluorine has the highest electronegativity due to small size and high effective nuclear charge, attracting bonding electrons strongly.
K is larger because it has an extra electron shell (Na: 2,8,1; K: 2,8,8,1), increasing atomic radius down the group.
Mg has higher ionization energy than Na because Mg has a greater effective nuclear charge for the valence electrons (despite both being in same period), holding electrons more tightly.
Chlorine is smaller and more electronegative, so it attracts electrons more strongly than bromine — making it a stronger oxidizing agent and more reactive.
Trends are not uniform: metallic elements often have high boiling points; molecular non-metals have low boiling points. Down a group metallic bonding may weaken or strengthen depending on structure — many exceptions exist.
Metals are on the left and center, non-metals on the right; metalloids lie along the zig-zag (staircase) line between metals and non-metals (e.g., B, Si, Ge).
Elements in a period have the same number of shells but different valence electrons, hence different properties. Example: Na (metal) vs Cl (non-metal) in period 3.
Because properties vary predictably with position (atomic number), we can predict reactivity, valency, ionic charge, metallic vs non-metallic behavior and fuels for element behavior and reactions.
Cl (17): 1s² 2s² 2p⁶ 3s² 3p⁵ → valence = 7 → group 17, period 3.
Moseley measured atomic numbers using X-ray spectra and showed that atomic number (not weight) is the correct basis for the table — leading to the modern periodic law.
Period 3, group 16 → like sulfur: valence electrons = 6, forms −2 anions or covalent compounds, non-metallic, higher boiling point than oxygen.
Atomic number (Z) is the number of protons in the nucleus of an atom and equals the number of electrons in a neutral atom.
Isotopes are atoms of the same element with different mass numbers (different neutrons). They have the same chemical properties because they have same electronic configuration; periodic classification is by atomic number, so isotopes do not change position.
If valence is 2 (forms M²⁺), oxide will be M O (balance 2+ and 2−) → formula MO (e.g., MgO).
Electropositivity is tendency to lose electrons (opposite of electronegativity). It increases down a group. Cesium (and francium theoretically) are highly electropositive among common elements.
2,8,3 → Aluminium (Al), period 3, group 13, valency commonly 3 (forms Al³⁺).
Atomic number determines electronic configuration; elements with similar valence electrons show similar chemical reactivity — allowing predictions about compound formation, valency and reactions.
- Modern periodic law: properties depend on atomic number.
- Groups: similar valence electrons; Periods: same shells.
- Trends: atomic radius, ionization energy, electronegativity, metallic character.
- Special blocks: s, p, d (transition), f (inner transition).
- Mendeleev predicted properties; Moseley established atomic number basis.