On this page
Groups 15, 16 and 17 contain some of the most familiar substances in chemistry: nitrogen, ammonia, oxygen, sulphuric acid, chlorine. They also contain a large share of the "explain why" questions in inorganic chemistry. The good news is that most of those questions have the same answer. This post sets out the trends, explains the anomalies of the first element in each group, and works through oxoacid oxidation states, a reliable source of marks.
What does each exam ask for in the p-block?
The JEE (Main) 2026 syllabus and the NEET (UG) 2026 syllabus cover groups 13 to 18 under a "general introduction": electronic configuration, trends in physical and chemical properties across periods and down groups, and the unique behaviour of the first element in each group. The JEE (Advanced) 2026 brochure goes further. For group 15 it lists ammonia, nitric acid, phosphine, PCl₃, PCl₅, oxides of nitrogen and oxoacids of phosphorus; for group 16 ozone, SO₂, sulphuric acid and oxoacids of sulphur; for group 17 chlorine, HCl, interhalogen compounds, oxoacids of halogens and bleaching powder. The current NCERT Class 12 Part I contents no longer include a p-block unit, so plan where you will study it from. Check the current year's syllabus before you start.
What are the main trends in groups 15 to 17?
Valence configurations are ns²np³ (group 15), ns²np⁴ (group 16) and ns²np⁵ (group 17).
Electronegativity falls down each group and rises across. PubChem's periodic table data give 3.04 for nitrogen, 3.44 for oxygen and 3.98 for fluorine; for the halogens, 3.16 for chlorine, 2.96 for bromine and 2.66 for iodine.
Metallic character rises down each group. In group 15, nitrogen and phosphorus are non-metals and bismuth is a metal.
Hydrides get less stable and more reducing down the group. Bond angles shrink too. NIST's computational chemistry database lists experimental H–N–H angles of 106.7° for ammonia and H–P–H of 93.3° for phosphine. In ammonia the bonding pairs sit close to the small, electronegative nitrogen and repel each other strongly; in phosphine they are farther out and the angle closes up.
Halogens get weaker as oxidising agents down the group. The NCERT table of standard electrode potentials gives 2.87 V for F₂, 1.36 V for Cl₂, 1.09 V for Br₂ and 0.54 V for I₂.
Why are nitrogen, oxygen and fluorine different?
Three properties explain almost every anomaly: small size, high electronegativity and no d orbitals in the valence shell.
Nitrogen forms strong pπ–pπ multiple bonds, so it exists as N≡N. Phosphorus forms P₄ with single bonds instead. With no d orbitals, nitrogen cannot exceed a covalency of four, which is why NCl₅ does not exist but PCl₅ does.
Oxygen forms hydrogen bonds, which is why water is a liquid at room temperature while H₂S is a gas.
Fluorine shows only the −1 oxidation state, while chlorine, bromine and iodine also show +1, +3, +5 and +7.
The electron affinity surprise. PubChem lists the electron affinity of chlorine as 3.62 eV and of fluorine as 3.34 eV. Fluorine's 2p shell is so compact that an added electron is repelled by the electrons already there.
The common mistake: writing that fluorine has the highest electron affinity because it has the highest electronegativity. They measure different things.
Why is HF the weakest of the hydrogen halides?
In water, acid strength runs HF < HCl < HBr < HI. The H–X bond gets longer and weaker down the group, so it breaks more easily. HF also forms strong hydrogen bonds. Students who go by electronegativity alone predict the opposite order.
Worked example 1: oxoacids of phosphorus
Find the oxidation state of phosphorus and the basicity (number of replaceable H⁺) of H₃PO₂, H₃PO₃ and H₃PO₄.
Oxidation states, with H = +1 and O = −2:
H₃PO₂: 3 + x − 4 = 0, so x = +1.
H₃PO₃: 3 + x − 6 = 0, so x = +3.
H₃PO₄: 3 + x − 8 = 0, so x = +5.
Basicity depends on structure, not formula. Only hydrogens on oxygen (P–OH) are acidic. H₃PO₄ has three P–OH groups and is tribasic. H₃PO₃ has two P–OH and one P–H, so it is dibasic. H₃PO₂ has one P–OH and two P–H, so it is monobasic. The P–H bonds also make H₃PO₂ and H₃PO₃ good reducing agents.
The common mistake: counting every hydrogen as acidic. Draw the structure first.
Worked example 2: two traps in sulphur and chlorine
Peroxodisulphuric acid, H₂S₂O₈. Treating every oxygen as −2 gives 2 + 2x − 16 = 0, so x = +7. Sulphur cannot exceed +6, its group number. The molecule contains one O–O peroxide link, so two oxygens are −1: 2 + 2x − 12 − 2 = 0, giving x = +6.
Chlorine in hot alkali. Chlorine disproportionates: 3Cl₂ + 6NaOH → 5NaCl + NaClO₃ + 3H₂O. Check: six chlorine atoms on each side; one goes from 0 to +5 (loses 5 electrons) while five go from 0 to −1 (gain 5 in total). In cold, dilute alkali the products are NaCl and NaOCl instead. Our guide to balancing redox reactions covers the method.
A practice routine for groups 15 to 17
Write one page of trends with a reason for each, then a second page of first-element anomalies.
Draw the structures of the oxoacids of phosphorus, sulphur and chlorine and assign every oxidation state.
Predict shapes of ClF₃, BrF₅, XeF₄ and PCl₅ with VSEPR; our post on VSEPR and molecular shapes has the method.
If you are taking JEE Advanced, add the named compounds in the brochure and write one preparation reaction for each.
Quiz yourself from memory every three days.
If you do not know where to begin with a chapter this large, Fermi's AI tutor works out what you already know and builds a plan for the session, so you start with the gaps rather than page one.
Today, draw the three phosphorus oxoacids from memory and mark which hydrogens are acidic.
