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Chemistryp-Block Elements
Identify the incorrect statement from the following :
Options
1
Carbon has the ability to form pπ-pπ multiple bond with itself.
2
ECl₃ (E = B and Al) is a monomer when E = B and a dimer when E = Al.
3
Oxygen exhibits only −2 oxidation state.
4
The order of catenation of Group 14 elements is C >> Si > Ge ≈ Sn.
Incorrect Statement (Answer)
Option 3 — Oxygen shows MORE than just −2
Why Option 3 is INCORRECT
1

Claim: "Oxygen exhibits ONLY −2 oxidation state."

This is FALSE. Oxygen shows multiple oxidation states:

❌ WRONG statement: Oxygen shows ONLY −2 oxidation state

✅ CORRECT fact: Oxygen shows −2, −1, 0, and +2 oxidation states

2

Oxidation states of Oxygen with examples:

−2: H₂O, CO₂, MgO, Na₂O (most common)

−1: H₂O₂ (hydrogen peroxide), Na₂O₂ (sodium peroxide)

0: O₂ (molecular oxygen, elemental state)

+2: OF₂ (oxygen difluoride — F is more electronegative than O)

3

Why other options are CORRECT:

✅ Option 1: Carbon forms C=C, C≡C (pπ-pπ bonds) — correct

✅ Option 2: BCl₃ = monomer (B completes octet via backbonding); AlCl₃ = dimer Al₂Cl₆ — correct

✅ Option 4: Catenation: C >> Si > Ge ≈ Sn — correct (C-C bonds very strong)

Theory: Oxidation States of p-Block Elements
1. Oxidation States of Oxygen — Complete List

📌 −2: H₂O, MgO, CO₂, SO₃ — most common; O is more electronegative than most elements

📌 −1: H₂O₂ (hydrogen peroxide), Na₂O₂ (sodium peroxide), BaO₂ (barium peroxide) — peroxides contain O-O single bond

📌 0: O₂ (elemental oxygen) — zero in its natural form

📌 +2: OF₂ (oxygen difluoride) — F is more electronegative than O, so O is positive here

📌 +1: O₂F₂ (dioxygen difluoride) — exotic, unstable compound

📌 Oxygen does NOT show +4 or +6 because it lacks d orbitals (Period 2)

📌 Key rule: Fluorine is ALWAYS −1 (most electronegative); Oxygen is usually −2 but not in OF₂

2. Why Does Oxygen NOT Show Positive Oxidation States (Except with F)?

Oxygen is the second most electronegative element (3·44 on Pauling scale) after fluorine (3·98). Oxygen can only show positive oxidation state when bonded to fluorine (the only element more electronegative than oxygen). In OF₂: F pulls electrons toward itself → O becomes electron-deficient (+2). In all other compounds (with C, H, N, S, metals), oxygen is more electronegative → gets negative oxidation state. Oxygen also cannot expand its octet (no d orbitals, Period 2) → maximum valency = 2. This contrasts with S (Period 3) which shows +2, +4, +6 using d orbitals.

3. Structure of Hydrogen Peroxide (H₂O₂)

H₂O₂ has oxidation state of O = −1. Structure: H-O-O-H (each O bonded to one H and one O). It has a non-planar structure with a dihedral angle of about 111° in gas phase. H₂O₂ is: a weak acid (Ka ≈ 2·4×10⁻¹²), a good oxidising agent (bleaching, antiseptic), and a reducing agent in strongly oxidising conditions (e.g., with KMnO₄). It decomposes: 2H₂O₂ → 2H₂O + O₂ (accelerated by catalysts like MnO₂). Commercially available as 3% (antiseptic), 30% (laboratory), 90% (rocket propellant). Pure H₂O₂ is pale blue, syrupy liquid.

4. Catenation — Ability to Form Self-Chains

Catenation is the ability of an element to form bonds with itself in long chains, rings, or networks. Carbon shows highest catenation because C-C bond is very strong (347 kJ/mol) and stable. Order: C >> Si > Ge ≈ Sn > Pb. Silicon: forms Si-Si bonds in silanes (SiₙH₂ₙ₊₂) but max ~10 Si atoms. Carbon: forms chains up to millions of atoms (polymers, DNA). Why C is special: small size → strong C-C σ bond + ability to form pπ-pπ bonds (alkenes, alkynes). Oxygen: very weak O-O bond (only in peroxides) → poor catenation. Sulphur: moderate S-S bond → chains in S₈ ring, polysulphides.

5. BCl₃ vs AlCl₃ — Monomer vs Dimer

BCl₃ exists as monomer: Boron has empty p orbital → accepts lone pair from Cl via pπ-pπ backbonding (Cl lone pair → B empty p) → partial double bond character → planar structure → no tendency to dimerize. Structure: trigonal planar. AlCl₃ exists as dimer (Al₂Cl₆): Al has empty d orbitals but weaker backbonding with Cl. Al is Lewis acid → forms dimer where each Al atom accepts lone pair from bridging Cl of the other Al. Structure: two AlCl₄ tetrahedra sharing an edge (Cl-Al-Cl-Al bridge). In gas phase at very high T, Al₂Cl₆ dissociates to AlCl₃.

6. Oxidation States of Sulphur

📌 −2: H₂S, Na₂S (sulphides)

📌 0: S (elemental, S₈ ring)

📌 +2: SCl₂, SF₂ (rare)

📌 +4: SO₂, H₂SO₃, SF₄, SOCl₂ (thionyl chloride)

📌 +6: SO₃, H₂SO₄, SF₆, SO₂Cl₂ (sulphuryl chloride)

📌 S shows +4 and +6 because it has d orbitals (Period 3) — unlike O

📌 This is the key difference between O and S: O max valency = 2; S max valency = 6

7. Oxidation States of Nitrogen

Nitrogen shows all oxidation states from −3 to +5: −3 (NH₃, amines), −2 (N₂H₄ hydrazine), −1 (NH₂OH hydroxylamine), 0 (N₂), +1 (N₂O laughing gas), +2 (NO), +3 (HNO₂, NO₂⁻, NCl₃), +4 (NO₂, N₂O₄), +5 (HNO₃, NO₃⁻). Unlike P, N cannot show +5 valency in fluorides (NF₅ doesn't exist) because N has no d orbitals — maximum valency = 4 in NH₄⁺ (using all 4 valence orbitals). But in NO₃⁻ and N₂O₅, N is +5 through oxidation state counting without violating octet rule.

8. Carbon — Why It Shows Maximum Catenation

Three reasons: (1) C-C bond energy = 347 kJ/mol (very high) → stable chains. Compare: Si-Si = 226 kJ/mol, Ge-Ge = 188 kJ/mol. (2) Carbon is small (covalent radius 77 pm) → good orbital overlap. (3) Carbon forms both σ and π bonds (pπ-pπ) → can make C=C and C≡C without d orbitals. This allows aromatic rings, alkenes, alkynes — adding enormous structural diversity. Life on Earth is carbon-based precisely because of this unique combination of strong C-C bonds + diverse bonding (single, double, triple) + ability to bond to H, O, N, S, halogens.

Frequently Asked Questions
1. What is the oxidation state of O in OF₂ and why is it +2? ⌄
In OF₂: F is more electronegative than O (F=3·98, O=3·44 on Pauling scale). So electrons are pulled toward F. Let O = x: x + 2(−1) = 0 → x = +2. This is the ONLY compound where O shows +2 oxidation state. OF₂ is called oxygen difluoride (not fluorine oxide) because O is the central atom. It's a highly reactive, toxic gas. This shows that oxidation state depends on relative electronegativity — whoever pulls electrons more gets the negative oxidation state.
2. What is the oxidation state of O in H₂O₂? ⌄
In H₂O₂: Let O = x. H is +1 (in non-metal compounds). So: 2(+1) + 2x = 0 → 2x = −2 → x = −1. Each O has oxidation state −1. This makes sense because in H-O-O-H, each O is bonded to one H (O takes electrons from H: O is −1 relative to H) and one O (no electron transfer between identical atoms). The O-O bond is called peroxide bond. H₂O₂ can act as both oxidising agent (getting reduced to −2: H₂O) and reducing agent (getting oxidised to 0: O₂).
3. Why can't oxygen show +4 or +6 oxidation states? ⌄
Oxygen is in Period 2: electronic configuration [He]2s²2p⁴. It has no d orbitals in its valence shell. To show +4 or +6 oxidation state, electrons would need to be promoted to d orbitals — but there are no d orbitals available at the n=2 level. This limits oxygen's valency to 2. Compare sulphur (Period 3): has 3d orbitals available → can show +4 (SO₂) and +6 (SO₃, H₂SO₄). This is the key periodic trend: Period 2 elements (C, N, O, F) cannot expand their octet.
4. What is pπ-pπ backbonding in BCl₃? ⌄
Boron in BCl₃ has an empty 2p orbital (only 6 electrons around B — incomplete octet). The Cl atoms have lone pairs in their 3p orbitals. The Cl lone pair donates into the empty B 2p orbital → pπ-pπ backbonding. This gives BCl₃ partial double bond character (B-Cl bond order slightly > 1). Effect: BCl₃ is trigonal planar (not pyramidal), and Cl lone pairs are delocalised → less electron density on Cl → Cl is less "available" to bridge between two B atoms → BCl₃ remains monomeric. AlCl₃ has weaker backbonding (3p→3p less effective) → Al remains Lewis acidic → dimerizes.
5. What is the structure of Al₂Cl₆? ⌄
Al₂Cl₆ (aluminium chloride dimer): two AlCl₃ units sharing two bridging Cl atoms. Structure: each Al has 4 Cl ligands (2 terminal + 2 bridging). Both Al atoms become tetrahedral. The bridging Cl atoms donate lone pairs to the adjacent Al atom → coordinate/dative bonds. In Al₂Cl₆: Al-Cl-Al bridge bonds are longer/weaker than terminal Al-Cl bonds. At high temperature (~400°C), Al₂Cl₆ dissociates to monomeric AlCl₃. In aqueous solution, Al³⁺ exists as [Al(H₂O)₆]³⁺. AlCl₃ is widely used as Lewis acid catalyst in Friedel-Crafts reactions.
6. What is the oxidation state of S in thiosulphate (S₂O₃²⁻)? ⌄
Thiosulphate ion S₂O₃²⁻: 2x + 3(−2) = −2 → 2x = 4 → x = +2. But this is an average! The two S atoms are NOT equivalent: one S is at +5 (central S bonded to 3 O) and one is at −1 (terminal S, like in sulphide). Average = (+5 + (−1))/2 = +2. This is called a "mean oxidation state." Similarly in S₄O₆²⁻ (tetrathionate): average S = +2.5, but individual S atoms have different oxidation states. This concept of non-equivalent atoms in a molecule having different oxidation states is important in NEET.
7. Which element shows maximum number of oxidation states and why? ⌄
Manganese (Mn, Z=25) shows the maximum number of oxidation states: +1 to +7 (and 0, −1 in some carbonyls). This is because Mn ([Ar]3d⁵4s²) can use all 7 valence electrons (5 in 3d + 2 in 4s) → maximum is +7 (in KMnO₄, MnO₄⁻). Among non-metals: Cl shows −1, 0, +1, +3, +5, +7 (6 states). Among p-block: S shows −2, 0, +2, +4, +6 (5 states). The more valence electrons and available d orbitals an element has, the more oxidation states it can display.
8. Why is carbon unique among all elements for forming life? ⌄
Carbon's uniqueness for life: (1) Strong, stable C-C bonds → long chains possible (polymers, proteins, DNA). (2) Forms C-H, C-O, C-N, C-S bonds → diverse functional groups. (3) Tetravalent → 3D structures possible (diamond, proteins). (4) Forms single, double, triple bonds → alkanes, alkenes, alkynes, aromatics. (5) Moderate reactivity — C-C bonds don't hydrolyse in water (unlike Si-O-Si). (6) Small atomic size → excellent orbital overlap → strong bonds. Silicon is often called the "next best" but Si-O bonds are very strong (hydrolysed easily) making long chains difficult in aqueous environments.
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