Antoine Lavoisier (18th century) first attempted to classify known elements as metals and nonmetals.
Döbereiner (1829) grouped elements into triads – in each triad, the atomic weight of the middle element was roughly the average of the other two. Example: Li (7), Na (23), K (39).
John Newlands (1864) arranged 62 elements in increasing atomic mass and observed that every eighth element had similar properties – Law of Octaves.
Lothar Meyer (1869) plotted atomic volume vs. atomic weight; the curves showed periodicity.
Dmitri Mendeleev (1869) arranged 63 elements by atomic mass, left gaps for undiscovered elements, and accurately predicted their properties. He is considered the father of the Periodic Table.
Henry Moseley (1913) determined atomic numbers using X‑ray emission and modified the periodic law: “Properties of elements are periodic functions of their atomic numbers.”
Metals: Tend to lose electrons to form positive ions. Examples: Fe, Cu, Au, Ag.
Non‑metals: Tend to gain electrons to form negative ions. Examples: Cl, S, P.
Metalloids (semi‑metals): Exhibit properties of both metals and non‑metals. They lie along the stair‑step line: B, Si, Ge, As, Sb, Te, Po. Elements to the left of the line are metals; those to the right are non‑metals. Hydrogen is an exception (non‑metal).
s‑block (Groups 1‑2): Valence electrons in s‑orbital. Includes alkali and alkaline earth metals.
p‑block (Groups 13‑18): Valence electrons in p‑orbital. Includes halogens, noble gases, and metalloids.
d‑block (Groups 3‑12): Transition elements with valence electrons in d‑orbital.
f‑block (Lanthanides & Actinides): Valence electrons in f‑orbital; placed at the bottom.
Alkali metals (Group 1, except H): 1 valence e⁻, most reactive, form alkalis with water.
Alkaline earth metals (Group 2): 2 valence e⁻, less reactive, found in earth’s crust.
Transition elements (d‑block): Variable oxidation states, coloured compounds, catalytic properties.
Chalcogens (Group 16): O, S, Se, Te, Po, Lv – called “chalcogens” because most copper ores are oxides or sulfides (Greek chalkos = copper).
Halogens (Group 17): F, Cl, Br, I, At, Ts – “salt‑formers”; high electron affinity, 7 valence e⁻.
Noble gases (Group 18): He, Ne, Ar, Kr, Xe, Rn, Og – complete octet, very low reactivity; Xe forms some compounds (e.g., XePtF₆).
Period number = principal quantum number (n) = number of electron shells.
Group number = number of valence electrons (for main group elements).
Example: Element in period 3, group 2 → Mg (1s²2s²2p⁶3s²).
Element in period 2, group 14 → Carbon (1s²2s²2p²).
Element in period 4, group 17 → Br (1s²2s²2p⁶3s²3p⁶3d¹⁰4s²4p⁵).
Modern Periodic Law: “The physical and chemical properties of elements are periodic functions of their atomic numbers.”
Definition: Half the distance between two identical bonded atoms. Units: picometers (pm) or Angstroms (Å).
Factors: Number of shells, effective nuclear charge, shielding effect.
Trends: Decreases across a period (↑ nuclear charge pulls electrons closer). Increases down a group (↑ shells, more shielding).
Cation: Smaller than parent atom because loss of electrons reduces repulsion and nuclear pull increases.
Anion: Larger than parent atom because gain of electrons increases repulsion and electron cloud expands.
Trends: Ionic radii decrease across a period (for cations and anions) due to increasing nuclear charge. Increase down a group because more shells are added.
First IE: Energy required to remove 1 mole of electrons from 1 mole of gaseous atoms: Na(g) → Na⁺(g) + e⁻, ΔHᵢ₁ = 494 kJ/mol.
Second IE: Ca⁺(g) → Ca²⁺(g) + e⁻, ΔHᵢ₂ = 1150 kJ/mol.
Factors: Nuclear charge (↑ IE), atomic size (↓ size ↑ IE), electronic arrangement (half‑filled/fully filled orbitals are stable → ↑ IE), shielding (↑ shielding ↓ IE), spin‑pair repulsion (paired electrons easier to remove).
Trends: IE increases across a period (↑ nuclear charge, same shell). IE decreases down a group (↑ size, ↑ shielding).
First EA: Enthalpy change when 1 mole of electrons is added to 1 mole of gaseous atoms: Cl(g) + e⁻ → Cl⁻(g), ΔHₑₐ₁ = –348.8 kJ/mol.
Second EA: O⁻(g) + e⁻ → O²⁻(g), ΔHₑₐ₂ = +798 kJ/mol (endothermic due to repulsion).
Factors: Atomic size (↓ size ↑ EA), nuclear charge (↑ charge ↑ EA), electronic configuration (half‑filled subshells have low EA).
Trends: EA becomes more negative across a period (↑ nuclear charge, ↓ size). EA becomes less negative down a group (↑ size, more shielding).
Definition: Measure of an atom’s attraction for bonding electrons in a molecule. Pauling scale: F = 4.0 (highest), Cs = 0.8 (lowest).
Factors: Atomic size (↓ size ↑ EN), effective nuclear charge (↑ charge ↑ EN).
Trends: EN increases across a period (↑ nuclear charge, ↓ size). EN decreases down a group (↑ size, ↑ shielding).
Definition: Tendency to lose electrons and form positive ions.
Trends: Metallic character decreases across a period (↑ nuclear charge makes it harder to lose electrons). Metallic character increases down a group (↑ size, ↓ IE, electrons held more loosely).
Sodium: Reacts vigorously with cold water: 2Na(s) + 2H₂O(l) → 2NaOH(aq) + H₂(g).
Magnesium: Reacts slowly with cold water: Mg(s) + 2H₂O(l) → Mg(OH)₂(aq) + H₂(g). With steam, it reacts vigorously: Mg(s) + 2H₂O(g) → MgO(s) + 2H₂(g).
Sodium: Burns with golden yellow flame to form sodium peroxide: 2Na(s) + O₂(g) → Na₂O₂(s). Under limited oxygen: 4Na(s) + O₂(g) → 2Na₂O(s).
Magnesium: Burns with intense white flame to form magnesium oxide: 2Mg(s) + O₂(g) → 2MgO(s).
Sodium: 2Na(s) + Cl₂(g) → 2NaCl(s) (white solid, golden yellow flame).
Magnesium: Mg(s) + Cl₂(g) → MgCl₂(s) (white solid).
Oxides of period 3: Ionic character decreases from left to right. Na₂O, MgO are ionic (giant lattices). Al₂O₃ is amphoteric. SiO₂, P₂O₅, SO₂, SO₃ are covalent (molecular).
Basic oxides: React with water to give alkalis. Example: Na₂O(s) + H₂O(l) → 2NaOH(aq).
Acidic oxides: React with water to give acids. Example: SO₃(g) + H₂O(l) → H₂SO₄(aq); react with bases: SO₂(g) + 2NaOH(aq) → Na₂SO₃(aq) + H₂O(l).
Amphoteric oxides: React with both acids and bases. Example: Al₂O₃(s) + 6HCl(aq) → 2AlCl₃(aq) + 3H₂O(l) ; Al₂O₃(s) + 2NaOH(aq) → Na₂Al₂O₄(aq) + H₂O(l).
Neutral chlorides: NaCl, MgCl₂ – dissolve in water to give neutral solutions (pH ≈ 7).
Acidic chlorides: AlCl₃, SiCl₄, PCl₅ – hydrolyse in water to produce HCl, making acidic solutions (pH < 7).
Oxidation number is the formal charge on an atom in a molecule or ion. For 3rd period elements, the maximum oxidation number equals the group number.
In oxides: Na₂O (+1), MgO (+2), Al₂O₃ (+3), SiO₂ (+4), P₂O₅ (+5), SO₃ (+6).
In chlorides: NaCl (+1), MgCl₂ (+2), AlCl₃ (+3), SiCl₄ (+4), PCl₅ (+5), SCl₂ (–2 for S).
Phosphorus and sulfur show variable oxidation states because they can expand their octet by using empty 3d orbitals.
📘 Study guidelines: Read each section carefully, use the mnemonic tips, and practice with the 30 MCQs below. Adjust font size for comfortable reading. Toggle night mode for late‑night study. Focus on trends – they are the key to understanding periodicity.