Are you struggling to memorize the vast trends, anomalous behaviors, and chemical reactions of the p-block elements? You aren’t alone. In the Inorganic Chemistry section of competitive exams like JEE Main, JEE Advanced, and NEET, the p-block (Groups 13 to 18) carries significant weightage, often accounting for several direct questions every year.
To help you secure those crucial marks, we have compiled these p-block elements notes to simplify complex concepts into a “one-shot” revision format. Whether you are a Class 11 or Class 12 student, these notes cover everything from the Inert Pair Effect and Catenation to the detailed properties of the Nitrogen, Oxygen, and Halogen families.
What’s Included in These Notes?
- NCERT-Based Summary: Every line is aligned with the latest syllabus to ensure no important trend is missed.
- JEE & NEET Focus: Special emphasis on “Exceptions” and “Anomalous Properties”—the favorite topics of examiners.
- Visual Aids: Comparative tables for atomic radii, ionization enthalpy, and electronegativity across Groups 13-18.
- Quick Revision: Designed for “One Shot” learning, perfect for last-minute exam preparation.
Stop flipping through heavy textbooks. Dive into these structured notes and master the p-block elements with clarity and confidence.
P-Block Elements: Introduction
P-block elements include elements of groups 13, 14, 15, 16, 17, and 18. These elements are called p-block elements because their last electron is added to a p-orbital.
Group – 13
- This group is known as Boron family.
- The members if group 13 are Boron (B), Aluminium (Al), Gallium (Ga), Indium (In) and Thallium (Tl).
- General electronic configuration: ns2 np1.
- Order of atomic radii: B < Al > Ga < In < Tl.
In general, size is increasing on moving down the group due to the addition of an extra shell but Gallium is an exception. Gallium is smaller than Aluminium due to the poor shielding effects of d-electron in Gallium. - General Oxidation state: +3.
- Trend of ionization energy: From B to Al, it decreases due to an increase in size but then from Al to Ga, it decreases slightly due to poor shielding of d-electrons. From Ga to In, it decreases slightly and then again increases for Tl due to poor shielding of f-electrons.
- Trend of electronegativity: Decreases from B to Al and then increases from Al to Tl. This trend is again due to the poor shielding of d and f electrons.
- Metallic character: Boron has some chracter of metalloids. Aluminium is most metallic and then metallic character is also present Ga to Tl.
- Reactivity of Boron family with oxygen: These elements react with oxygen at higher temperatures to form trioxides M2O3.
4 Al + 3 O2 → 2 Al2O3
4 B + 3 O2 → 2 B2O3 - Reactivity with water: At high temperature only, boron reacts with steam to form boron oxide.
2 B + 3 H2O → B2O3 + H2 - Aluminium reacts with cold water and liberates hydrogen gas. Gallium (Ga) and Indium (In) do not react with water. Thallium (Tl) form TlOH in the moist air.
4 Tl + 2 H2O + O2 → 4 TlOH - Reactivity towards metals: Only boron reacts with the metal to form borides (M3B2).
3 Mg + 2 B → Mg3B2 - Reaction with acid and alkali: Boron reacts with hot and concentrated nitric acid to form boric acid and nitrogen dioxide.
B (s) + 3 HNO3 (aq) → H3BO3 (aq) + 3 NO2 (g) - Other elements of this group react with acids to produce hydrogen gas.
- At temperatures above 773K, boron reacts with alkalis to form borates.
2 B (s) + 6 KOH (aq) → 2 K3BO3 (aq) + 3 H2 (g)
Important Compounds of Boron
Orthoboric Acid
- It is a weak acid and also a monobasic acid of boron.
- Formula: H3BO3 or B(OH)3
- Geometry: Trigonal Planar
- Structure:
Preparation
- From borax: By heating a concentrated solution of borax with sulphuric acid or hydrochloric acid.
Na2B4O7.10H2O + 2 HCl → 4 H3BO3 + 2 NaCl + 5 H2O - By hydrolysis of diborane: B2H6 + 6 H2O → 2 H3BO3 + 6 H2
- By hydrolysis of borane trihalides: BX3 + 3 H2O → H3BO3 + 3 HX
Properties
- Action of heat:
- Reaction with Metal Oxide:
where M stands for a bivalent metal
- Reaction with Ammonium borofluoride:
Borax
- Also known as Sodium Tetraborate.
- Formula: Na2B4O7.10H2O
Preparation
- From Boric acid: 4 H3BO3 + Na2CO3 → Na2B4O7 + 6 H2O + CO2
Properties
- Basic Nature: The aqueous solution of borax is alkaline in nature.
Na2B4O7 + 3 H2O → NaBO2 + 3 H3BO3
NaBO2 + 2 H2O → NaOH + H3BO3
- Action of Heat:
Diborane
- Formula: B2H6 or (BH3)2
- Structure: 2 B-H-B bonds are called as Banana bonds. Diborane contains 2 3c-2e bonds and 4 2c-2e bonds.
Preparation
- Reduction of Boron trifluoride: BF3 + 3 LiAlH4 → 2 B2H6 + 3 LiAlF4
- From NaBH4: 2 NaBH4 + H2SO4 → B2H6 + 2 H2 + Na2SO4
2 NaBH4 + H3PO4 → B2H6 + 2 H2 + NaH2PO4
Properties
- Reaction with water: B2H6 + H2O → 2 H3BO3 + 6 H2
- Combustion: B2H6 + 3 O2 → B2O3 + 3 H2O
Important Compounds of Aluminium
Aluminium Oxide
- Also known as alumina.
- Formula: Al2O3.
- It is amphoteric in nature.
Preparation
- Aluminium oxide is produced by heating aluminium hydroxide or aluminium sulphate.
2 Al(OH)3 + Heat → Al2O3 + 3 H2O
Aluminium Chloride
- Formula: AlCl3
- It exists as dimer.
- Structure:
Properties
- White, hygroscopic solid. Thus, it absorbs moisture from air.
- There are weaker intermolecular forces due to which it sublimes at 183 oC.
- Hydrolysis: AlCl3 + 3H2O → Al(OH)3 + 3 HCl + 3 H2O
- Action of Heat: 2 AlCl3.6H2O → 2 Al(OH)3 + Al2O3 + 6 HCl + 3 H2O
Group – 14
- This group is known as Carbon family.
- The members of group 14 are Carbon (C), Silicon (Si), Germanium (Ge), Tin (Sn) and Lead (Pb).
- General elctronic configuration: ns2 np2.
- Valency: 4
- General Oxidation state: +4 and +2 (inert pair effect).
- Carbon shows different properties as compared to the rest of the elements because of its small size and its catenation property.
- Size of the atoms increases on going down the group but after Si, there is only a slight increase in the size due to the poor shielding effect of d-electrons.
- Trend in ionization energy: C>Si>Ge>Pb> Sn.
Group 14 ionization energy generally decreases down the group (C to Pb) due to increasing atomic size and enhanced shielding, though with a notable exception where lead (
- The size of Pb and Sn are comparable, so due to higher charge density, itis more difficult to ionize lead.
- On going down the group, the metallic character increases. Carbon and Silicon are non-metals, Germanium is a metalloid, and Tin and Lead are metals.
- Catenation is the property of carbon atoms to bond chains with other carbon atoms.
Important Compounds of Carbon
Carbon Monoxide
- Formula: CO
- CO is a very poisonous gas due to the formation of carboxy haemoglobin in the blood.
Preparation
- On heating the mixture of powdered zinc and calcium carbonate (Lab method): Zn + CaCO3 → ZnO + CaO + CO
- By dehydrating methanoic acid in the presence of sulphuric acid: HCOOH + H2SO4 → CO + H2O
- In industries, carbon monoxide is generated by passing air over red hot coke: 2 C + O2 → 2 CO + Heat
- In industries, it is also produced by the reaction of steam and carbon: 2 C + H2O + Heat → 2 CO + H2
Properties
- CO behaves as strong reducing agent and reduces metal oxides.
Fe2O3 + 3 CO → 2 Fe + 3 CO2
CuO + CO → Cu + CO2 - CO reacts with nickel and forms tetracarbonyl nickel.
Ni + 4 CO → Ni(CO)4 - CO reacts with water vapours at high temperature to produce carbon dioxide.
CO + H2O → CO2 + H2
Carbon dioxide
- Formula: CO2
- It is a greenhouse gas, and thus causes global warming.
- Geometry: Linear
- Structure:
Preparation
- By action of acids on carbonates: CaCO3 + 2 HCl → CaCl2 + H2O + CO2
- By combustion of carbon: C + O2 → CO2
Properties
- When carbon dioxide is passed to lime water, it turns milky due to the formation of calcium carbonate, and the milkiness disappeared in the presence of excess of CO2
Ca(OH)2 + CO2 → CaCO3 + H2O
CaCO3 + CO2 → Ca(HCO3)2 - It behaves as acid and reacts with base to form carbonates and bicarbonates.
CO2 + NaOH → NaHCO3
NaHCO3 + NaOH → Na2CO3 + H2O
Important Compounds of Silicon
Sodium Silicate
- Structure of the silicate is SiO44-, in which four oxygen atoms are bonded to one silicon atom. Assembling the silicate unit forms the ring, chain, and 3D structure.
- Glass and cement are two important silicates made by humans.
- Hydrolysis of alkyl or aryl substituted chlorosilane forms silicon polymer with Si-O-Si bonds.
Preparation
- By melting soda ash in pure sand at high temperature: Na2CO3 + SiO3 → Na2SiO3 + CO2
Group – 15
- This group is known as Nitrogen family.
- Also known as Pnicogens.
- The members of group 15 are nitrogen (N), phosphorus (P), arsenic (As), antimony (Sb) and bismuth (Bi).
- General electronic configuration: ns2 np3
- General oxidation states: From -3 to +5, but +3 is most stable.
- Trend in ionisation energy: N > P > As > Sb > Bi.
- Size of atoms increases on going down the group.
- Down the group, metallic character increases.
- Nitrogen is a diatomic gas, while others are solid in nature.
- Except Bi, all other elements show allotropy.
- Phosphorous is present in many allotropic structures. Two of these important allotrope structures are red phosphorous and white phosphorous.
- Arsenic exists in three important allotrope structures, i.e., black, yellow and grey.
- Antimony has three allotropes, i.e., yellow, explosive and metallic.
Important Compunds of Nitrogen
Dinitrogen
- Formula: N2
- Most abundant chemical species in air.
Preparation
- Laboratory method: In the laboratory N2 is produced by heating an aqueous solution of ammonium chloride and sodium nitrite.
NH4Cl (aq) + NaNO2 (aq) → NaCl (aq) + 2 H2O (l) + N2 (g) - By heating red crystals of ammonium dichromate: (NH4)2Cr2O7 + Heat → N2 + 4 H2O + Cr2O3
- Oxidation of ammonia: Nitrogen is produced when ammonia is oxidised by red hot copper oxide or chlorine.
2 NH3 + 3 CuO → N2 + 3 H2O + 3 Cu
8 NH3 + 3 Cl2 → N2 + 6 NH4Cl - Very pure nitrogen can be obtained by heating sodium azide.
NaN3 → 2 Na + 3 N2
Properties
- N2 binds to some strong electropositive metals at high temperatures to form their nitrides.
6 Li + N2 → 2 Li3N2
3 Mg + N2 → Mg3N2
2 Al + N2 → 2 AlN - N2 combines with O2 at temperature above 3273 K to form nitric oxide.
N2 + O2 → 2 NO
Oxides of Nitrogen
Oxyacids of Nitrogen
Ammonia
- Chemical Formula: NH3
- Shape: Pyramidal
- Structure:
Preparation
- Ammonium salt is heated with a strong alkali: NH4Cl + NaOH → NaCl + H2O
- By hydrolysis of magnesium nitride: Mg3N2 + 6 H2O → 3 Mg(OH)2 + 2 NH3
- Haber’s Process: N2 (g) + 3 H2 (g) → 2 NH3 (g)
Properties
- Basic nature: Ammonia is basic in nature.
- Reaction with halogen:
8 NH3 + 3 Cl2 → 6 NH4Cl + N2
NH3 + 3 Cl2 (excess) → NCl3 + 3 HCl
8 NH3 + 3 Br2 → 6 NH4Br + N2
NH3 + 3 Br2 (excess) → NBr3 + 3 HBr
2 NH3 + 3 I2 → NH3.NI3 + 3 HI
8 NH3.NI3 → 6 NH4I + 9 I2 + 6 N2 - Complex formation:
Ag+ + NH3 → [Ag(NH3)]2+
Cu2+ + 4 NH3 → [Cu(NH3)4]2+
Cd2+ + 4 NH3 → [Cd(NH3)4]2+ - Precipitation of heavy metal ions from the aqueous solution of their salts:
FeCl3 + 3NH4OH → Fe(OH)3 (Brown ppt) + 3 NH4Cl
AlCl3 + 3 NH4OH → Al(OH)3 (White ppt) + 3 NH4Cl
CrCl3 + 3NH4OH → Cr(OH)3 (Green ppt) + 3 NH4Cl
Group – 16
- This group is known as Oxygen family.
- Also known as Chalcogens.
- The members of group 16 are oxyegn (O), sulphur (S), selenium (Se), tellurium (Te), polonium (Po).
- General electronic configuration: ns2 np4.
- General oxidation state: -2 to +6, but -2 is common.
- Atomic and ionic radii: As the number of shells increases, the atomic and ionic radii increase from top to bottom within the group.
- Ionisation enthalpy: Due to expansion of atoms, the ionization enthalpy decreases within the group.
- Electron gain enthalpy: Due to compact nature of oxygen, the electron gain enthalpy is lower than that of sulphur. After sulphur, the electron gain enthalpy is reduced within the group.
- Electronegativity: It decreases within the group, i.e., the metallic properties increase from oxygen to plutonium.
- Reactivity to Hydrogen: All elements in the group form hydrides H2E (E = S, Se, Te, Po).
- Thermal stability: The thermal stability of the hydrides is H2O > H2S > H2Se > H2Te > H2Po.
- Acidity: The acidic character of hydrides increases down the group, i.e., H2O< H2S< H2Se< H2Te.
- Reducing character: The reducing character also decreases down the group due to the decreasing bond dissociation enthalpy, i.e., H2O < H2S < H2Se < H2Te < H2Po.
Group – 17
- This group is known as halogen family or halogens.
- Halogens are highly reactive non-metals.
- General electronic configuration: ns2 np5.
- These elements are very similar in their properties.
- The members of group 17 are fluorine (F), chlorine (Cl), bromine (Br), iodine (I) and astatine (As).
- Astatine is the only radioactive element in the group.
- All halogen group elements show an oxidation state of -1. However, elements such as chlorine, bromine, and iodine also exhibit +1, +3, +5, and+7 states.
- This higher oxidation state of chlorine, bromine, and iodine is realized when these halogens are combined with small and highly electronegative atoms, e.g., fluorine and oxygen.
- Fluorine, the most electronegative element, only has an oxidation state of-1.
- These elements have a higher ionisation enthalpy in their respective periods. This value continues to decrease as we move down the group.
- The electron gain enthalpy of these elements becomes less negative as you move down the group. Fluorine has a lower enthalpy than chlorine. We can attribute this to the small size and smaller 2p subshell of the fluorine atom.
- Halogens show high values of electronegativity. However, it slowly decreases as you move down the group from fluorine to iodine. This can be attributed to the increase in nuclear radii as one moves down the group.
Properties
- Physical state: Group 17 elements are found in various physical states. For example, fluorine and chlorine are gases. On the other hand, bromine is a liquid, and iodine is a solid.
- Colour: These elements have different colours. For example, while fluorine has a pale yellow colour, iodine has a deep purple colour.
- Solubility: Fluorine and chlorine are soluble in water. On the other hand, bromine and iodine are much less soluble in water.
- Melting and boiling points: These elements melting and boiling points increase as we move up the group from fluorine to iodine. Therefore, fluorine has the lowest boiling and melting points.
- Oxidizing power: All halogens are excellent oxidizing agents. Of the list, fluorine is the most effective oxidizing agent. It is able to oxidize all halide particles to halogen. The oxidizing power decreases as we move down the group. Halide particles also act as reducing agents. However, their reduction capacity also decreases in the group.
- Reaction with hydrogen: All halogens react with hydrogen to produce acidic hydrogen halides. The acidity of these hydrogen halides increases from HF to HI. Fluorine reacts violently and chlorine requires sunlight. On the other hand, bromine reacts when heated and iodine needs a catalyst.
- Reaction with oxygen: Halogens react with oxygen to form oxides. However, it was found that the oxides are not permanent. In addition to oxides, halogens also form a number of halogen oxoacids and oxoanions.
- Reaction with metals: Because halogens are very reactive, they react with most metals immediately to form the resulting metal halides. For example, sodium reacts with chlorine gas to form sodium chloride. Metal halides are ionic in nature. This is due to the highly electronegative nature of halogens and the high electropositive nature of metals. This ionic character of halide is reduced from fluorine to iodine.
Group – 18
- This group is known as rare gases or noble gases.
- The members of group 18 are helium (He), neon (Ne), argon (Ar), krypton (Kr), xenon (Xe) and radon (Rn).
- These elements are chemically inert, i.e., they do not participate in any chemical reaction.
- General electronic configuration: ns2 np6. The exception is helium, it has 1s2 electronic configuration.
- General oxidation state: Zero, however xenon show +2, +4 and +6 in compounds with oxygen and fluorine.
- Electron gain enthalpy: Group 18 elements exhibit very stable electronic configurations. They do not tend to accept electrons.
- Ionisation enthalpy: They have a high ionization potential due to their complete electronic configurations. This value decreases as you move down the group due to the expansion of size.
Properties
- Due to their stable nature, we find these elements as monatomic gases in their free state.
- They are colourless, tasteless and odourless gases. Particles of these elements have weak Van der Waals forces. This power increases as you move down the group. This is due to the expansion of the polarization capacity of the molecules.
- They exhibit low melting and boiling points. We can attribute this to weak Van der Waals forces. Melting and boiling points increase as we move down the group.
- We can condense these elements at extremely low temperatures. As the size of atoms in a group increases, the ease of liquefaction also increases.
- In 1962, Neil Bartlett hypothesized that xenon should react with platinum hexafluoride. He was the first to create a compound of xenon, called xenon hexafluoro platinate (V). Later, many xenon compounds were integrated, including fluorides, oxyfluorides, and oxides.
- The ionisation enthalpy of helium, neon and argon are too high to form compounds.
- Krypton only forms krypton difluoride, because its ionisation enthalpy is slightly higher than that of xenon.
- Although, radon has a lower ionization enthalpy than xenon, it forms onlya few compounds, such as radon difluoride, and a few complexes because radon has no stable isotopes. In any case, xenon forms a particularly significant number of compounds, e.g., XeF2, XeF4, XeF6, XeO3, XeOF4, XeO2F2.
Conclusion
Inorganic chemistry often makes or breaks your final rank, and the p-block is its heaviest pillar. Review the group trends, drill the exceptions, and practice the key reactions outlined above. You have the exact material you need for JEE and NEET right here; now it’s just about putting in the reps. Keep revising, stay consistent with your mock tests, and best of luck with your preparation!
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