Periodic table and periodicity

Why classify

Too many elements to learn one by one. Put similar outer configurations in the same column so one member teaches you the family.

Periodic table: arrangement where similar properties fall under the same group.

Short history path

Dobereiner’s triads: groups of three; middle atomic weight ≈ average of the other two. Only a few sets fit.

Newlands’ octaves: every eighth element echoed the first (noble gases were unknown then, so modern counting looks like 1st and 9th).

Mendeleev: properties are a periodic function of atomic weight. Groups I–VIII with A/B subgroups; left gaps that later filled with real elements.

Mendeleev’s pain points: hydrogen’s place, isotopes sharing one box, Te before I by mass, and no clear “why” behind the periodicity.

Modern law (Moseley): properties are a periodic function of atomic number. Long form: 7 periods, 18 IUPAC groups.

Reading the long form

Period number ≈ highest principal quantum number n for that row. Period lengths: 2, 8, 8, 18, 18, 32… (7th still filling out in textbooks).

Lanthanoids and actinoids sit in panels at the bottom so the table stays usable.

Elements with Z > 100 get systematic IUPAC names built from digit roots + “ium”. You mainly need the idea, not every name memorised.

s, p, d, f blocks

s-block: Groups 1–2. Outer config ns¹–². Soft, reactive metals; low IE; often colour flames (Be, Mg exceptions). Strongly electropositive.

p-block: Groups 13–18. Outer ns² np¹–⁶. With s-block → representative / main-group elements. End of each period: noble gas ns² np⁶.

d-block: Groups 3–12. (n−1)d¹–¹⁰ ns⁰–². Transition metals: coloured ions, variable oxidation states, often catalysts, high melting points.

f-block: lanthanoids Ce–Lu, actinoids Th–Lr. Inner transition. (n−2)f filling. Actinoids radioactive; post-U = transuranium.

Atomic and ionic size

Atomic radius: nucleus centre → outermost electron cloud (definitions differ: covalent, van der Waals, metallic — know which your question means).

Across a period: nuclear charge up, same shell filling → size generally shrinks.

Down a group: new shells → size grows.

Cation < parent atom (less e^-, same Z pulls harder). Anion > parent atom (extra e^-, more repulsion).

Isoelectronic ions (same e^- count): higher nuclear charge → smaller ion. Classic set: N^3- > O^2- > F^- > Na⁺ > Mg²⁺ > Al³⁺.

IE, electron gain enthalpy, electronegativity

Ionisation enthalpy ΔᵢH: energy to remove the outermost e^- from gaseous atom. Across period: generally up. Down group: generally down.

Exceptions students must flag: Be, Mg, N, P, noble gases look “high” because of stable configs. He is extreme. Second IE of Na is huge (you’re ripping from Ne-like Na⁺).

Electron gain enthalpy Δ_egH: enthalpy when gaseous atom gains an e^-. More negative = easier anion formation. Cl is the classic “most negative” in many tables; F is less negative than Cl because of small-size repulsion.

Successive electron gains are endothermic (you’re charging a negative ion further).

Electronegativity: pull on a shared pair in a bond. F highest, Cs among the lowest. Not the same thing as IE — you’re in a molecule, not an isolated atom.

Chemical trends to recite

Across a period (L → R): metallic character down; non-metallic up; IE and electronegativity up; oxide basicity down, acidity up.

Down a group: metallic character up; non-metallic down; size up; IE and electronegativity down; oxide basicity up for metals.

Alkali metals: reactivity increases down. Halogens: reactivity decreases down.

Diagonal relationship: Li–Mg, Be–Al, B–Si type pairs — similar charge/size ratio → similar chemistry.