S- and P- Block Elements

The s-block (Groups I and II) and p-block (Groups III–VIII / 13–18) together make up the "main-group" elements. The PMDC MDCAT 2026 syllabus expects you to know periodic trends across and down these blocks, the characteristic reactions of Groups I, II and IV, and how to classify any element into s, p, d or f block by its electronic configuration.

PMC Table of Specifications. This chapter covers five PMDC subtopics — Properties & Trends, Reactions of Group I/II/IV elements, and S/P/D/F block overview. Skim the headings below to confirm full coverage.

S/P/D/F Block Elements (overview)

The block to which an element belongs is decided by the orbital being filled by the last electron in its ground-state configuration.

s-block — last electron enters an s orbital. Groups I (alkali metals) and II (alkaline-earth metals); also H and He. General configuration ns¹ or ns². Soft, low density, low IE, highly reactive metals (except H/He).
p-block — last electron enters a p orbital. Groups III to VIII / 13–18. Configuration ns²np^1–6. Includes metals (Al, Sn, Pb), metalloids (B, Si, Ge, As), and all the non-metals plus noble gases.
d-block (transition elements) — last electron enters a (n−1)d orbital. Groups 3–12. Configuration (n−1)d^1–10ns^0–2. Hard, high melting points, variable oxidation states, often coloured compounds.
f-block (inner transition) — last electron enters an (n−2)f orbital. Lanthanides (4f, Ce–Lu) and actinides (5f, Th–Lr). Many actinides are radioactive.

Properties and Trends

Periodic trends arise from the interplay between effective nuclear charge (Z_eff) and the principal quantum number (n) of the outermost electrons.

Atomic radius

Across a period (L → R): decreases. New electrons enter the same shell while nuclear charge rises — Z_eff ↑ pulls electrons closer.
Down a group (T → B): increases. New shells are added; the inner shells shield the outer electrons.

Ionisation energy (IE)

The energy needed to remove the most loosely bound electron from a gaseous atom. Opposite trend to atomic radius:
Across a period: increases (smaller atoms hold electrons more tightly).
Down a group: decreases (outer electrons are farther and better shielded).
Anomalies: Be > B and N > O because of stable filled-2s and half-filled-2p configurations.

Electron affinity

Energy released when a gaseous atom gains an electron. Generally becomes more negative across a period, less negative down a group. Cl has a more negative EA than F (F's small size and electron–electron repulsion in the small 2p shell).

Electronegativity (EN)

Tendency of an atom in a bond to attract the shared electron pair. Trend mirrors IE: increases across, decreases down. F is the most electronegative element on the Pauling scale (4.0).

Metallic / non-metallic character

Metallic character decreases across a period (right side becomes non-metals) and increases down a group (heavier elements lose electrons more easily). Most reactive metal → Cs (or Fr); most reactive non-metal → F.

Reactions of Group I Elements

Group I = alkali metals: Li, Na, K, Rb, Cs, Fr. Configuration ns¹. They lose the single s electron easily, giving M⁺. Reactivity rises down the group: Cs > Rb > K > Na > Li.

Characteristic reactions

With water: 2M + 2H₂O → 2MOH + H₂↑. Vigorous; K and below burn with a flame; Cs is explosive even with cold water.
With oxygen: Li gives normal oxide Li₂O; Na gives peroxide Na₂O₂; K, Rb, Cs give superoxides MO₂ (KO₂).
With halogens: 2M + X₂ → 2MX (ionic salts; NaCl, KBr).
With dilute acids: 2M + 2HCl → 2MCl + H₂ (extremely vigorous, dangerous).
Flame colours: Li → crimson, Na → golden yellow, K → lilac, Rb → red-violet, Cs → blue. Used as an analytical test.

Hydroxides are strong bases, increasing in basicity down the group: LiOH < NaOH < KOH < RbOH < CsOH.

Reactions of Group II Elements

Group II = alkaline-earth metals: Be, Mg, Ca, Sr, Ba, Ra. Configuration ns². They form M²⁺ ions. Reactivity rises down the group; Be is markedly less reactive than the rest, and BeCl₂ is essentially covalent (small size, high charge density — Fajans' rules) while MgCl₂…BaCl₂ are ionic.

Characteristic reactions

With water: Be does not react. Mg reacts only with steam: Mg + H₂O(g) → MgO + H₂. Ca, Sr, Ba react steadily with cold water giving M(OH)₂ + H₂.
With oxygen: 2M + O₂ → 2MO. Mg burns with a brilliant white flame.
With halogens: M + X₂ → MX₂.
With nitrogen (heated): 3M + N₂ → M₃N₂ (e.g. Mg₃N₂).
With acids: M + 2HCl → MCl₂ + H₂.
Flame colours: Ca brick-red, Sr crimson, Ba apple-green; Be and Mg show no flame colour (high IE).

Solubility of sulphates decreases down the group (BeSO₄ soluble → BaSO₄ insoluble); solubility of hydroxides increases down the group (Be(OH)₂ insoluble → Ba(OH)₂ soluble).

Reactions of Group IV Elements

Group IV (Group 14) = C, Si, Ge, Sn, Pb. Configuration ns²np². Properties span non-metal (C) → metalloid (Si, Ge) → metal (Sn, Pb).

Inert pair effect

Down the group, the ns² pair becomes increasingly reluctant to take part in bonding because of poor shielding of the inner d/f electrons. Therefore the +2 oxidation state becomes more stable than +4 for heavier elements:

C and Si almost exclusively show +4 (CO₂, SiO₂, CCl₄, SiCl₄).
Ge and Sn show both +2 and +4, with +4 more stable for Sn (Sn²⁺ is a reducing agent).
Pb shows both, but +2 is more stable (PbCl₂ stable, PbCl₄ decomposes; PbO₂ is a strong oxidiser).

Characteristic reactions

With oxygen: M + O₂ → MO₂ for C, Si (CO₂, SiO₂); Sn gives SnO₂; Pb gives PbO (preferentially).
With halogens: Form MX₄ for C, Si, Ge, Sn (CCl₄, SiCl₄); Pb forms only PbX₂ easily (PbCl₄ exists but is unstable).
With acids: C and Si do not react with dilute acids. Sn and Pb dissolve in HCl to give MCl₂; Pb passivates in dilute H₂SO₄ (PbSO₄ coat).
Hydrides: CH₄ (very stable) → SiH₄ → GeH₄ → SnH₄ → PbH₄ (extremely unstable). Stability decreases down the group as M–H bond weakens.
Catenation: C >> Si > Ge > Sn (Pb negligible). Strong C–C bonds are why organic chemistry is dominated by carbon.

Common trap. Beryllium chloride (BeCl₂) is covalent, not ionic, even though Be is in Group II. Be²⁺ is so small with such high charge density that it severely polarises the chloride ion (Fajans' rules). The "all Group II compounds are ionic" assumption fails here.

Memory aid. "Lead loves +2." The inert pair effect makes the lower oxidation state increasingly stable down Group IV, so Pb²⁺ compounds dominate while PbO₂ is a strong oxidiser.

Worked MCQs

Five MCQs that capture the high-yield testing patterns for this chapter. Read the explanation even when you get the answer right — it's where the deeper concept lives.

Q1. Across a period from left to right, atomic radius generally:

Decreases
Increases
Remains constant
First decreases then increases

Across a period, electrons are added to the same shell while nuclear charge increases. Effective nuclear charge rises and pulls outer electrons closer, so atomic radius decreases.

Q2. The order of reactivity of alkali metals with water is:

Li > Na > K > Rb > Cs
Cs > Rb > K > Na > Li
Na > K > Li > Rb > Cs
K > Na > Cs > Li > Rb

Going down Group I, atomic radius increases and ionisation energy decreases — the single ns¹ electron is lost more easily. Therefore reactivity with water rises Cs > Rb > K > Na > Li.

Q3. Which compound is predominantly covalent rather than ionic?

NaCl
KCl
BeCl₂
BaCl₂

Be²⁺ has very small size and high charge density. By Fajans' rules it strongly polarises Cl⁻, giving BeCl₂ a largely covalent character (chains in the solid, monomeric/dimeric in the vapour).

Q4. The increasing stability of the +2 oxidation state down Group IV (Pb²⁺ > Sn²⁺) is explained by:

Increase in electronegativity
Decrease in atomic radius
The inert pair effect
Greater hydration energy

The inert pair effect describes the reluctance of the ns² pair to participate in bonding for heavier p-block elements. As a result, the lower oxidation state (e.g. +2 for Pb) becomes increasingly stable relative to the group oxidation state (+4).

Q5. An element with electronic configuration [Ar] 3d¹⁰ 4s² 4p³ belongs to which block?

s-block
p-block
d-block
f-block

The last electron entered the 4p orbital, so the element is in the p-block (it is As, Group V). Block assignment depends on which orbital receives the differentiating electron, not on which orbitals are filled overall.

Quick Recap

Block = which orbital receives the last electron (s/p/d/f).
Atomic radius: ↓ across, ↑ down. Ionisation energy: ↑ across, ↓ down. Electronegativity follows IE.
Group I (ns¹): reactivity Cs > Rb > K > Na > Li. Strong reducers; vigorous with water giving MOH + H₂.
Group II (ns²): less reactive than Group I; BeCl₂ covalent (Fajans). Solubility of sulphates ↓ down; hydroxides ↑.
Group IV: C non-metal → Pb metal. Inert pair effect ⇒ Pb²⁺ stable, PbO₂ oxidiser. Hydride stability ↓ down (CH₄ > PbH₄).
Catenation strongest for carbon.

Test yourself. Take a timed S/P-Block quiz or browse all Chemistry MCQs to lock these concepts in.