Class 12 Chemistry · NCERT chapter 7 · 8 topics · free to read
Group 15: Nitrogen Family
Group 15 elements (N, P, As, Sb, Bi) have ns²np³ configuration with 5 valence electrons. Nitrogen is anomalous due to small size, high electronegativity, absence of d-orbitals, and ability to form pπ-pπ multiple bonds. Oxidation states range from −3 to +5. NH₃ is pyramidal, basic, and acts as a ligand. Oxides of nitrogen range from N₂O (+1) to N₂O₅ (+5) with varying acid-base character.
N_2O_5 + H_2O arrow 2HNO_3
Electronic Configuration & Properties
Group 15 elements have a half-filled p-subshell (ns²np³), which gives extra stability. This explains their high ionization enthalpies and reluctance to lose electrons easily.
All Group 15 elements have the general electronic configuration ns^2np^3 . The three p -electrons occupy three separate p -orbitals (one each in p_x , p_y , p_z ) following Hund's rule — giving a half-filled, extra-stable configuration.
Elements: N ( Z=7 ), P ( Z=15 ), As ( Z=33 ), Sb ( Z=51 ), Bi ( Z=83 )
Common oxidation states: -3, +3, +5
Stability of +5 state decreases down the group (inert pair effect) — Bi prefers +3
Nitrogen shows -3 to +5 oxidation states (all values from -3 to +5 are possible)
First ionization enthalpy is higher than Group 14 and Group 16 — half-filled stability
Anomalous Behavior of Nitrogen
Nitrogen differs from other Group 15 members due to its small size, high electronegativity, absence of d-orbitals, and unique ability to form pπ-pπ multiple bonds. It cannot expand its covalence beyond 4.
Nitrogen is the oddball of Group 15. Four reasons make it anomalous: 1. Small size — atomic radius of N is much smaller than P 2. High electronegativity — 3.04 (highest in Group 15) 3. No d-orbitals — cannot expand octet beyond 4 bonds 4. pπ-pπ bonding — forms N=N and N , which P/As/Sb/Bi cannot do effectively
Forms N triple bond (very strong)
Forms P_4 (single bonds, tetrahedron)
Max covalence = 4 (no d-orbitals)
Max covalence = 5 or 6 ( d -orbitals available)
NF_3 only — no NF_5
PF_3 AND PF_5 both exist
Exists as N_2 (diatomic gas)
Exists as P_4 (tetrahedral solid)
Forms strong H-bonds ( NH_3 )
No H-bonding in PH_3
Ammonia
NH₃ is prepared by Haber process (N₂ + 3H₂ → 2NH₃). It has a trigonal pyramidal shape (sp³, bond angle 107°), acts as a Lewis base due to the lone pair, and forms NH₄⁺ by donating this lone pair to H⁺.
Ammonia ( NH_3 ) is the most important hydride of nitrogen. It's manufactured industrially by the Haber process — one of the most significant chemical reactions in human history (Fritz Haber won the 1918 Nobel Prize for this).
N_2(g) + 3H_2(g) [Fe/Mo] 200 atm, 700 K 2NH_3(g) H = -92.4 kJ/mol
Shape: Trigonal pyramidal ( sp^3 hybridized, bond angle 107.8° )
Lone pair on N → Lewis base → donates to H^+ forming NH_4^+ (tetrahedral)
Test: Turns moist red litmus blue; gives white fumes with HCl ( NH_4Cl )
Quick recall
Why is the bond angle in NH₃ (107.8°) less than the tetrahedral angle (109.5°)?
The lone pair on nitrogen repels the bonding pairs more strongly (lp-bp > bp-bp repulsion), compressing the H-N-H angle from 109.5° to 107.8°. Same logic for H₂O (104.5°) — two lone pairs compress even more.
Oxides of Nitrogen
Nitrogen forms oxides in every oxidation state from +1 to +5: N₂O, NO, N₂O₃, NO₂, N₂O₄, N₂O₅. Their acid-base nature shifts from neutral to acidic as oxidation state increases. NO₂ is the brown gas responsible for photochemical smog.
Nitrogen forms six well-known oxides spanning oxidation states +1 to +5 . Each has unique properties, colors, and acid-base character. NEET tests all of them — especially their nature (acidic/neutral) and physical state.
N_2O (Nitrous oxide, +1)
Colorless gas, neutral, 'laughing gas', supporter of combustion
NO (Nitric oxide, +2)
Colorless gas, neutral, paramagnetic, turns brown in air → NO_2
N_2O_3 (Dinitrogen trioxide, +3)
Blue solid, acidic, anhydride of HNO_2
NO_2 (Nitrogen dioxide, +4)
Brown gas, acidic, paramagnetic, dimerizes to N_2O_4
N_2O_4 (Dinitrogen tetroxide, +4)
Colorless solid/liquid, acidic, diamagnetic (dimer of NO_2 )
N_2O_5 (Dinitrogen pentoxide, +5)
White solid, acidic, anhydride of HNO_3
NO_2 is brown (causes brown color of smog); N_2O_4 is colorless — heating shifts equilibrium toward brown NO_2
NO + 1 2 O_2 → NO_2 (colorless to brown in air — classic lab observation)
N_2O_3 is the acid anhydride of HNO_2 (nitrous acid)
N_2O_5 is the acid anhydride of HNO_3 (nitric acid)
N_2O supports combustion (relights a glowing splint) — don't confuse with O_2
Oxoacids of Nitrogen & Phosphorus
Nitrogen forms HNO₂ (nitrous acid) and HNO₃ (nitric acid). HNO₃ is prepared by Ostwald process and is a strong oxidizing acid. Phosphorus forms a bewildering array of oxoacids — H₃PO₄ (orthophosphoric), H₃PO₃ (phosphorous acid, dibasic!), H₃PO₂ (hypophosphorous acid, monobasic). Phosphine (PH₃) is a toxic, non-flammable gas. PCl₃ and PCl₅ are important halides with different geometries.
4NH_3 + 5O_2 Pt/Rh 4NO + 6H_2O
Nitric Acid (HNO₃)
HNO₃ is manufactured by the Ostwald process (NH₃ → NO → NO₂ → HNO₃). It is a strong monobasic acid and powerful oxidizing agent that attacks almost all metals except Au and Pt.
Nitric acid ( HNO_3 ) is prepared industrially by the Ostwald process — a three-step oxidation of ammonia. This is one of the most important industrial processes in NEET chemistry.
Step 1: 4NH_3 + 5O_2 [500 K] Pt/Rh 4NO + 6H_2O
Step 2: 2NO + O_2 arrow 2NO_2
Step 3: 3NO_2 + H_2O arrow 2HNO_3 + NO
Catalyst: Pt-Rh gauge (platinum-rhodium, 90:10 alloy) at ~500 K
Concentrated HNO₃ is a strong oxidizing agent — oxidizes most metals (except Au, Pt, Rh, Ir)
Aqua regia ( 3HCl + 1HNO_3 ) dissolves even gold and platinum
Brown ring test: FeSO_4 + NO arrow [Fe(H_2O)_5NO]^ 2+ — tests for NO_3^- ions
Structure of HNO_3 : N is sp^2 hybridized, one N=O and one N-OH
Oxoacids of Phosphorus
Phosphorus forms many oxoacids. The basicity depends on the number of P–OH bonds, NOT the total H atoms. H₃PO₄ is tribasic (3 P–OH), H₃PO₃ is dibasic (2 P–OH + 1 P–H), H₃PO₂ is monobasic (1 P–OH + 2 P–H).
Phosphorus forms a spectacular variety of oxoacids. The key NEET concept is this: the basicity of a phosphorus oxoacid equals the number of P–OH bonds (ionizable H atoms), NOT the total number of H atoms in the formula. Hydrogen atoms bonded directly to P ( P–H ) are non-ionizable.
H_3PO_4 (Orthophosphoric acid)
Tribasic — 3 P–OH bonds, 0 P–H bonds
H_3PO_3 (Phosphorous acid)
Dibasic — 2 P–OH bonds, 1 P–H bond
H_3PO_2 (Hypophosphorous acid)
Monobasic — 1 P–OH bond, 2 P–H bonds
H_4P_2O_7 (Pyrophosphoric acid)
Tetrabasic — 4 P–OH bonds
HPO_3 (Metaphosphoric acid)
Monobasic — 1 P–OH bond
H_3PO_4 → gives 3 series of salts: NaH_2PO_4 , Na_2HPO_4 , Na_3PO_4
H_3PO_3 → only 2 series: NaH_2PO_3 , Na_2HPO_3 (no Na_3PO_3 !)
H_3PO_3 is a strong reducing agent (due to P–H bond) — reduces Ag^+ to Ag
H_3PO_2 is an even stronger reducing agent (2 P–H bonds)
All phosphorus oxoacids have tetrahedral geometry around P ( sp^3 )
Quick recall
H₃PO₃ has 3 hydrogen atoms. Why is it dibasic and not tribasic?
Only 2 of the 3 H atoms are attached to oxygen (P–OH bonds → ionizable). The third H is directly bonded to phosphorus (P–H bond → non-ionizable). Basicity = number of P–OH bonds = 2.
Phosphine, PCl₃ & PCl₅
Phosphine (PH₃) is a toxic pyramidal gas. PCl₃ is pyramidal (sp³) and PCl₅ is trigonal bipyramidal (sp³d). PCl₅ exists because P has empty 3d orbitals for octet expansion — nitrogen cannot do this.
Phosphine ( PH_3 ) is the phosphorus analog of ammonia. It's far more toxic, less basic, and has a rotten fish smell. Unlike NH_3 , phosphine does NOT form hydrogen bonds (P is not electronegative enough).
PH_3 : pyramidal, bond angle 93.6° (less than NH_3 's 107.8° — less lp-bp repulsion due to longer bonds)
Group 16 elements (O, S, Se, Te, Po) have ns²np⁴ configuration. Oxygen is anomalous due to small size, high electronegativity, and absence of d-orbitals. Allotropes include O₂ (paramagnetic) and O₃ (diamagnetic, bent). Oxidation states: −2 (most common), +2, +4, +6. H₂SO₄ is the 'king of chemicals' made by the Contact process. SO₂ is an acidic gas used as a bleaching agent.
2SO_2 + O_2 V_2O_5 2SO_3
Electronic Configuration & Allotropy
Group 16 elements have ns²np⁴ configuration. Oxygen exhibits allotropy: O₂ (dioxygen, paramagnetic) and O₃ (ozone, diamagnetic, angular). The paramagnetic nature of O₂ is explained only by MOT, not VBT.
Group 16 elements are called chalcogens ('ore-forming elements'). They have the general configuration ns^2np^4 — two electrons short of a noble gas configuration, which is why -2 is the most common oxidation state.
Elements: O ( Z=8 ), S ( Z=16 ), Se ( Z=34 ), Te ( Z=52 ), Po ( Z=84 , radioactive)
Common oxidation states: -2 (all), +2, +4, +6 (S, Se, Te) — O shows -2 almost exclusively
Oxygen shows +2 in OF_2 only (F is more electronegative)
+6 state stability decreases down the group — inert pair effect
First two elements (O, S) are nonmetals; Se, Te are metalloids; Po is metallic
Allotropy of oxygen: O_2 (dioxygen) — colorless, odorless, paramagnetic (attracted to magnet). Bond order = 2. O_3 (ozone) — pale blue gas, pungent smell, diamagnetic, angular/bent shape ( O–O–O bond angle 117° ).
Anomalous Behavior of Oxygen
Oxygen differs from S, Se, Te due to small size, high electronegativity, absence of d-orbitals, and ability to form pπ-pπ multiple bonds. Max covalence of O is 2 (or 4 with coordination), unlike S which can be 6.
Just like nitrogen in Group 15, oxygen is anomalous in Group 16. The reasons are identical in principle — small size, high EN, no d-orbitals — but the consequences are different.
Exists as O_2 (double bond, diatomic)
Exists as S_8 (single bonds, ring)
Max covalence = 2 (no d-orbitals)
Max covalence = 6 ( SF_6 exists)
Forms strong O=O double bond ( p - p )
Forms weak S=S — prefers single bonds + chains
Very high electronegativity (3.44)
Lower electronegativity (2.58)
Strong hydrogen bonding in H_2O
No H-bonding in H_2S
H_2O is liquid at room temperature
H_2S is gas (b.p. -60°C )
OF_2 → O is +2 (only case where O has positive oxidation state, because F is more electronegative)
H_2O has anomalously high b.p. (100°C) vs H_2S ( -60°C ), H_2Se ( -41°C ), H_2Te ( -2°C )
Trend in hydride b.p.: H_2O H_2Te > H_2Se > H_2S (H-bonding anomaly for water)
Thermal stability of hydrides: H_2O > H_2S > H_2Se > H_2Te (decreasing bond strength)
Acidic strength of hydrides: H_2Te > H_2Se > H_2S > H_2O (larger atom → weaker H–X bond → easier dissociation)
Sulphuric Acid
H₂SO₄ is manufactured by the Contact process: S → SO₂ → SO₃ → oleum → H₂SO₄. It is a strong dibasic acid, powerful dehydrating agent, and oxidizing agent. It chars sugar, dehydrates organic compounds, and liberates volatile acids from their salts.
Sulphuric acid ( H_2SO_4 ) is called the 'king of chemicals' because a country's industrial development is often measured by its H₂SO₄ production. It's made by the Contact process.
S + O_2 arrow SO_2
2SO_2 + O_2 [720 K, 2 atm] V_2O_5 2SO_3
SO_3 + H_2SO_4 arrow H_2S_2O_7 (oleum)
H_2S_2O_7 + H_2O arrow 2H_2SO_4
Catalyst: V_2O_5 (vanadium pentoxide) at ~720 K, 2 atm pressure
SO₃ is NOT directly dissolved in water (forms mist) → dissolved in conc. H_2SO_4 to form oleum first
Hot conc. H_2SO_4 is an oxidizing agent: Cu + 2H_2SO_4 arrow CuSO_4 + SO_2 + 2H_2O
Structure of H_2SO_4 : S is sp^3 hybridized, tetrahedral, 2 S=O + 2 S–OH
SO₂ and H₂S
SO₂ is an acidic gas and reducing agent (bleaches by reduction). H₂S is a weak diprotic acid, foul-smelling ('rotten eggs'), and the most common reducing agent in qualitative analysis. SO₂ is the anhydride of H₂SO₃.
SO₂ and H₂S are the two most commonly tested sulfur compounds after H₂SO₄. Both are pungent gases, but they differ dramatically in their chemistry.
SO_2 bleaching is temporary (reduction) — color returns on exposure to air. Cl₂ bleaching is permanent (oxidation)
SO_2 + K_2Cr_2O_7 (acidified) → turns orange to green ( Cr^ 3+ ) — confirmatory test
H_2S test: turns lead acetate paper black ( PbS ) — used in qualitative analysis
H_2S is used to precipitate Group II ( Cu^ 2+ , Pb^ 2+ , As^ 3+ ) and Group IV ( Co^ 2+ , Ni^ 2+ , Zn^ 2+ ) cations in salt analysis
Oxoacids of Sulphur
Sulphur forms numerous oxoacids with S in oxidation states +2 to +6. Key ones: H₂SO₃ (sulphurous, +4), H₂SO₄ (sulphuric, +6), H₂S₂O₃ (thiosulphuric, +2), H₂S₂O₇ (oleum/pyrosulphuric), H₂S₂O₈ (peroxodisulphuric — Marshall's acid, contains S-O-O-S peroxide linkage). All have tetrahedral geometry around each S atom.
H_2SO_4 arrow H^+ + HSO_4^- arrow 2H^+ + SO_4^ 2-
Structures of Sulphur Oxoacids
All sulphur oxoacids have S in sp³ hybridization with tetrahedral geometry. The differences lie in the number of S atoms, presence of S–S or O–O (peroxo) bonds, and the oxidation state of S.
Sulphur forms a large family of oxoacids. The common thread: every S atom is tetrahedrally surrounded ( sp^3 ). NEET doesn't ask you to draw every structure from scratch, but you must know the key features that distinguish them.
H_2SO_3 (Sulphurous acid, S = +4)
1 S atom, pyramidal around S, 1 S=O + 2 S–OH
H_2SO_4 (Sulphuric acid, S = +6)
1 S atom, 2 S=O + 2 S–OH, tetrahedral
H_2S_2O_7 (Oleum/Pyrosulphuric)
2 S atoms linked by S–O–S bridge
H_2S_2O_3 (Thiosulphuric acid)
Like SO_4^ 2- but one O replaced by S — contains S–S bond
H_2S_2O_8 (Peroxodisulphuric/Marshall's acid)
2 S atoms linked by S–O–O–S peroxide bridge
H_2SO_5 (Caro's acid/Peroxomonosulphuric)
Contains one S–O–O–H peroxide bond
Thiosulphate ( S_2O_3^ 2- ): replace one O in sulphate with S → the two S atoms are NOT equivalent (one central, one terminal)
Na_2S_2O_3 (hypo) + I_2 → Na_2S_4O_6 + 2NaI — used in iodometric titrations
Marshall's acid ( H_2S_2O_8 ) is the strongest oxidizing sulphur oxoacid — used in Caro's acid preparation
All S atoms in all these oxoacids maintain sp^3 hybridization
Quick recall
What is the key structural difference between H₂S₂O₇ (oleum) and H₂S₂O₈ (Marshall's acid)?
Oleum ( H_2S_2O_7 ) has an S–O–S bridge (single O linking two SO₃ groups). Marshall's acid ( H_2S_2O_8 ) has an S–O–O–S peroxide bridge (two O atoms linking the SO₃ groups). The extra O in the bridge makes it a peroxo compound.
H₂SO₄ Properties
Concentrated H₂SO₄ acts as: (1) dehydrating agent (chars sugar, removes water from compounds), (2) oxidizing agent when hot and concentrated, (3) non-volatile acid that displaces volatile acids from salts.
H₂SO₄ wears three hats in chemistry — and NEET tests all three. Understanding which role it plays in a given reaction is crucial.
Dehydrating agent: Removes water from compounds — C_ 12 H_ 22 O_ 11 conc. H_2SO_4 12C + 11H_2O (chars sugar to carbon)
Halogens (F, Cl, Br, I) have ns²np⁵ configuration and are the most reactive nonmetals. Fluorine is anomalous (small size, no d-orbitals, highest EN). Oxidation states: −1 (all), +1 to +7 (Cl, Br, I, not F). HF is a weak acid despite F being most electronegative. Interhalogen compounds (ClF₃, BrF₅, IF₇) form due to size and EN differences. Bleaching powder (CaOCl₂) is industrially important.
Acid strength: HF HCl < HBr < HI
Halogen Properties & Trends
Halogens have highest electron gain enthalpies (except F < Cl due to small size). Reactivity, oxidizing power, and electronegativity decrease down the group. All are colored and toxic.
The halogens are the most reactive nonmetals — they need just one electron to complete their octet. Their general configuration is ns^2np^5 . They exist as diatomic molecules ( F_2, Cl_2, Br_2, I_2 ).
Elements: F ( Z=9 ), Cl ( Z=17 ), Br ( Z=35 ), I ( Z=53 ), At ( Z=85 , radioactive)
Physical states: F_2 = pale yellow gas, Cl_2 = greenish-yellow gas, Br_2 = reddish-brown liquid, I_2 = dark violet solid
Electron gain enthalpy (magnitude): Cl > F > Br > I — F is LESS than Cl due to small size and electron-electron repulsion
Electronegativity: F (3.98) > Cl (3.16) > Br (2.96) > I (2.66)
Oxidizing power: F_2 > Cl_2 > Br_2 > I_2 (F₂ is the strongest oxidizing agent among all elements)
Bond dissociation energy: Cl–Cl > Br–Br > F–F > I–I — F–F is anomalously weak!
Quick recall
Why does Cl have a higher electron gain enthalpy than F, even though F is more electronegative?
Fluorine is so tiny that adding an electron to the already electron-dense 2p subshell creates severe electron-electron repulsion. Chlorine's larger 3p orbitals accommodate the incoming electron more comfortably. So: _ eg H of Cl ( -349 kJ/mol) > F ( -328 kJ/mol).
Anomalous Behavior of Fluorine
Fluorine differs from other halogens: only −1 oxidation state (no d-orbitals), highest EN, forms H-bonds in HF, anomalously weak F–F bond, and cannot form polyhalide ions. HF is a weak acid despite F being most electronegative.
Fluorine breaks almost every trend you'd predict. It's the most electronegative element on the periodic table, yet its bond with itself is surprisingly weak, and HF is a weak acid. Here's why.
Acid strength of hydrogen halides: HI > HBr > HCl HF (decreasing bond strength)
Thermal stability: HF > HCl > HBr > HI (HI decomposes on heating)
Reducing power of HX : HI > HBr > HCl > HF (HI is strongest reducing agent)
Boiling points: HF (19.5°C) HI ( -35°C ) > HBr ( -67°C ) > HCl ( -85°C ) — HF anomaly due to H-bonding
Interhalogen Compounds
Interhalogen compounds (AB, AB₃, AB₅, AB₇) form between two different halogens where A is the larger, less electronegative atom. Shapes: AB=linear, AB₃=T-shaped, AB₅=square pyramidal, AB₇=pentagonal bipyramidal.
When two different halogens combine, they form interhalogen compounds. The larger halogen (A) is always the central atom, bonded to an odd number of smaller halogen atoms (B). The types are: AB, AB_3, AB_5, AB_7 .
All interhalogens are more reactive than constituent halogens (except F_2 )
They are covalent, diamagnetic, and volatile
Hydrolysis: AB_n + n 2 H_2O arrow HA + nHB (approximately)
ClF_3 — T-shaped due to 2 lone pairs on Cl ( sp^3d )
IF_7 — the ONLY AB_7 interhalogen; pentagonal bipyramidal ( sp^3d^3 )
A = larger, less EN halogen; B = smaller, more EN halogen (always!)
Chlorine & Bleaching Powder
Cl₂ is prepared industrially by electrolysis of brine (Castner-Kellner cell). Lab preparation from MnO₂ + conc. HCl. Bleaching powder (CaOCl₂) is prepared by passing Cl₂ over Ca(OH)₂. It bleaches by oxidation and is used for water purification.
Chlorine ( Cl_2 ) is the most industrially important halogen. Its lab and industrial preparations, plus bleaching powder chemistry, are NEET essentials.
Active bleaching agent in bleaching powder: OCl^- (hypochlorite ion)
Available chlorine in fresh bleaching powder: 35 %
Bleaches by oxidation
Bleaches by reduction
Permanent bleaching
Temporary bleaching (color returns)
Destroys the dye molecule
Reduces the dye (can re-oxidize)
Used for cotton, paper
Used for delicate materials like silk, wool
Oxoacids of Halogens
Halogens form four series of oxoacids: HOX (hypohalous, +1), HXO₂ (halous, +3), HXO₃ (halic, +5), HXO₄ (perhalic, +7). Acid strength increases with oxidation state (more O atoms stabilize the conjugate base). Fluorine forms only HOF. Oxidizing power decreases as acid strength increases in the series HOCl > HClO₂ > HClO₃ > HClO₄.
Acid strength: HOCl < HClO_2 < HClO_3 < HClO_4
Four Series of Oxoacids
Chlorine forms all four types: HOCl (+1), HClO₂ (+3), HClO₃ (+5), HClO₄ (+7). Bromine and iodine form most of them. Fluorine forms only HOF (unstable). The number of terminal oxygen atoms determines both acid strength and naming.
Halogens (except F) form oxoacids in four series, each with increasing oxidation state of the halogen. Chlorine is the model example — it forms all four types.
HClO_4 (perchloric acid) is the strongest known oxyacid — even stronger than H_2SO_4
F forms only HOF (fluoric acid, very unstable) — because F has no d-orbitals for higher oxidation states
Among hypohalous acids: HOCl > HOBr > HOI (acid strength decreases as halogen size increases)
Structures & Naming Logic
In each oxoacid, the halogen is at the center bonded to O–H groups and terminal O atoms. Acid strength depends on the number of terminal (non-OH) oxygen atoms. More terminal O atoms = stronger acid = more stable conjugate base.
Each halogen oxoacid has the halogen at the center, bonded to one O–H group and varying numbers of terminal oxygen atoms (=O, double-bonded to X). The terminal oxygens are the key — they pull electron density away from the O–H bond.
HOCl (hypochlorous)
0 terminal O atoms, Cl bonded to O–H only, sp^3
HClO_2 (chlorous)
1 terminal O atom, sp^3
HClO_3 (chloric)
2 terminal O atoms, sp^3
HClO_4 (perchloric)
3 terminal O atoms, sp^3 , tetrahedral around Cl
Terminal O rule: More terminal O atoms → stronger inductive effect → weaker O–H bond → stronger acid
ClO_4^- (perchlorate) is extremely stable — charge delocalized over 4 equivalent O atoms
Periodate ( IO_4^- ) exists, but iodine can also form H_5IO_6 (paraperiodic acid) — I is large enough to coordinate 6 O atoms
All these oxoacids are strong oxidizing agents — they find use in disinfection, bleaching, and as rocket propellants ( NH_4ClO_4 )
Bleaching agents: NaOCl (sodium hypochlorite) is the active ingredient in household bleach
Quick recall
HClO₄ is a stronger acid than HClO₃. But which is a stronger oxidizing agent?
HClO_3 is the stronger oxidizing agent. Oxidizing power DECREASES as acid strength INCREASES for halogen oxoacids. The lower oxidation state species ( HOCl, HClO_2 ) are the strongest oxidizers because they gain stability by being further oxidized or reduced.
Group 18: Noble Gases
Noble gases have completely filled ns²np⁶ configuration (He: 1s²). They were called 'inert' until Neil Bartlett synthesized XePtF₆ in 1962. Xenon forms compounds with F and O: XeF₂ (linear, sp³d), XeF₄ (square planar, sp³d²), XeF₆ (distorted octahedral, sp³d³), XeO₃ (pyramidal, sp³). Lighter noble gases (He, Ne, Ar) form no stable compounds under normal conditions.
Xe + 2F_2 400°C, Ni XeF_4
Noble Gas Properties
Noble gases have completely filled outer shells, giving them exceptional stability, high ionization enthalpies, and no tendency to form bonds. Bartlett (1962) shattered the 'inert' myth by making XePtF₆.
For over a century, Group 18 elements were called 'inert gases' — believed to be completely unreactive. Then in 1962, Neil Bartlett noticed that O_2 (IE = 1175 kJ/mol) reacted with PtF_6 . Since Xe has a similar IE (1170 kJ/mol), he reasoned Xe should react too.
Elements: He ( Z=2 ), Ne ( Z=10 ), Ar ( Z=18 ), Kr ( Z=36 ), Xe ( Z=54 ), Rn ( Z=86 , radioactive)
Configuration: ns^2np^6 (He: 1s^2 ) — completely filled, most stable
All are monatomic gases — no intermolecular bonding needed
Ionization enthalpies are the highest in each period → very reluctant to lose electrons
Boiling points increase down the group: He ( -269°C ) < Ne < Ar < Kr < Xe ( -108°C ) < Rn — increasing van der Waals forces
He has the lowest boiling point of any substance ( -269°C = 4.2 K)
Xenon Fluorides
Xenon forms three fluorides by direct reaction with F₂ at different conditions. XeF₂ is linear (sp³d, 3 lone pairs), XeF₄ is square planar (sp³d², 2 lone pairs), XeF₆ has distorted octahedral geometry (sp³d³, 1 lone pair).
Xenon reacts directly with F_2 to form three fluorides. The product depends on the Xe:F₂ ratio and reaction conditions.
Xe + F_2 [400°C, Ni] 1:1 XeF_2
Xe + 2F_2 [400°C, Ni] 1:2 XeF_4
Xe + 3F_2 [300°C, 60-70 atm] 1:5 XeF_6
XeF_2 (2 bp + 3 lp)
sp^3d — Linear shape, 3 lone pairs equatorial in TBP arrangement
For XeF_6 preparation, excess F_2 (1:5 ratio) is used at higher pressure (60-70 atm)
Xenon-Oxygen Compounds
Xenon forms oxide and oxyfluoride compounds. XeO₃ is pyramidal (sp³), XeOF₂ is T-shaped (sp³d), XeOF₄ is square pyramidal (sp³d²). XeO₃ is dangerously explosive. These form by hydrolysis of xenon fluorides.
Xenon-oxygen compounds are formed by partial hydrolysis of xenon fluorides. They contain Xe=O double bonds and are generally more reactive (and more dangerous!) than the fluorides.
XeO_3 (3 bp + 1 lp)
sp^3 — Pyramidal (like NH_3 )
XeOF_2 (3 bp + 2 lp)
sp^3d — T-shaped (like ClF_3 )
XeOF_4 (5 bp + 1 lp)
sp^3d^2 — Square pyramidal (like BrF_5 )
XeO_3F_2 (5 bp + 0 lp)
sp^3d — Trigonal bipyramidal
XeO_3 is formed by complete hydrolysis of XeF_6 : XeF_6 + 3H_2O arrow XeO_3 + 6HF
XeO_3 is highly explosive — detonates at the slightest mechanical shock
XeOF_4 is formed by partial hydrolysis of XeF_6 : XeF_6 + H_2O arrow XeOF_4 + 2HF
XeOF_2 is formed by partial hydrolysis of XeF_4 : XeF_4 + H_2O arrow XeOF_2 + 2HF
In all Xe=O bonds, the oxygen is double-bonded to Xe
Quick recall
What is the shape of XeOF₂ and what is the hybridization of Xe?
XeOF₂ has 3 bond pairs (1 Xe=O + 2 Xe–F) and 2 lone pairs around Xe. Total = 5 electron domains → sp^3d hybridization → T-shaped geometry (like ClF₃). The lone pairs occupy equatorial positions in the trigonal bipyramidal arrangement.
Comparative Trends Across Groups 15–18
Systematic comparison of periodic trends across Groups 15–18: electronegativity, ionization enthalpy, oxidation states, metallic character, and the anomalous behavior of the first element in each group (N, O, F, Ne). The inert pair effect becomes prominent for heavier elements. Hydride stability decreases down each group while acidic character increases.
IE_1: Group 18 > Group 17 > Group 15 > Group 16
Electronegativity & Ionization Enthalpy
Electronegativity and IE₁ generally decrease down each group. Across groups 15→18 in a period, IE follows: Group 18 > 17 > 15 > 16 (the half-filled bump at Group 15 and noble gas stability at Group 18).
When comparing trends across Groups 15–18, two key anomalies break the expected left-to-right increase in IE:
IE₁ across period: Group 18 (fully filled) > Group 17 > Group 15 ( np^3 half-filled) > Group 16 ( np^4 , one electron beyond half-filled)
The 'dip' at Group 16: IE₁ of O < N, S < P — because removing one electron from the np^4 configuration gives the extra-stable np^3
Down each group: IE₁ decreases (larger atoms, electrons farther from nucleus)
Electronegativity across a period: F > O > N (Group 17 > 16 > 15) — follows expected trend
Electronegativity down each group: decreases (N > P > As > Sb > Bi, etc.)
Electronegativity
Increases: Group 15 < 16 < 17 (Group 18 not usually compared)
IE₁
Group 18 > 17 > 15 > 16 (anomaly at Group 15 and 16)
Electron Gain Enthalpy
Group 17 most negative (highest magnitude), Group 18 ≈ 0 (positive)
Decreases left to right: Group 15 has some metals, Group 18 has none
Oxidation States & Inert Pair Effect
The highest oxidation state equals the group number minus 10 (e.g., Group 15 → +5, Group 17 → +7). The inert pair effect makes the lower oxidation state more stable for heavier elements: Bi (+3 > +5), Pb (+2 > +4), Tl (+1 > +3).
Each group has a characteristic range of oxidation states. The maximum oxidation state equals the number of valence electrons. But the stability of the highest oxidation state changes dramatically down the group.
Group 15 ( ns^2np^3 )
-3, +3, +5 . Heaviest (Bi): +3 is stable, +5 is rare → inert pair effect
-1, +1, +3, +5, +7 . F shows only -1 (no d-orbitals)
Group 18 ( ns^2np^6 )
0 (mostly). Xe shows +2, +4, +6, +8 in fluorides/oxides
The inert pair effect is the reluctance of the ns^2 electrons to participate in bonding for heavier elements. This happens because poor shielding by d and f electrons increases the effective nuclear charge, holding the s -electrons tighter.
N : shows all oxidation states from -3 to +5 (most versatile in Group 15)
Bi : +3 is stable, Bi_2O_5 is a strong oxidizing agent (wants to go back to +3 )
For Group 17: F shows ONLY -1 (most EN element, no d-orbitals)
Positive oxidation states in halogens: Cl, Br, I show +1, +3, +5, +7
HI is unstable and a strong reducing agent — I readily goes to higher oxidation states
Anomalous Behavior of First Element
The first element of each group (N, O, F, Ne) is anomalous due to: small size, high electronegativity, absence of d-orbitals, and ability to form pπ-pπ multiple bonds. This affects their covalence, allotropy, hydride properties, and compound formation.
A universal rule across Groups 15–18: the first element is always anomalous. The causes are the same four factors in each group, but the consequences differ.
N (Group 15)
Forms N_2 (triple bond) vs P_4 (single bonds). Max covalence = 4. No NF_5 .
O (Group 16)
Forms O_2 (double bond) vs S_8 (single bonds). Max covalence = 2. No OF_6 .
F (Group 17)
Only -1 oxidation state. HF is weak acid. F-F bond anomalously weak.
Ne (Group 18)
Forms NO compounds at all (IE too high). Kr, Xe form fluorides.
Common cause: No d-orbitals in the second period ( n=2 ) → cannot expand octet
Multiple bonding: N, O, F form strong p - p bonds; heavier elements prefer single bonds + chains/rings
H-bonding: NH_3 , H_2O , HF all form H-bonds — their heavier analogs don't
Metallic character: First element of each group is always nonmetallic; metallic character increases down the group
Quick recall
Give two consequences of the absence of d-orbitals for each: nitrogen, oxygen, fluorine.
Nitrogen: (1) Cannot form NF₅ — max covalence = 4, (2) Forms N≡N (pπ-pπ bond) instead of N₄ chains. Oxygen: (1) Cannot form OF₆ — max covalence = 2, (2) Forms O=O instead of O₈ rings. Fluorine: (1) Shows only −1 oxidation state — cannot expand octet, (2) F–F bond is anomalously weak due to lone pair repulsion at short distance.
Hydride Trends
Down each group: hydride thermal stability decreases, acidic character increases, reducing power increases, and boiling points generally decrease (except the first member which shows H-bonding anomaly). These four trends are interconnected through bond strength.
Hydrides of p-block elements show beautifully systematic trends — and every trend connects back to one factor: bond strength between H and the central atom.
Thermal stability
Decreases — larger atom → weaker E–H bond → easier to break
Acidic character (in water)
Increases — weaker E–H bond → easier to release H^+
Reducing power
Increases — weaker bond → more willing to donate electrons
Boiling point
Generally decreases (less van der Waals) BUT first member is anomalously high (H-bonding)
Bond dissociation energy
Decreases — N-H > P-H > As-H > Sb-H > Bi-H
Group 15 stability: NH_3 > PH_3 > AsH_3 > SbH_3 > BiH_3 ( BiH_3 decomposes at room temperature)
Group 16 acidity: H_2Te > H_2Se > H_2S > H_2O
Group 17 acid strength: HI > HBr > HCl HF
Reducing power (Group 17): HI > HBr > HCl > HF — conc. H_2SO_4 oxidizes HI → I_2 and HBr → Br_2 , but NOT HCl