The p-Block ElementsNEET Chemistry · Class 12 · NCERT Chapter 13

Group 15: Nitrogen Family

Group 15 elements are nitrogen (N), phosphorus (P), arsenic (As), antimony (Sb), and bismuth (Bi). The general valence shell configuration is ns² np³, with three half-filled p orbitals. This half-filled configuration gives these elements extra stability, which is why they have higher ionization energies than their neighbours (Group 14 and Group 16).

Key Physical Properties

Oxidation States

All Group 15 elements show oxidation states of -3, +3, and +5. The -3 state is found in ionic hydrides and covalent hydrides (NH3, PH3). The +3 and +5 states are found in halides and oxoacids.

Important trend: inert pair effect. Going down the group, the +5 oxidation state becomes less stable and +3 becomes more stable. This is because the ns² electrons become increasingly reluctant to participate in bonding as you go down the group (due to poor shielding by the intervening d and f electrons). Bismuth shows +3 almost exclusively in its stable compounds.

Why Nitrogen is Anomalous

Nitrogen differs from the rest of Group 15 in several ways:

• No pentavalent halides: N cannot form NF5 or NCl5 because nitrogen is in period 2 and has no d orbitals. It cannot expand its octet. Phosphorus can form PCl5 because P has available 3d orbitals (sp3d hybridization).
• Forms p-p pi bonds: N forms strong multiple bonds (N≡N, C=N, N=O) using 2p orbitals that overlap effectively. Heavier elements (P, As) prefer single bonds and show catenation differently.
• No d orbitals: N cannot act as a Lewis acid in certain reactions where P can (e.g., P forms adducts by using d orbitals).
• High ionization energy: The half-filled 2p³ configuration of N is extra-stable, giving N a higher first ionization energy than even O.

Trends in Hydrides: NH3 to BiH3

NH3 has the highest boiling point because it forms hydrogen bonds (N-H...N). Bond angle decreases down the group as the lone pair is in a larger, more diffuse orbital and pushes less. Stability decreases down because the M-H bond strength decreases. Basicity decreases because the lone pair is less available (larger, more diffuse orbitals in heavier elements).

Nitrogen: Compounds and Reactions

Dinitrogen (N2)

Nitrogen exists as N2, held together by a very strong triple bond (N≡N, bond enthalpy = 945 kJ/mol). This makes N2 extremely unreactive at room temperature. To "fix" nitrogen (convert it to useful compounds), you need very high energy or special catalysts.

Haber Process (Synthesis of Ammonia)

The Haber process converts atmospheric nitrogen to ammonia, which is essential for fertilisers:

N2(g) + 3H2(g) ⇌ 2NH3(g) ΔH = -92 kJ/mol (exothermic)

Properties of Ammonia (NH3)

Ammonia is a colourless, pungent gas with a characteristic smell. It is highly soluble in water: NH3 + H2O ⇌ NH4+ + OH- (alkaline solution). This makes litmus paper turn blue in the presence of NH3 vapour.

Reducing property: 4NH3 + 5O2 → 4NO + 6H2O (catalytic oxidation with Pt catalyst, used in Ostwald process for HNO3 manufacture).

Complexing property: NH3 acts as a Lewis base (electron pair donor) and forms complexes: Cu2+ + 4NH3 → [Cu(NH3)4]2+ (deep blue colour, Schweitzer's reagent).

Oxides of Nitrogen

Nitric Acid (HNO3): Ostwald Process

HNO3 is made industrially by the Ostwald process:

1. Step 1: 4NH3 + 5O2 → 4NO + 6H2O (Pt catalyst, 850°C)
2. Step 2: 2NO + O2 → 2NO2
3. Step 3: 4NO2 + O2 + 2H2O → 4HNO3

Concentrated HNO3 is a powerful oxidising agent. It reacts with most metals (except platinum and gold). With copper: 3Cu + 8HNO3(dil.) → 3Cu(NO3)2 + 2NO + 4H2O. With concentrated HNO3: Cu + 4HNO3(conc.) → Cu(NO3)2 + 2NO2 + 2H2O. Zinc dissolves in very dilute HNO3 to give N2O or NH4NO3. Aqua regia (3 parts HCl + 1 part HNO3 conc.) dissolves gold and platinum.

Oxoacids of Nitrogen

Phosphorus: Allotropes and Compounds

Allotropes of Phosphorus

Phosphorus has several allotropes. The three main ones tested in NEET are:

White Phosphorus (P4)

White phosphorus consists of discrete P4 molecules with a tetrahedral structure. Each P atom is bonded to the other three, with bond angles of only 60°. This causes severe ring strain (normal sp3 bond angle is 109.5°), making white P highly reactive.

• Waxy, white/yellow solid. Translucent.
• Soft. Melting point 44°C.
• Extremely toxic. Causes severe burns.
• Ignites spontaneously in air above 30°C (pyrophoric).
• Glows in the dark (chemiluminescence) due to slow oxidation.
• Stored under water to prevent contact with air.
• Soluble in CS2 (carbon disulfide).

Red Phosphorus

Red phosphorus is formed by heating white P at 250°C in the absence of air. It is a polymeric form where P4 units are linked into long chains by covalent bonds. There are no discrete small molecules.

• Red, amorphous powder. Denser than white P.
• Stable in air up to 260°C. Does not ignite spontaneously.
• Non-toxic. Used on the striking surface of safety match boxes.
• Insoluble in CS2 (unlike white P).
• Melting point: decomposes at about 400°C (converts back to white P then vaporises).

Black Phosphorus

Black phosphorus is the most stable and least reactive allotrope. It is formed by heating white phosphorus under very high pressure. It has a layered structure similar to graphite, with each P atom bonded to three neighbours in corrugated layers.

• Black solid, resembles graphite in appearance.
• Semiconductor. Layers held by van der Waals forces.
• Thermodynamically most stable. Least reactive.
• A 2D layer called phosphorene (analogous to graphene) is obtained from black P.

Halides of Phosphorus

Phosphorus forms two important chlorides: PCl3 and PCl5.

PCl3 (Phosphorus trichloride)

P is sp3 hybridized in PCl3. It has 3 bond pairs and 1 lone pair. Geometry is trigonal pyramidal, similar to NH3. Bond angle is about 100°.

PCl3 + 3H2O → H3PO3 + 3HCl (hydrolysis gives phosphorous acid).

PCl5 (Phosphorus pentachloride)

P is sp3d hybridized in PCl5. It has 5 bond pairs and 0 lone pairs. Geometry is trigonal bipyramidal. There are two types of bonds:

• 3 equatorial P-Cl bonds: bond angle 120°. Shorter and stronger.
• 2 axial P-Cl bonds: bond angle 90° with equatorial, 180° with each other. Longer and weaker (more repulsion from equatorial bonds).

PCl5 + H2O → POCl3 + 2HCl (partial hydrolysis). PCl5 + 4H2O → H3PO4 + 5HCl (complete hydrolysis gives phosphoric acid).

In the solid state, PCl5 exists as [PCl4]+ [PCl6]- (ionic), not as covalent PCl5.

Oxoacids of Phosphorus

Key rule for basicity of P oxoacids: only the OH groups directly bonded to P are acidic. Any P-H bond is NOT acidic (the hydrogen is not ionizable). So count the P-OH groups to find basicity.

Group 16: Oxygen Family

Group 16 elements are oxygen (O), sulphur (S), selenium (Se), tellurium (Te), and polonium (Po). The general valence shell configuration is ns² np⁴. They commonly show -2 oxidation state (in binary compounds). Heavier elements can expand their octets and show +4 and +6 oxidation states.

Why Oxygen is Anomalous

• No d orbitals: Oxygen (period 2) cannot expand its octet. It shows only -2 and exceptionally +2 (in OF2) oxidation states.
• High electronegativity: O is the second most electronegative element (after F). This is why water has such a high boiling point (H-bonding).
• Small size: O forms strong double bonds (O=O in O2, C=O, S=O) using 2p orbitals. Heavier S, Se, Te prefer single bonds with lone pairs.
• Allotropy: Oxygen has two allotropes: O2 (dioxygen, the common form) and O3 (ozone).

Ozone (O3)

Ozone is an allotrope of oxygen. Its structure has resonance: O-O=O (with a lone pair on the central O). The O-O-O bond angle is 116.8°. The central O is sp2 hybridized. It is a bent molecule.

Preparation: O2 + electrical discharge → O3 (by electrolytic method: in the ozoniser). O3 is formed in the upper atmosphere when UV radiation breaks O2: O2 + UV → 2O; O + O2 → O3.

Properties: Ozone is a powerful oxidising agent. It decomposes to give nascent oxygen: O3 → O2 + [O]. It bleaches by oxidation. It liberates iodine from KI: O3 + 2KI + H2O → 2KOH + I2 + O2 (test for ozone). It also bleaches indigo dye.

Ozone layer: The ozone layer (15-40 km altitude, stratosphere) absorbs harmful UV-B and UV-C radiation. CFCs (chlorofluorocarbons) and nitrogen oxides destroy ozone through catalytic cycles. One Cl atom can destroy over 100,000 ozone molecules.

Hydrogen Peroxide (H2O2)

H2O2 is a pale blue liquid in pure form. Structure: non-planar with O-O bond. The dihedral angle between the two O-H groups is about 111.5° in the liquid.

Preparation (industrial): autoxidation of 2-ethylanthraquinol.

Oxidising property: H2O2 + 2I- + 2H+ → I2 + 2H2O (oxidizes iodide in acidic medium). It acts as an oxidiser in acidic and basic media.

Reducing property: 2MnO4- + 5H2O2 + 6H+ → 2Mn2+ + 5O2 + 8H2O (H2O2 reduces permanganate in acidic medium). So H2O2 is amphoteric with respect to redox (acts as both oxidiser and reducer).

Bleaching: H2O2 bleaches hair and fabrics by oxidation (nascent O from decomposition).

Disproportionation: 2H2O2 → 2H2O + O2 (catalysed by MnO2, dust, or base). Store in dark bottles to prevent catalytic decomposition.

Sulphur: Allotropes and Oxoacids

Allotropes of Sulphur

Sulphur has two important crystalline allotropes:

Rhombic Sulphur (alpha-S)

• Stable form at room temperature (below 96°C).
• Yellow crystals with rhombic (orthorhombic) unit cell.
• Consists of crown-shaped S8 rings, puckered, with S-S-S angle 108°.
• Melting point: 114°C. Density: 2.06 g/cm3.
• Insoluble in water. Soluble in CS2.

Monoclinic Sulphur (beta-S)

• Stable above 96°C (transition temperature) up to 119°C.
• Needle-like monoclinic crystals, pale yellow.
• Also consists of S8 rings but different crystal packing.
• Melting point: 119°C. Density: 1.98 g/cm3.
• Converts back to rhombic S below 96°C.

Plastic Sulphur

• Formed by pouring boiling S into cold water.
• Amorphous, elastic (rubbery) solid. Long chains of S atoms.
• Metastable: slowly reverts to rhombic S on standing.

Sulphur Dioxide (SO2)

SO2 is produced when S burns in air: S + O2 → SO2. It is also produced in the roasting of sulfide ores: 2ZnS + 3O2 → 2ZnO + 2SO2.

Structure: SO2 is bent (angular), bond angle 119°. S is sp2 hybridized. There is a resonance structure with both S=O and S-O bonds.

Properties: Pungent smell. Acidic oxide (anhydride of H2SO3). A reducing agent: bleaches by reducing colored compounds, unlike Cl2 which bleaches by oxidation. The bleaching by SO2 is temporary (the colorless reduced compound can be re-oxidized).

Sulphuric Acid (H2SO4): Contact Process

H2SO4 is made industrially by the contact process:

1. S + O2 → SO2 (burning sulphur or roasting pyrite)
2. 2SO2 + O2 ⇌ 2SO3 (catalyst: V2O5, 450-500°C, 1-2 atm)
3. SO3 + H2SO4 → H2S2O7 (SO3 absorbed in 98% H2SO4, forms oleum)
4. H2S2O7 + H2O → 2H2SO4 (oleum diluted with water)

Note: SO3 cannot be dissolved directly in water because the reaction is too vigorous and produces a fine acidic mist. It is first absorbed in concentrated H2SO4 to form oleum.

Properties of Concentrated H2SO4

Concentrated H2SO4 (98%, density 1.84 g/mL) is a colourless, oily, dense liquid.

• Strong acid: In water it fully ionises. First dissociation is complete; second is partial (Ka2 = 0.012).
• Dehydrating agent: Removes water from organic compounds. Charcoal forms from sugar: C12H22O11 → 12C + 11H2O (conc. H2SO4 removes H2O from sucrose). Absorbs water of crystallization from CuSO4.5H2O to give white anhydrous CuSO4.
• Oxidising agent: Hot concentrated H2SO4 oxidises metals and non-metals. Cu + 2H2SO4(hot, conc.) → CuSO4 + SO2 + 2H2O. S + 2H2SO4(hot, conc.) → 3SO2 + 2H2O.
• Mixing with water is exothermic: Always add acid to water (not water to acid) to prevent localised boiling and splashing.

Oxoacids of Sulphur

Group 17: Halogens

Group 17 elements are fluorine (F), chlorine (Cl), bromine (Br), iodine (I), and astatine (At). The general valence configuration is ns² np⁵. They need one electron to complete their octet, which is why they are excellent oxidising agents and form -1 ions readily. All halogens are non-metals; astatine is radioactive.

Trends in Physical Properties

Bond Dissociation Energy Anomaly of F2

You might expect bond energy to decrease smoothly from F2 to I2 (as bond length increases). However, F2 (159 kJ/mol) has a LOWER bond dissociation energy than Cl2 (243 kJ/mol). This is a major NEET anomaly.

Reason: Fluorine atoms are very small. In the F2 molecule, the two F atoms are so close that their lone pairs (all three of them on each F) are also very close together. These lone pairs repel each other strongly, weakening the F-F bond. In Cl2, Br2, and I2, the atoms are larger so the lone pairs are further apart and the repulsion is less. This extra lone pair repulsion in F2 is why the F-F bond is weaker than expected.

Boiling Point Order of Hydrogen Halides

The boiling points are: HF (19.5°C) >> HI (-35.4°C) > HBr (-66.8°C) > HCl (-85.1°C).

Why HF is anomalous: Fluorine is the most electronegative element. The H-F bond is very polar, and HF molecules form strong intermolecular hydrogen bonds (F-H...F). These H-bonds require a lot of energy to break, giving HF a very high boiling point. HCl, HBr, and HI cannot form significant H-bonds. Their boiling points follow molecular mass order: HCl (lowest MW) to HI (highest MW) due to increasing van der Waals forces.

Acid Strength of HX

In water, the acid strength order is: HF << HCl < HBr < HI.

HF is a weak acid in water (pKa = 3.2), while HCl, HBr, and HI are strong acids. The H-F bond is so strong that it does not dissociate completely in water. Going from HCl to HI, the H-X bond gets weaker and the acid gets stronger.

Reducing power of HX: The reducing power order is HF < HCl <HBr < HI. HI is the strongest reducing agent because it has the weakest H-I bond. HF cannot reduce even MnO4- or MnO2, but HI can.

Reactions of Halogens

Displacement reactions: A more reactive halogen displaces a less reactive one. F2 displaces Cl-, Br-, I-. Cl2 displaces Br- and I-. Br2 displaces only I-.

Cl2 + 2KBr → 2KCl + Br2 Cl2 + 2KI → 2KCl + I2 Br2 + 2KI → 2KBr + I2

Important Compounds of Halogens

Bleaching Powder (Ca(OCl)Cl)

Prepared by passing Cl2 over dry slaked lime at 40°C: 2Ca(OH)2 + 2Cl2 → Ca(OCl)Cl + CaCl2 + 2H2O. Bleaching powder is actually Ca(OCl)Cl, not pure Ca(OCl)2. The active bleaching component is the hypochlorite ion (OCl-), which releases nascent O in acid medium.

Chlorine Water

Cl2 + H2O ⇌ HCl + HOCl. Chlorine water is acidic and a weak bleach due to hypochlorous acid (HOCl). HOCl releases [O] on exposure to light, which oxidises and bleaches coloured materials.

Sodium Hypochlorite (NaOCl)

NaOH + Cl2(cold) → NaCl + NaOCl + H2O. Used in household bleach (Javelle water). Also used for disinfecting swimming pools.

Iodine Extraction from Seaweed

Seaweed is burnt, and the ash (containing NaI and KI) is treated with Cl2: Cl2 + 2NaI → 2NaCl + I2. Iodine can be purified by sublimation.

Oxoacids of Halogens (Chlorine series)

Key point for NEET: As the number of oxygen atoms increases, acid strength increases (more electron withdrawal makes the O-H bond more polar, easier to donate H+). But oxidising power decreases as the OS of Cl increases (Cl is already in a high oxidation state, less tendency to be reduced further).

Interhalogen Compounds

Interhalogen compounds are formed between two different halogen atoms. The general formula is XYn where X is the larger (heavier, less electronegative) halogen, Y is the smaller (lighter, more electronegative) halogen, and n = 1, 3, 5, or 7.

Why They Form

Interhalogens are more reactive than the parent halogens because the X-Y bond is weaker than X-X or Y-Y bonds (it is a polar bond with less effective overlap). The heavier halogen has available d orbitals, allowing it to accommodate more than 4 bonds.

Types and Structures

Notable Points on Interhalogen Structures

• ClF3 and BrF3 (XY3): T-shaped. sp3d hybridization gives a trigonal bipyramidal electron geometry. The 2 lone pairs occupy the equatorial positions (less repulsion with bonds at 90°). The 3 F atoms go axial (1) and equatorial (2). Result: T-shape with bond angles close to 90° and 180°.
• BrF5 and ClF5 (XY5): Square pyramidal. sp3d2 hybridization gives an octahedral electron geometry. 1 lone pair occupies one position. The other 5 positions are bonded to F. Result: square pyramidal shape.
• IF7 (XY7): Pentagonal bipyramidal. sp3d3 hybridization. Only IF7 exists in this type because only iodine is large enough to accommodate 7 bonds with the small F atoms.

Properties of Interhalogen Compounds

• They are all reactive oxidising agents and halogenating agents.
• They react with water to give halide and oxyhalide acids.
• ICl (iodine monochloride) is used for the Wij's test (detecting C=C bonds in fats, Wij's solution).
• Many are used as fluorinating agents in industrial processes (e.g., ClF3 for uranium hexafluoride preparation in nuclear industry).

Group 18: Noble Gases

Group 18 elements are helium (He), neon (Ne), argon (Ar), krypton (Kr), xenon (Xe), and radon (Rn). They have completely filled valence shells: He is 1s², all others have ns² np⁶. This gives them exceptional stability and very low chemical reactivity, hence the older name "inert gases."

Occurrence and Discovery

Noble gases were discovered in the late 19th and early 20th century:

• Argon (Ar): Discovered by Rayleigh and Ramsay (1894). Most abundant noble gas (0.93% of air).
• Helium (He): First detected spectroscopically in the sun by Lockyer (1868) before being found on Earth.
• Neon, Krypton, Xenon: Found by Ramsay and Travers (1898).
• Radon: Discovered as a radioactive decay product.

Physical Properties

Noble Gas Compounds (Xenon Fluorides)

For a long time, noble gases were considered completely unreactive. In 1962, Neil Bartlett synthesised the first noble gas compound: Xe[PtF6]. This was followed by the synthesis of xenon fluorides. Among the noble gases, only xenon and krypton form well-characterised compounds (radon too, but it is radioactive).

Xenon Fluorides

Hydrolysis of Xenon Fluorides

XeF2 + H2O: 2XeF2 + 2H2O → 2Xe + 4HF + O2 (complete hydrolysis in excess water).

XeF4 + H2O (partial hydrolysis): 6XeF4 + 12H2O → 4Xe + 2XeO3 + 24HF + 3O2.

XeF6 + H2O: XeF6 + 3H2O → XeO3 + 6HF. Xenon trioxide XeO3 is formed. With excess NaOH: XeO3 + 2NaOH → Na2XeO4 + H2O (sodium perxenate).

Other Xenon Compounds

Why Are Noble Gas Compounds Rare?

Noble gases have fully filled valence shells (ns2 np6 for Ne through Rn), so they have no tendency to gain or lose electrons. They do not form ionic bonds easily. Their ionization energies are the highest in each period. However, the heavier noble gases (Kr, Xe) have larger atomic radii and their electrons are farther from the nucleus, making it easier to distort the electron cloud. Combined with very electronegative atoms like F and O (which can accept the electron density), Xe can form compounds. He and Ne have no known stable compounds because their IEs are too high.

Worked NEET Problems

Summary Cheat Sheet

Group 15: Key Points

Group 16: Key Points

Group 17 (Halogens): Key Points

Group 18 (Noble Gases): Key Points

Interhalogen Compounds: Quick Reference

Must-Know Reactions for NEET