Group 13 to 18 elements (except helium), in which the last electron enters the p-orbitals, constitute the p-block. In this chapter will be study systematic group wise details of p-block elements.
Elements: B, Al, Ga, In, Tl.
Trends in Chemical Reactivity and Oxidation States
Most common oxidation states are +3 and +1. Stability of +1 oxidation state increases on going down the group from aluminium to Tl. Thus, in aqueous solution Tl+ is more stable.
Chemical Behaviour
- Metallic character increases down the group due to decreasing I.E. and easy loss of electrons. Boron having small size cannot easily lose its electron and hence is an exception to metallic character.
- Al, Ga, In and Tl having vacant d-orbitals form [M(OH)4]-, [M(OH)4(H2O)2]- and [M(H2O)6]3+ type of complexes with H2O.
- MH3 type of hydrides are formed by group 13 elements. These hydrides act as Lewis acids and form adducts with strong Lewis bases. LiAlH4 is a white crystalline solid obtained as follows:
4LiH + AlCl3Et2O→ LiAlH4 + 3LiCl
- All halides of group 13 elements are known except TlI3. The halides have halogen bridged dimeric structures. Boron trihalides exist as monomers only as boron is too small to co-ordinate with four large halide ions. In case of BF3, the energy required to break p–p bond is not released during bridge formation. The ionic character of trihalides increases on moving down the group due to polarization effects according to Fajan's rule.
- Reaction with alkalies: Boron dissolves to give borates with evolution of H2 gas.
2B + 6NaOH Fuse→ 2Na3BO3 + 3H2Al and Ga with conc. alkalies form metaaluminate and gallate respectively.2Al + 2NaOH + 6H2O → 2Na[Al(OH)4] + 3H2
Sodium metaaluminate
2Ga + 2NaOH + 6H2O → 2Na[Ga(OH)4] + 3H2Sodium gallate
- Oxides: All elements form oxides of formula M2O3, except Tl which forms Tl2O due to inert pair effect.
Boron shows anomalous behaviour due to its small size and high electronegativity. Chemistry of boron is discussed in detail.
BORON
Boron is the only non-metal in its group due to high I.E., it cannot form B3+ instead it forms three covalent bonds.
Properties of Boron
It exists in mainly two allotropic forms, i.e. amorphous dark brown powder and crystalline black hard solid.
Reaction with air:
Boron nitride
Reaction with metals:
Magnesium boride
Boron does not react with H2O.
It reacts with oxidizing agents like HNO3, H2SO4 etc.
Boron reduces CO2 to carbon and SiO2 to Si.
Compounds of Boron
Borax or sodium tetraborate decahydrate (Na2B4O7.10H2O)
Borax naturally occurs as tincal, which contains 50% borax and is obtained from mineral colemanite.
Borax
Its aqueous solution is alkaline due to hydrolysis.
Borax when heated with ethyl alcohol and conc. H2SO4 gives volatile vapours of triethyl borate which burns with green edged flame.
Triethyl borate
On heating, borax loses water of crystallization and swells up to form fluffy mass. On further heating, transparent glassy bead is formed.
Sodium metaborate Boric anhydride
Borax bead
Boric acid or orthoboric acid (H3BO3) or B(OH)3:
It can be obtained from borax or from colemanite as follows,
It behaves as a weak monobasic acid due to the following reaction in water.
With ethyl alcohol and conc. H2SO4, it gives triethyl borate, which burns with a green edged flame.
On heating, boric acid loses water to form B2O3 in three stages.
Illustration 1: Borax structure contains
(A) two BO4 groups and two BO3 groups
(B) four BO4 groups only
(C) four BO3 groups only
(D) three BO4 and one BO3 groups
Solution: (A). Depends on the charge on the anion.
Illustration 2: The order of increasing acidic strength of BF3, BCl3 and BBr3 is
(A) BF3 > BCl3 > BBr3
(B) BF3 < BCl3 < BBr3
(C) BBr3 < BF3 < BCl3
(D) BCl3 > BBr3 > BF3
Solution: (B). Moving down the group, size of halogen increases, hence extent of overlap decreases, so back donation from halogen to boron decreases.
Illustration 3: The bonds present in borazole are
(A) 12σ, 3π
(B) 9σ, 6π
(C) 6σ, 6π
(D) 9σ, 9π
Solution: (A). Borazole is B3N3H6.
Exercise 1
- Inorganic benzene is
(A) BN (B) BF4
(C) B2H6 (D) B3N3H6
- BCl3 does not exist as dimer but BH3 exists as dimer (B2H6) because
(A) chlorine is more electronegative than hydrogen
(B) there is p-p back bonding in BCl3 but BH3 does not contain such multiple bonding
(C) large sized chlorine atoms do not fit in between the smaller boron atoms whereas small sized hydrogen atoms get fitted in between boron atoms
(D) None of the above
COMPOUNDS OF ALUMINIUM
Aluminium chloride (AlCl3):
It can be prepared by passing dry Cl2 or HCl gas over heated Al or by heating a mixture of alumina and carbon in a current of dry chlorine.
It exists as dimer Al2Cl6 in which each Al atom is tetrahedrally surrounded by four Cl atoms.
Anhydrous AlCl3 fumes in moist air due to formation of HCl.
It behaves as a Lewis acid and forms adduct with donor molecules. For example,
Alums:
Alums are double salts of formula X2SO4.Y2(SO4)3. 24H2O where X is a monovalent cation and Y is a trivalent cation. For example, K2SO4. Al2(SO4)3. 24H2O is potash alum.
- All alums are Isomorphours.
- All alums form acidic solutions due to formation of H2SO4.
Illustration 4: Al2O3 formation involves large quantity of heat evolution, which makes its use in
(A) deoxidizer
(B) confectionary
(C) indoor photography
(D) thermite welding
Solution: (D).
GROUP 14 ELEMENTS
Elements: C, Si, Ge, Sn, Pb
These elements have valence shell electronic configuration ns2, np2. Common oxidation states observed for group 14 elements are +4 for silicon and +4 and +2 for, Ge, Sn and Pb. Stability of +2 oxidation state increases in the sequence, Ge < Sn < Pb due to inert pair effect.
Trends in Chemical Reactivity
Due to vacant d-orbitals, elements of the group, except carbon, form compounds having co-ordination numbers higher than four like (SiF5)-, (SiF6)2- and (PbCl6)2-.
As we go down the group, the size of atom increases and the interatomic bond strength decreases Thus, the order of catenation decreases as we go down the group.
Tetrahedral, covalent tetrahalides of the type MX4 are formed by Si, Ge, Si and Pb. Ge, Sn and Pb also form dihalides with increasing stability in the sequence, GeX2 < SnX2 < PbX2 due to increasing stability of divalent state going down the group (inert pair).
Oxides of composition MO2 are formed by Si, Ge, Sn and Pb. SiO2 (silica) is a three dimensional network solid of silica and oxygen atoms connected by single covalent bonds with each silicon atom bonded to four oxygen atoms in tetrahedral arrangement. Silica exists both in crystalline and amorphous forms. Monoxides of Sn and Pb are known. Red lead (Pb3O4) is a combination of Pb(II) and Pb(IV) oxides, i.e. 2PbO.PbO2.
Silicates
Silicates are formed by heating metal oxides or metal carbonates with sand, e.g.
Silicates have basic unit of SiO44- where each silicon atom is bonded with four oxide ions tetrahedrally.
Type of silicates:
| Type | Units | Example |
| 1. Orthosilicates | Single SiO44- unit. | ZrSiO4, Mg2SiO4 |
| 2. Pyrosilicates | Two units of SiO44- joined along a corner oxygen. Hence, they are formed of Si2O76- units. | Sc2Si2O7, Zn3(Si2O7)Zn(OH)2.H2O |
| 3. Cyclic silicates | General formula: (SiO32-)n or (SiO3)n2n- | Si3O96- in Ca3Si3O9, Si6O1812-, Be3Al2Si6O18 |
| 4. Chain silicates | (SiO3)n2n- and (Si4O11)n6n- | Li2SiO3, MgSiO3 |
| 5. Two dimensional sheet silicates | (Si2O5)n2n- | Mg(Si2O5)2Mg(OH)2, Al2(OH)4(Si2O5) |
| 6. Three dimensional sheet silicates | All 4 oxygen atoms shared with adjacent SiO44- tetrahedral. | Quartz, zeolites, etc. |
CARBON
It exists in two or more allotropes having different physical properties but identical chemical properties. Amorphous forms are coal, coke and charcoal etc., while crystalline forms are diamond, graphite and fullerenes.
Among group 14 elements, carbon forms pπ – pπ multiple bond easily with itself (in graphite). Hence, carbon shows pronounced ability to form pπ – pπ multiple bonds.
Compounds of Carbon
Oxides:
Carbon forms two stable oxides, i.e. carbon dioxide (CO2) and carbon monoxide (CO). Less stable oxides are C3O2, C5O2 and C12O9 known as suboxides.
Carbon monoxide:
Can be prepared by heating oxalic acid or from formic acid or potassium ferrocyanide as follows:
- H2C2O4H2SO4-H2O→ CO + CO2 + H2O
- HCOOH H2SO4373K→ CO(pure) + H2O
- K4[Fe(CN)6] + 6H2SO4 + 6H2O → 2K2SO4 + FeSO4 + 3(NH4)2SO4 + 6CO
Carbon monoxide reduces metal oxides to metals.
Due to presence of lone pairs of electrons, CO acts as a Lewis base and this is the reason for the formation of metal carbonyls.
Phosgene, a poisonous gas is formed when CO combines with Cl2 in the presence of sunlight
Carbon dioxide:
Carbon dioxide is prepared from calcium carbonate.
Glucose and fructose fermentation also produces CO2
CO2 turns lime water milky due to the formation of insoluble CaCO3 which dissolves by passing excess of CO2 due to soluble Ca(HCO3)2 formation.
In photosynthesis, CO2 is converted into glucose, starch and O2.
Glucose
CO2 is also used in urea formation under a pressure of 220 atm and 453 – 473 K temperature.
Urea
Carbon forms other compounds, namely,
- carbides, e.g. Be2C, Al4C3
- acetylides having (C ≡ C)2- ions, e.g. MgC2, CaC2
- allylide having C34- ions, e.g. Mg2C3
- halides, e,g, CF4, CCl4, CBr4, CFC
- sulphides, e.g. CS, CS2, C3S2
GROUP 15 ELEMENTS
Elements: N, P, As, Sb, Bi.
Nitrogen and phosphorus are non-metals. The metallic character increases down the group due to lower I.E. and larger size. Hence, bismuth shows metallic character. The group state electronic configuration is ns2, np3.
Trends in Chemical Reactivity
Most common oxidation states shown by group 15 elements are – 3, +3 and +5. The stability of highest oxidation state (+5) decreases down the group.
The covalent character goes on decreasing as we move down the group in the sequence, P > As > Sb > Bi. This is due to increasing size of atom which refers to Fajan's rules.
Sb and Bi are the heavier elements of the group and form M3+ cations due to decrease in ionization enthalpy.
In contrast to nitrogen, the size of phosphorus atom forms pπ bonding and forms both cyclic and open chain compounds. Hence, it shows catenation.
NITROGEN
Nitrogen can be obtained from ammonium nitrite, ammonium dichromate, ammonia, urea and barium azide in the laboratory as follows.
Unstable
NH4NO2 → N2 + 2H2O
Barium azide
Nitrogen combines with oxygen under high temperature conditions to form nitric oxide
It combines with H2 in presence of a catalyst at 200 atm and 400 – 500°C temperature to form ammonia.
With metals it forms nitrides.
Some non-metals also combine with nitrogen to form nitrides.
Compounds of Nitrogen
Ammonia:
It is prepared from ammonium salt, magnesium nitride and by Haber's process as follows.
In water it forms a basic solution.
Ammonia reacts with halogens in different stoichiometic ratios.
Explodes in dry state
Ammonia acts as a Lewis base in complex formation.
Oxides of nitrogen
Nitrogen forms six different oxides, with oxidation states ranging from +1 to +5. Various oxides with their Lewis formula are listed below:
Preparations of above oxides are as follows:
Chemical reactivity of various oxides can be summarized as follows:
(i) Nitrous oxide is neutral towards litmus and reacts with sulphur, phosphorus, metals and hydrogen as follows:
(ii) (a) Nitric oxide is a stable oxide and decomposes only when heated at 800°C.
(b) It dissolves in cold ferrous sulphate solution.
Hydrated nitrosyl
(Brown ring)
(c) It also acts as an oxidizing agent.
(iii) Nitrogen tetroxide is a reddish brown pungent smelling gas, which associates and dissociates with change in temperature.
NO2 + 2Cu → Cu2O + NO }
(iv) Nitrogen pentoxide (N2O5) is a strong oxidizing agent.
(a) It is decomposed by alkali metals.
(b) With aq. NaCl, the reaction proves that N2O5 exists as ionic nitronium nitrate (NO2+NO3-).
Oxides of nitrogen
Nitrogen forms six different oxides, with oxidation states ranging from +1 to +5. Various oxides with their Lewis formula are listed below:
| Formula | Oxidation no. of nitrogen | Lewis formula |
| N2O Nitrous oxide |
+1 | :N≡N→Ö: |
| NO Nitric oxides |
+2 | :N=Ö: |
| N2O3 Dinitrogen trioxide |
+3 | O \ N—N / O |
| NO2 Nitrogen dioxide |
+4 | N / \ O O |
| N2O4 Dinitrogen tetroxide |
+4 | O O \ / N—N / \ O O |
| N2O5 Dinitrogen pentoxide |
+5 | O ‖ O N O \ / N / \ O O |
Preparations of above oxides are as follows:
Chemical reactivity of various oxides can be summarized as follows:
(i) Nitrous oxide is neutral towards litmus and reacts with sulphur, phosphorus, metals and hydrogen as follows:
(ii) (a) Nitric oxide is a stable oxide and decomposes only when heated at 800°C.
(b) It dissolves in cold ferrous sulphate solution.
Hydrated nitrosyl
(Brown ring)
(c) It also acts as an oxidizing agent.
(iii) Nitrogen tetroxide is a reddish brown pungent smelling gas, which associates and dissociates with change in temperature.
↓
300 K
N2O4
Yellow liquid
↓
260 K N2O4
Colourless
solid
(a) It acts both as oxidizing and reducing agent.
NO2 + 2Cu → Cu2O + NO }
(iv) Nitrogen pentoxide (N2O5) is a strong oxidizing agent.
(a) It is decomposed by alkali metals.
(b) With aq. NaCl, the reaction proves that N2O5 exists as ionic nitronium nitrate (NO2+NO3-).
Oxyacids of nitrogen:
Out of the five oxyacids of nitrogen, nitric acid (HNO3) is most important.
It can be prepared in the lab by heating NaNO3 or KNO3.
It can also be manufactured by Oswald's process.
Reactions of nitric acid are summarized in the following table.
| Concentration | Metal | Main Product |
| Very dilute HNO3 |
Mg, Mn | H2 + M(NO3)x |
| Fe, Zn, Sn | NH4NO3 + M(NO3)x Metal nitrate |
|
| Dilute HNO3 |
Pb, Cu, Ag, Hg | NO + M(NO3)x |
| Fe, Zn | N2O + M(NO3)x | |
| Conc. HNO3 |
Zn, Fe, Pb, Cu, Ag | NO2 + M(NO3)x |
| Sn | NO2 + H2SO4 Metastannic acid |
Oxides of phosphorus:
Phosphorus pentoxide (P4O10) is prepared by burning white phosphorus in excess of air or oxygen.
It acts as a dehydrating agent.
Final product or reaction with water is as follows.
Metaphosphoric Pyrophosphoric Orthophosphoric
Phosphorous trioxide (P4O6) is prepared by burning white phosphorous in limited supply of air.
It reacts with cold as well as hot water.
It burns in chlorine to form oxy – chlorides.
Oxyacids of phosphorus are summarized below:
(HPO3)n → It exists in polymeric form.
Exercise 2
- Which is true among the following?
(A) PH3 is stronger base than NH3
(B) Bond angle in PH3 is more than that of NH3
(C) PH3 is stronger reducing agent than NH3
(D) Boiling point of PH3 is more than that of NH3
- Orthophosphoric acid on strong heating forms
(A) P4O10 (B) H4P2O7
(C) (HPO3)x (D) P2O5
GROUP 16 ELEMENTS
Elements: O, S, Se, Te, Po.
First four elements are called chalcogens. The valence shell electronic configuration is ns2, np4, which shows each element has two electrons short of the next noble gas configuration.
Trends in Chemical Reactivity and Oxidation States
- Important oxidation states of S, Se and Te are -2, +2, +4 and +6. Oxidation states +4 and +6 are more pronounced for sulphur and heavier chalcogens, e.g. SF6, SeCl4, Te(OH)6.
- As we go down the group, the electronegativity for S, Se, Te and Po decreases due to increase in size.
- Due to easy loss of electron with increasing size, the metallic character increases on descending down the group. Hence, O, S are non-metals; Se, Te are metalloids whereas Po shows metallic character.
- Tendency for catenation decreases with increasing size as we move down the group.
- Thermal stability of hydrides decreases in the order: H2O > H2S > H2Se > H2Te > H2Po.
- Tendency to form multiple bonds decreases as we go down the group. Thus, S = C = S is quite stable as compared to Se = C = Se, which is unstable and decomposes readily while Te = C = Te is unknown.
OXYGEN
Oxygen can be prepared from various oxides and salts as follows:
Ozone:
In the laboratory ozone can be prepared by passing oxygen through a strong electric field.
Being a metastable allotrope, it always has a tendency to convert back to oxygen.
It acts as a powerful oxidizing agent.
The amount of ozone in a gas mixture can be determined by passing it through KI solution (at constant pH 9.2). The iodine liberated is titrated with sodium thiosulphate (hypo) solution.
O3 layer in the upper atmosphere is destroyed by exhaust gases containing nitrogen oxides. Halogens also damage O3 layer.
Also,
SULPHUR
Sulphur can be extracted from underground deposits by the Frasch process. It can also be prepared from H2S gas as follows:
H2S + 3⁄2O2 → SO2 + H2O
2H2S(g) + SO2(g) — Fe2O3 (catalyst), 673 K —→3⁄8S8(g) + 2H2O(g)
Allotropes of Sulphur
Sulphur exists in three allotropic forms:
- Rhombic sulphur (octahedral sulphur) which exists as S8 molecules. It is a bright yellow solid, soluble in CS2.
- Monoclinic sulphur or β-sulphur is dull yellow and soluble in CS2. Below 369 K it slowly changes to rhombic sulphur. At 369 K (the transition temperature) both forms can coexist.
- Plastic sulphur (amorphous sulphur or γ-sulphur) is a rubber-like, transparent yellow material, insoluble in CS2 and H2O. It is regarded as a super-cooled liquid that exists in long, random, inverted chains of sulphur atoms.
Compounds of Sulphur
Sulphur Dioxide
Among the several oxides of sulphur, SO2 (sulphur dioxide) and SO3 (sulphur trioxide) are the most important.
Preparation
- S8 + 8O2 → 8SO2
- 4FeS2 + 11O2 → 2Fe2O3 + 8SO2
SO2 shows both oxidizing and reducing properties.
Reactions illustrating its dual behaviour
- 2H2S + SO2 → 2H2O + 3S ↓ (oxidizing property of SO2)
- SO2 + 2Mg → 2MgO + S ↓ (oxidizing property of SO2)
- 2KMnO4 + 5SO2 + 2H2O → K2SO4 + 2MnSO4 + 2H2SO4 (SO2 acts as a reducing agent)
- 2FeCl3 + SO2 + 2H2O → H2SO4 + 2FeCl2 + 2HCl (SO2 acts as a reducing agent)
Sulphur Trioxide, Sulphuric Acid & Sodium Thiosulphate
Sulphur trioxide
(SO3) is prepared by catalytic oxidation of sulphur dioxide.
2SO2(g) + O2(g) — V2O5 catalyst, ~723 K —→ 2SO3(g)
It exists in three allotropic forms: α-SO3, β-SO3, γ-SO3.
SO3 is an acidic oxide:
SO3 + H2O → H2SO4 + heat
It forms oleum with concentrated sulphuric acid:
H2SO4 + SO3 → H2S2O7 (oleum)
Sulphuric acid
Sulphur forms many oxyacids; one important acid is sulphuric acid (H2SO4), also called oil of vitriol. It is prepared by the contact process and the lead chamber process.
Contact process (outline)
2SO2 + O2— catalyst —→ 2SO3
H2SO4 + SO3 → H2S2O7 (Oleum)
H2S2O7 + H2O → 2H2SO4
Properties
Sulphuric acid is a low-volatile liquid, a strong acid with a strong affinity for water. It removes water from wet gases that do not react with the acid and dehydrates many organic compounds (e.g., carbohydrates), causing charring.
C12H22O11 + 11 H2SO4 → 12C + 11 H2SO4 + 11 H2O
Both metals and non-metals are oxidized by concentrated sulphuric acid; the acid itself is reduced to SO2.
C + 2H2SO4 → CO2 + 2SO2 + 2H2O
Cu + 2H2SO4 → CuSO4 + SO2 + 2H2O
Sodium thiosulphate
Preparation
Prepared by boiling aqueous solutions of metal sulphites with elemental sulphur.
Na2SO3 + 1⁄8S8— H2O, ~378 K —→ Na2S2O3
It can also be manufactured by Spring’s reaction, treating sodium sulphide and sodium sulphite with iodine:
Na2S + Na2SO3 + I2 → Na2S2O3 + 2NaI
Reaction with acids
Na2S2O3 + 2HCl → 2NaCl + H2O + S ↓ + SO2
Group 17 (Halogens) – Notes & MCQs
Illustrations (MCQs)
(B) it is not a dehydrating agent
(C) it reacts with moisture to give an acid which reacts with NH3 (base)
(D) it is basic while NH3 is acidic
GROUP 17 ELEMENTS (HALOGENS)
Elements: F, Cl, Br, I, At. Astatine is radioactive. The melting and boiling points of halogens increase down the group due to increasing atomic size and hence stronger van der Waals forces. Thus, fluorine and chlorine are pale-yellow and greenish-yellow gases, bromine is a deep reddish-brown liquid, and iodine is a lustrous greyish-black crystalline solid.
Bond energy (X–X) order: F–F < Cl–Cl > Br–Br > I–I. The smaller bond energy of F–F is due to strong lone-pair repulsions in F2. From Cl2 to I2, increasing size reduces effective overlap and the bond strength decreases.
Trends in Oxidation States and Chemical Reactivity
- Fluorine, the most electronegative element, shows oxidation state −1 in all compounds and easily accepts electrons. Halogens act as oxidizing agents and can oxidize halide ions of halogens below them in the group.
Halogens react with metals and non-metals to form halides. Fluorine forms compounds with all elements except He, Ar and Ne.
Halogens form metal halides, M–X, whose ionic character decreases in the order M–F > M–Cl > M–Br > M–I. With increasing halide ion size, polarizability increases, covalent character rises, and ionic character falls. For a metal with multiple oxidation states, the halide of the higher oxidation state is more covalent due to the smaller cation size. Thus, SnCl4, PbCl4 and SbCl5 are more covalent than SnCl2, PbCl2 and SbCl3, respectively.
Hydrogen halides (HF, HCl, HBr, HI) are hydrohalic acids. Only HF is a liquid due to H-bonding. In aqueous solution their acid strength increases in the order HF < HCl < HBr < HI, following the decreasing H–X bond strength.
Other Comparative Properties
- Dipole moment: HI < HBr < HCl < HF.
- Bond length: HF < HCl < HBr < HI.
- Bond strength: HI < HBr < HCl < HF.
- Thermal stability: HI < HBr < HCl < HF.
- Acid strength: HF < HCl < HBr < HI.
- Reducing power: HF < HCl < HBr < HI (HI strongest).
Oxyacids
Halogens form various oxyacids. Fluorine does not form oxyacids. Important oxyacids of chlorine are listed below.
Oxyacids of Chlorine
| Formula | Name | Oxidation no. of Cl |
| HClO | Hypochlorous acid | +1 |
| HClO2 | Chlorous acid | +3 |
| HClO3 | Chloric acid | +5 |
| HClO4 | Perchloric acid | +7 |
Acidic strength: Increases with oxidation number of the halogen: HClO4 > HClO3 > HClO2 > HClO.
Group 17 (Halogens) – Notes & MCQs
Illustrations (MCQs)
(B) it is not a dehydrating agent
(C) it reacts with moisture to give an acid which reacts with NH3 (base)
(D) it is basic while NH3 is acidic
GROUP 17 ELEMENTS (HALOGENS)
Elements: F, Cl, Br, I, At. Astatine is radioactive. The melting and boiling points of halogens increase down the group due to increasing atomic size and hence stronger van der Waals forces. Thus, fluorine and chlorine are pale-yellow and greenish-yellow gases, bromine is a deep reddish-brown liquid, and iodine is a lustrous greyish-black crystalline solid.
Bond energy (X–X) order: F–F < Cl–Cl > Br–Br > I–I. The smaller bond energy of F–F is due to strong lone-pair repulsions in F2. From Cl2 to I2, increasing size reduces effective overlap and the bond strength decreases.
Trends in Oxidation States and Chemical Reactivity
- Fluorine, the most electronegative element, shows oxidation state −1 in all compounds and easily accepts electrons. Halogens act as oxidizing agents and can oxidize halide ions of halogens below them in the group.
Halogens react with metals and non-metals to form halides. Fluorine forms compounds with all elements except He, Ar and Ne.
Halogens form metal halides, M–X, whose ionic character decreases in the order M–F > M–Cl > M–Br > M–I. With increasing halide ion size, polarizability increases, covalent character rises, and ionic character falls. For a metal with multiple oxidation states, the halide of the higher oxidation state is more covalent due to the smaller cation size. Thus, SnCl4, PbCl4 and SbCl5 are more covalent than SnCl2, PbCl2 and SbCl3, respectively.
Hydrogen halides (HF, HCl, HBr, HI) are hydrohalic acids. Only HF is a liquid due to H-bonding. In aqueous solution their acid strength increases in the order HF < HCl < HBr < HI, following the decreasing H–X bond strength.
Other Comparative Properties
- Dipole moment: HI < HBr < HCl < HF.
- Bond length: HF < HCl < HBr < HI.
- Bond strength: HI < HBr < HCl < HF.
- Thermal stability: HI < HBr < HCl < HF.
- Acid strength: HF < HCl < HBr < HI.
- Reducing power: HF < HCl < HBr < HI (HI strongest).
Oxyacids
Halogens form various oxyacids. Fluorine does not form oxyacids. Important oxyacids of chlorine are listed below.
Oxyacids of Chlorine
| Formula | Name | Oxidation no. of Cl |
| HClO | Hypochlorous acid | +1 |
| HClO2 | Chlorous acid | +3 |
| HClO3 | Chloric acid | +5 |
| HClO4 | Perchloric acid | +7 |
Acidic strength: Increases with oxidation number of the halogen: HClO4 > HClO3 > HClO2 > HClO.
Interhalogens & Noble Gases – Notes and MCQs
Reason: This is because the loss of H+ ion from each oxyacid forms the conjugate base as (oxo-anion). The greater the number of oxygen atoms in the ion, the greater is the dispersal of negative charge and hence the higher is the stability of the resulting ion. Hence, the ease of formation of the resulting ions increases accordingly.
Interhalogen Compounds
Halogens can also form interhalogen compounds by reacting with each other. General formula: ABn where n = 1, 3, 5, 7; A = heavier halogen and B = lighter halogen; e.g., ICl3, IF7, etc. All interhalogen compounds are covalent and are generally more reactive than the individual halogens.
The stability of interhalogen compounds increases with the size of the central atom. Shapes of some interhalogen compounds are given below:
| Interhalogen compound | Shape | Hybridization |
| Type AB (n = 1) — BrF, BrCl, IBr, IF | Linear | sp |
| Type AB3 (n = 3) — BrF3, ICl3, IF3 | T-shape | sp3d |
| Type AB5 (n = 5) — BrF5, ICl5, IF5 | Square pyramidal | sp3d2 |
| Type AB7 (n = 7) — IF7 | Pentagonal bipyramidal | sp3d3 |
Pseudo Halide Ions and Pseudo Halogen
Ions consisting of two or more atoms, of which at least one is nitrogen, and having properties similar to halide ions are called pseudohalide ions. Some of these pseudohalide ions can be oxidized to form covalent dimers comparable to halogens. Such dimers are called pseudohalogens; e.g., CN− is a pseudohalide ion and (CN)2 is called cyanogen. SCN− is a pseudohalide ion (thiocyanate) and (SCN)2 is thiocyanogen.
Illustrations (MCQs)
Illustration 8: Which radical can bring about the highest oxidation state of a transition metal?
(A) F− (B) Cl− (C) Br− (D) I−
Solution: (A). F− because it has the smallest size and highest electronegativity.
Illustration 9: Which forms an “acid salt” with base?
(A) HCl (B) HBr (C) HI (D) HF
Sol.: (D). HF exists as H2F2 and forms KHF2 (acid salt) and KF (normal salt).
Illustration 10: Among the fluorides given below which will further react with F2?
(A) NaF (B) CaF2 (C) SF6 (D) IF5
Solution: (D).
Illustration 11: The following acids have been arranged in the order of decreasing acidic strength. Identify the correct order.
(A) ClOH > BrOH > IOH (B) BrOH > ClOH > IOH (C) IOH > BrOH > ClOH (D) ClOH > IOH > BrOH
Solution: (A). Electronegativity of halogen is inversely proportional to basic strength and directly proportional to acidic strength.
Exercise 3
- When Cl2 water is added to an aqueous solution of potassium halide in the presence of CHCl3, a violet colour is obtained. On adding more Cl2 water, the violet colour disappears and a colourless solution is obtained. This test confirms the presence of which in the aqueous solution?
(A) Iodide (B) Bromide (C) Chloride (D) Iodide and Bromide - Which of the following species is linear?
(A) I3− (B) XeF2 (C) ICl2− (D) All of the above
GROUP 18 ELEMENTS
Elements: He, Ne, Ar, Kr, Xe, Rn. These are noble gases. Their valence shells are fully occupied (ns2 np6), so gain or loss of electrons is difficult and they are chemically least reactive. All noble gases are monoatomic.
The natural abundance of noble gases in dry air is around 1%, of which argon is the major component. The main commercial source of He is natural gas. Radon is obtained as the decay product of the Radium-226 isotope.
Isolations: Rare gases are isolated from air by Ramsay and Rayleigh’s method. Several xenon compounds with very electronegative elements like fluorine and oxygen have been synthesized, while no true compounds of He, Ne and Ar are known. Krypton forms compounds rarely (notably KrF2). This trend is explained by increasing atomic size down the group, decreasing ionization energy, and the availability of vacant d-orbitals.
Xenon Fluoride and Xenon Oxygen Compounds
XeF2, XeF4 and XeF6 are the main fluorides of xenon.
XeFx are stable only in Ni containers. They are strong fluorinating agents and are hydrolysed even by traces of water. Xenon fluorides react with fluoride-ion acceptors to form cationic species and with fluoride-ion donors to form fluoroanions.
Hydrolysis of xenon fluorides produces xenon oxygen compounds. XeO3 with aqueous alkali forms hydrogen xenate ion.
Structure of xenon compounds
| Compound | Structure | Hybridization |
| XeF2 | Linear | sp3d |
| XeF4 | Square planar | sp3d2 |
| XeF6 | Distorted octahedron | sp3d3 |
| XeO3 | Pyramidal | sp3 |
| XeOF4 | Square pyramidal | sp3d2 |
| XeO4 | Tetrahedral | sp3 |
| XeO2F2 | Distorted octahedron | sp3d |
Separation of Rare Gases
Rare gases are separated individually from their mixture by Dewar’s charcoal adsorption method. After removing O2, CO2 and N2 from dry air, the remaining mixture of rare gases is adsorbed by coconut-charcoal at about −100 °C in a Dewar flask.
Illustration 12 :
Geometry of XeOF4 molecule is
(A) square planar (B) square pyramidal
(C) triangular bipyramidal (D) distorted octahedron
Solution. (B). It involves sp3d2 hybridization with one lone pair.
Exercise 4
i) The case of liquefaction of noble gases decreases in the order
(A) He > Ne > Ar > Kr > Xe (B) Xe > Kr > Ar > Ne > He
(C) Kr > Xe > He > Ar > Me (D) Ar > Kr > Xe > He > Me
ii) The gaseous mixture used by divers for respiration is
(A) N2 + O2 mixture (B) He + O2 mixture
(C) Ar + O2 mixture (D) Ne + O2 mixture
ANSWER TO EXERCISE
Exercise 1. 1) D 2) C
Exercise 2. 1) C 2) D
Exercise 3. 1) A 2) D
Exercise 4. 1) B 2) B
SOLVED EXAMPLES
1. Alumina on heating with carbon in nitrogen atmosphere gives
(A) Al + CO (B) Al + CO2
(C) AlN + CO (D) Al + CO + N2
Sol. (C).
2. Thallium shows different oxidation states
(A) as it is a transition metal (B) due to inert pair effect
(C) because of its amphoteric character (D) due to its high reactivity
Sol. (B).
3. Aluminum vessels should not be washed with materials containing washing soda because
(A) washing soda is expensive
(B) washing soda is easily decomposed
(C) washing soda reacts with aluminum to form soluble aluminates
(D) washing soda reacts with aluminum to form insoluble aluminum oxide
Sol. (C).
4. Which of the following is pseudoalum?
(A) (NH4)2SO4·Fe2(SO4)3·24H2O (B) K2SO4·Al2(SO4)3·24H2O
(C) MnSO4·Al2(SO4)3·24H2O (D) None of the above
Sol. (C).
5. Which is used in high temperature thermometry?
(A) Na (B) Tl (C) Ga (D) Hg
Sol. (C).
6. The main factor responsible for weak acidic nature of B–F bonds in BF3 is
(A) large electronegativity of fluorine
(B) three centred two electron bonds in BF3
(C) pπ–dπ back bonding
(D) pπ–pπ back bonding
Sol. (D).
7. Which of the following is methanide?
(A) Be2C (B) Al4C3 (C) Mg2C3 (D) Both (A) and (C)
Sol. (D). Both Be2C and Al4C3 are called methanides because they react with water to give methane.
8. Which of the following nitrate will produce laughing gas on heating?
(A) NaNO3 (B) Ca(NO3)2 (C) Hg(NO3)2 (D) NH4NO3
Sol. (D).
9. Among the following oxides, least acidic is
(A) P4O6 (B) P4O10 (C) As4O6 (D) As4O10
Sol. (C). Trioxides are less acidic than pentoxides. Acidic character decreases down the group.
10. The equivalent weight of phosphoric acid (H3PO4) in the reaction — is
(A) 25 (B) 49 (C) 59 (D) 98
Sol. (D).
ASSIGNMENT PROBLEMS
1. Which of the following has the minimum heat of dissociation?
(A) (B) (C) (D)
2.The role of fluorspar (CaF2) which is added in small quantities in the electrolytic reduction of alumina dissolved in fused cryolite (Na3AlF6) is
(A) as a catalyst
(B) to make the fused mixture very conducting
(C) to increase the temperature of the melt
(D) to decrease the rate of oxidation of carbon at the anode
3. The dissociation of Al(OH)3 by a solution of NaOH results in the formation of
(A) (B) (C) (D)
4. B–H–B bridge in B2H6 is formed by the sharing of
(A) 2 electrons (B) 4 electrons (C) 1 electron (D) 3 electrons
5. Fluorine is more electronegative than both boron and phosphorus. What conclusion can be drawn from the fact that BF3 has no dipole moment but PF3 does?
(A) BF3 is not spherically symmetrical
(B) BF3 must be triangular planar
(C) BF3 must be linear
(D) Atomic radius of P is larger than atomic radius of B
6. Identify the reagent I and II respectively in the following sequence:
[sequence]
(A) Acid, Al (B) Acid, C (C) Acid, Fe (D) Acid, Mg
7. The structure of quartz, mica and asbestos have same common units of
(A) (SiO4)4− (B) (SiO3)2− (C) (SiO4)2− (D) SiO2
8. Antidote for CO poisoning is
(A) Carborundum (B) Pure CO2 (C) Carbogen (D) Carbonyl chloride
9. R3SiCl on hydrolysis forms
(A) R3SiOH (B) R3Si–O–SiR3 (C) R3Si=O (D) None of the above
10. The thermal stability order for group 14 halides is
(A) CX4 > SiX4 > GeX4 > SnX4
(B) SnX4 > GeX4 > SiX4 > CX4
(C) SiX4 > CX4 > GeX4 > SnX4
(D) None of the above
11. Molecular shapes of SF4, CF4 and XeF4 are
(A) same with 2, 0 and 1 lone pair of electrons respectively
(B) same with 1, 1 and 1 lone pair of electrons respectively
(C) different with 0, 1 and 2 lone pair of electrons respectively
(D) different with 1, 0 and 2 lone pair of electrons respectively
12. The number of moles of water required to completely hydrolyze one mole of P4O10 is
(A) two (B) three (C) four (D) six
13. Ammonium dichromate is used in some fireworks. The green coloured powder blown during explosion is
(A) CrO3 (B) Cr2O3 (C) Cr (D) CrO(O2)
14. Antichlor is a compound
(A) which absorbs chlorine
(B) which removes chlorine from a material
(C) which liberates Cl2 from bleaching powder
(D) which acts as a catalyst in the manufacture of chlorine
15. Which of the following is not a fertilizer?
(A) 4Ca(H2PO4)2
(B) 3Ca(H2PO4)2 + 7CaSO4 + 2HF
(C) Ca(NO3)2 + NH4NO3
(D) CaF2 + (NH4)3PO4
ANSWERS TO ASSIGNMENT PROBLEMS
1. D 2. B 3. B 4. A 5. B 6. D
7. A 8. C 9. B 10. A 11. D 12. D
13. B 14. B 15. B
Frequently Asked Questions
When preparing p-block for NEET, the key is smart selection and repeated revision rather than trying to memorise everything at once.
Step 1: Focus on NCERT
Almost all NEET questions are directly from NCERT chemistry. Read the p-block chapters word-for-word, underlining exceptions, preparation methods, and uses.
Step 2: Group Prioritisation
For NEET, Groups 15, 16, and 17 are high-yield. Questions often focus on preparation, properties, and uses of compounds like HNO₃, H₂SO₄, PCl₅, and halogen oxoacids.
Step 3: Diagram & Table Learning
NCERT tables (like anomalous properties, trends) are question goldmines. Learn them visually—draw and label trends.
Step 4: Previous Years’ Papers
Solve last 10 years of NEET chemistry questions for p-block. This reveals patterns—often direct NCERT lines are used.
Step 5: Revision Cycles
Do 3–4 revisions before the exam. In the last week, only read marked/highlighted parts.
This method ensures you hit all NEET-specific points without drowning in unnecessary detail.
Period 1 has only hydrogen and helium, filling only the 1s orbital. The p-orbital starts from period 2, so there are no p-block elements in the first period.
Because the p-orbital has 3 sub-orbitals, each holding 2 electrons—maximum 6 electrons. This creates six columns in the periodic table’s p-block.
The count comes from adding all elements in Groups 13–18 across periods 2–7. Since period 1 has no p-block, the total is 6 columns × number of periods containing them = 35.
- Variable Oxidation States due to available orbitals.
- Diversity in Physical States—metals, non-metals, metalloids.
- Allotropy—multiple forms of the same element.
- Wide Reactivity Range—from highly reactive halogens to inert noble gases.
Yes. Group 13 elements the boron group—are the first column of the p-block. They include both non-metals (B), metals (Al, Ga, In, Tl), and a radioactive element (Nh). Their outer electron configuration is ns²np¹.
P-block elements include:
- Group 13: B, Al, Ga, In, Tl, Nh
- Group 14: C, Si, Ge, Sn, Pb, Fl
- Group 15: N, P, As, Sb, Bi, Mc
- Group 16: O, S, Se, Te, Po, Lv
- Group 17: F, Cl, Br, I, At, Ts
- Group 18: He, Ne, Ar, Kr, Xe, Rn, Og
If you’re starting p-block in Class 12 for NEET:
- Read NCERT First: Don’t jump to notes or guides until you’ve read the NCERT text fully.
- Mark Important Points: Highlight uses, preparation methods, and exceptions directly in the book.
- Group Order: Study 15 → 16 → 17 → 18, because NEET questions are more frequent here than in 13 & 14 (covered in Class 11).
- Revise Regularly: Set a 15-minute daily slot for p-block review to keep it fresh.
- Use Flashcards: For formulas, structures, and names of compounds.
This method prevents last-minute cramming and ensures steady retention.
For JEE Main & Advanced, the p-block is important but requires a deeper approach than NEET because JEE includes conceptual and reaction-based questions.
High-Weightage Areas:
- Group 13 & 14: Boron compounds (borax, boric acid), aluminium properties, carbon allotropes, silicones, silicates.
- Group 15: Preparation & properties of ammonia, nitric acid, phosphorus allotropes.
- Group 16: Oxygen allotropes, sulfur compounds like SO₂, H₂SO₄.
- Group 17: Halogen preparation, interhalogen compounds, oxoacids.
- Group 18: Xenon compounds (XeF₂, XeF₄, XeO₃).
Approach:
- Understand trends and anomalies (why B is different from Al, why F is most reactive halogen).
- Practice reaction mechanisms for compound preparation.
- Apply conceptual reasoning for theoretical questions.
For JEE, problem-solving practice is as important as theory—apply knowledge to real chemical situations.
In chemistry, the p-block refers to a set of elements in the periodic table whose outermost electrons occupy the p-orbital of their electron configuration. The periodic table is divided into blocks—s, p, d, and f—based on the type of orbital that receives the last electron. The p-block elements are found in Groups 13 to 18 of the periodic table, on the right-hand side. These groups include a wide range of element types, such as metals, metalloids, and non-metals, as well as some of the most reactive and biologically significant elements.
The name “p-block” comes directly from the quantum mechanical description of atoms. In the electron configuration notation, the letter “p” refers to the principal orbital type. The p-orbital can hold a maximum of six electrons, meaning there are six columns of elements in this block. This is why periods (rows) that have p-block elements always have six members in this section.
One of the defining features of the p-block is its diversity. It contains essential life-supporting elements like oxygen and nitrogen, technologically important materials like aluminium and silicon, and reactive halogens like fluorine and chlorine. It also includes noble gases such as helium, neon, and argon, which are chemically inert under normal conditions. The p-block spans a wide range of oxidation states and chemical behaviours, making it central to understanding both inorganic and organic chemistry.
From an academic perspective, studying the p-block helps students grasp periodic trends such as electronegativity, ionization energy, and atomic radius, as these trends are strongly visible across these groups. For competitive exams like NEET and JEE, mastering p-block concepts is crucial due to the high weightage of questions related to its properties, reactions, and applications.
Studying the p-block can feel overwhelming because it covers many groups, each with different trends, exceptions, and compound types. However, with a structured approach, you can make it manageable—and even interesting.
Step 1: Learn Group-Wise
Break it down into Groups 13–18. Focus on one group at a time, understanding physical and chemical properties, oxidation states, and important compounds.
Step 2: Master Trends
Memorise periodic trends—electronegativity, ionisation energy, melting points, acidic/basic oxides—because these explain most reactions and exceptions.
Step 3: Use Mnemonics
For remembering element names, oxidation states, or anomalies, mnemonics are powerful. For example, for Group 15: “Naughty People Are Small Babies” (N, P, As, Sb, Bi).
Step 4: Link with Real-World Uses
Associate each element with something tangible—Aluminium in aircraft, Silicon in chips, Nitrogen in fertilisers.
Step 5: Active Recall & Revision
Write short notes and keep revisiting them. Attempt MCQs regularly to reinforce concepts.
The 'P' in p-block refers to the p-orbital—one of the four types of atomic orbitals (s, p, d, f) where electrons can be found. In the periodic table, elements are grouped into blocks depending on the type of orbital that receives the last electron in their electron configuration.
For p-block elements, the outermost electrons enter the p-subshell. The p-orbital can hold a maximum of six electrons, which explains why there are six columns of elements in this block (Groups 13–18). Each period starts with s-block elements, moves through d-block (in longer periods), and ends with p-block elements before reaching noble gases.
The letter ‘p’ itself comes from the word “principal”, as used in the old spectroscopic notation, indicating the type of angular momentum of the orbital (l = 1 for p-orbitals).
From a chemical perspective, the filling of the p-orbital influences bonding, reactivity, and oxidation states. For example, oxygen (2p⁴) is highly electronegative and forms covalent bonds, whereas aluminium (3p¹) behaves as a reactive metal.
So when we say "p-block", we’re essentially classifying elements by the quantum mechanical property of their electrons, which in turn determines their chemistry.
S-block Elements: Groups 1 and 2 — Alkali metals (Li, Na, K, Rb, Cs, Fr) and Alkaline earth metals (Be, Mg, Ca, Sr, Ba, Ra).
P-block Elements: Groups 13 to 18 — Includes boron group, carbon group, nitrogen group, oxygen group, halogens, and noble gases.
D-block Elements: Groups 3 to 12 — Transition metals such as Fe, Cu, Zn, Ag, Au, Pt, Ni, Cr, Mn.
These divisions are based on which orbital (s, p, d) is being filled with electrons.
Physical Properties:
- State: Can be gases (O₂, N₂), liquids (Br₂), or solids (C, S, I₂).
- Density & Hardness: Varies—carbon (diamond) is extremely hard, while noble gases are light and inert.
- Conductivity: Non-metals are poor conductors; metals like aluminium conduct electricity well.
- Allotropes: Many p-block elements exist in multiple forms—carbon (diamond, graphite), phosphorus, sulfur.
Chemical Properties:
- Show variable oxidation states due to availability of d-orbitals in heavier elements.
- Oxides range from acidic to basic depending on metallic character.
- Many form covalent compounds, but metals in the block also form ionic compounds.
- Exhibit reactivity trends—halogens are strong oxidising agents, oxygen is a vital oxidant, and nitrogen forms stable triple bonds.
Example of an Interhalogen:
An interhalogen compound is formed between two different halogens, e.g., ClF₃ (Chlorine Trifluoride). These are often more reactive than individual halogens and are used in industrial fluorination processes.
A comparative study of p-block elements focuses on trends across periods (left to right) and down groups (top to bottom). This helps explain why elements behave differently despite being in the same block.
Physical Properties:
- Metallic Character: Decreases across a period (Boron → Oxygen) and increases down a group (Nitrogen → Bismuth).
- Electronegativity: Increases across a period and decreases down a group. Fluorine is the most electronegative.
- Melting and Boiling Points: Vary widely—carbon has one of the highest melting points, while noble gases have extremely low boiling points.
- Density: Generally increases down the group due to increased atomic mass.
Chemical Properties:
- Acidic/Basic Nature of Oxides: Across a period, oxides change from basic (Al₂O₃) to amphoteric to acidic (SO₂). Down a group, basicity increases.
- Reactivity: Halogens are highly reactive non-metals, while noble gases are almost inert.
- Type of Bonding: Covalent bonding dominates in the upper period elements; metallic and ionic bonding become more common down the groups.
For example, in Group 15, nitrogen forms strong multiple bonds (N≡N), while bismuth prefers metallic bonding. Understanding these trends allows chemists to predict reactions—critical in designing materials, medicines, and industrial processes.
The chemical properties of p-block elements vary widely because the block contains metals, non-metals, and metalloids across Groups 13–18. However, these elements show predictable patterns based on their position in the periodic table and the number of valence electrons in the p-orbital.
1. Oxidation States:
P-block elements exhibit multiple oxidation states. For example, Group 13 elements typically show +3 and sometimes +1 oxidation states due to the inert pair effect, while Group 15 elements range from –3 (e.g., NH₃) to +5 (e.g., P₂O₅). Halogens generally have –1 oxidation states but can show positive states in interhalogen compounds.
2. Reactivity Trends:
- Non-metals like oxygen, nitrogen, and halogens are highly electronegative and form covalent compounds.
- Metals such as aluminium and lead form ionic or covalent bonds depending on the partner element.
- Metalloids like boron and silicon show mixed behaviour.
3. Nature of Compounds:
P-block elements form acids (HCl, HNO₃), bases (aluminates), and amphoteric oxides (Al₂O₃). Non-metal oxides are usually acidic, while metal oxides are basic.
4. Allotropes:
Many p-block elements exist in multiple allotropes—carbon (diamond, graphite), phosphorus (white, red, black), and sulfur (rhombic, monoclinic).
5. Industrial Importance:
They participate in a wide range of industrial reactions—chlorine in water treatment, nitrogen in fertilisers, silicon in electronics, etc.
Understanding these properties is key for predicting reactivity in organic, inorganic, and industrial chemistry, which is why p-block mastery is essential for exams like NEET and JEE.
Helium has the electron configuration 1s², which means its valence electrons are in the s-orbital. Logically, one might expect it to be placed in the s-block alongside hydrogen. However, in the modern periodic table, helium is placed in Group 18—the noble gases—within the p-block section. This arrangement is based on chemical properties and periodic trends, not just electron configuration.
Helium’s position in Group 18 is justified because it shares the same defining property as other noble gases: a completely filled outer shell, which makes it chemically inert under normal conditions. Even though its outer shell is an s-orbital, the “block” classification in the periodic table is often secondary to grouping by similar reactivity. Helium behaves more like neon and argon than hydrogen or lithium, so it is placed alongside them.
Another reason for helium’s p-block placement is consistency in the periodic table layout. If helium were placed above beryllium in Group 2 (s-block), it would disrupt the classification of groups by chemical behaviour. Instead, its noble gas characteristics, such as extremely low reactivity, low boiling point, and monatomic gaseous state, align perfectly with the rest of Group 18.
This decision reflects the periodic table’s dual purpose: to show electron configurations and chemical similarities. While helium’s configuration is unique, its placement in the p-block ensures it sits with elements that mirror its inert nature, making it easier to study periodic trends and group behaviours.
The p-block elements contain all elements in Groups 13, 14, 15, 16, 17, and 18 of the periodic table. These groups are home to an incredible variety of substances, ranging from light gases to heavy metals. Specifically:
- Group 13 (Boron Group): Boron (B), Aluminium (Al), Gallium (Ga), Indium (In), Thallium (Tl), and Nihonium (Nh).
These show a mixture of metallic and non-metallic characteristics, with boron as a metalloid and aluminium as a widely used lightweight metal. - Group 14 (Carbon Group): Carbon (C), Silicon (Si), Germanium (Ge), Tin (Sn), Lead (Pb), and Flerovium (Fl).
These elements form the backbone of organic chemistry (carbon) and modern electronics (silicon). - Group 15 (Nitrogen Group or Pnictogens): Nitrogen (N), Phosphorus (P), Arsenic (As), Antimony (Sb), Bismuth (Bi), and Moscovium (Mc).
Known for a range of oxidation states and vital biological roles (e.g., nitrogen in DNA and proteins). - Group 16 (Oxygen Group or Chalcogens): Oxygen (O), Sulfur (S), Selenium (Se), Tellurium (Te), Polonium (Po), and Livermorium (Lv).
Oxygen is essential for respiration, while sulfur plays a key role in amino acids. - Group 17 (Halogens): Fluorine (F), Chlorine (Cl), Bromine (Br), Iodine (I), Astatine (At), and Tennessine (Ts).
Extremely reactive non-metals, halogens form salts with metals (e.g., NaCl). - Group 18 (Noble Gases): Helium (He), Neon (Ne), Argon (Ar), Krypton (Kr), Xenon (Xe), and Radon (Rn).
They are chemically inert due to a complete octet in their outer shell.
This diversity makes the p-block unique. It contains gases, solids, metals, non-metals, and metalloids. It also includes elements vital for life (O, C, N, P) and industrial processes (Al, Si, Cl). Understanding the constituents of the p-block is crucial for predicting chemical reactivity, industrial applications, and environmental behaviour of elements.