NEET UG Inorganic Chemistry Syllabus

Why is lithium an anomalous alkali metal? Give two examples.

Due to its very small size and high polarizing power. Examples: Li forms mainly the monoxide (Li2O) not peroxide, LiNO3 decomposes to Li2O (others give nitrite), Li is the only alkali metal forming a nitride (Li3N), and many Li salts are covalent and soluble in organic solvents.

What are the products when Na2O2 (sodium peroxide) and KO2 (superoxide) react with water/CO2?

Na2O2 + 2H2O gives 2NaOH + H2O2; KO2 reacts with CO2 to release O2 (4KO2 + 2CO2 gives 2K2CO3 + 3O2), used in breathing apparatus.

Why are alkaline earth metals (Group 2) less reactive than alkali metals?

They have two valence electrons, higher ionization energy, smaller atomic radii, and higher nuclear charge than the corresponding Group 1 metals, so they lose electrons less readily.

How does the solubility of Group 2 hydroxides and sulphates change down the group?

Hydroxide solubility increases down the group (Mg(OH)2 least soluble, Ba(OH)2 most). Sulphate solubility decreases down the group (BeSO4/MgSO4 soluble, BaSO4 insoluble).

Why does beryllium behave anomalously and resemble aluminium?

Its very small size and high charge density give it high polarizing power, so Be forms covalent compounds, has a diagonal relationship with Al; both form amphoteric oxides and are passivated by nitric acid.

What is the composition of Plaster of Paris and how does it set?

Plaster of Paris is calcium sulphate hemihydrate, CaSO4.1/2H2O. On adding water it sets to a hard mass of gypsum, CaSO4.2H2O.

Why does boron differ from other Group 13 elements and resemble silicon?

Boron is a small, non-metallic element with no metallic character; it shows a diagonal relationship with silicon, forms covalent and electron-deficient compounds (e.g., BF3, boranes), and acid B(OH)3 resembles silica's acidic oxide behaviour.

Explain the inert pair effect and its effect on stability in Group 13/14.

The reluctance of the ns2 electrons to participate in bonding, increasing down the group. It makes the lower oxidation state more stable down the group (e.g., Tl+ more stable than Tl3+; Pb2+ more stable than Pb4+).

Why is CO2 a gas but SiO2 a high-melting solid?

CO2 is a discrete molecule with C=O double bonds (pπ-pπ) held by weak Van der Waals forces; SiO2 is a giant covalent (network) solid where each Si is singly bonded to four O, giving a 3D lattice.

What is the structure of diborane (B2H6) and what are 3-centre-2-electron bonds?

Diborane has two B and four terminal H in a plane and two bridging H atoms. The B-H-B bridges are 3-centre-2-electron (banana) bonds where two electrons hold three atoms, accounting for boron's electron deficiency.

Why is the +2 oxidation state of carbon (in CO) and +4 in CO2 contrasted with the trend in Group 14 lower oxidation states?

For carbon and silicon the +4 state is most stable; down the group the +2 state becomes increasingly stable due to the inert pair effect, so Pb2+ is more stable than Pb4+ and PbO2 is a strong oxidizing agent.

Why does nitrogen show a maximum covalency of 4 while phosphorus can reach 5 or 6?

Nitrogen has no d-orbitals in its valence shell (period 2), limiting covalency to 4; phosphorus has accessible 3d orbitals allowing expansion of its octet (e.g., PCl5, PF6-).

How does the stability of +5 oxidation state vary in Group 15 and why?

Stability of +5 decreases down the group (N, P stable in +5; Bi5+ rare) because of the inert pair effect; thus the +3 state becomes more stable down the group (Bi mainly +3).

What is the brown ring test for nitrate, and what causes the brown colour?

Adding FeSO4 and then concentrated H2SO4 (added carefully along the side) to a nitrate gives a brown ring at the junction. The brown ring is the complex [Fe(H2O)5NO]2+ (nitrosyl iron complex).

How does acidic/basic character of Group 15 trioxides change down the group?

They change from acidic to basic down the group: N2O3 and P4O6 are acidic, As4O6 is amphoteric, Sb2O3 is amphoteric, Bi2O3 is basic.

How does the boiling point of Group 16 hydrides vary (H2O, H2S, H2Se, H2Te)?

H2O has anomalously high b.p. due to hydrogen bonding; among the rest b.p. increases from H2S to H2Te as molecular size and Van der Waals forces increase.

What is the structure of ozone (O3) and is it diamagnetic or paramagnetic?

Ozone is angular (bent) with a bond angle of about 117 degrees and two equal O-O bonds due to resonance; it is diamagnetic.

Describe the Contact process for manufacturing H2SO4 (key step and catalyst).

SO2 is catalytically oxidized to SO3 over V2O5 (vanadium pentoxide) catalyst (2SO2 + O2 -> 2SO3), SO3 is absorbed in conc. H2SO4 to form oleum (H2S2O7), which is diluted to give H2SO4.

Why does fluorine have a lower electron affinity than chlorine despite being more electronegative?

Fluorine's very small size causes strong inter-electronic repulsion in its compact 2p subshell, so the incoming electron is less effectively attracted, making its electron gain enthalpy less negative than chlorine's.

How does the oxidizing power of halogens change down Group 17, and what is the displacement rule?

Oxidizing power decreases down the group (F2 > Cl2 > Br2 > I2). A halogen higher in the group displaces a halide of one lower (e.g., Cl2 displaces Br- and I-).

What are interhalogen compounds? Give the four general types with an example.

Compounds formed between two different halogens of type AX, AX3, AX5, AX7. Examples: ClF (AX), ClF3 (AX3), IF5 (AX5), IF7 (AX7).

Why were noble gases (Group 18) earlier called inert, and which one first formed a compound?

They have completely filled valence shells (ns2np6, He is 1s2), giving very high ionization energies and chemical inertness. Xenon was the first to form a true compound (Xe + PtF6 by Neil Bartlett, then XeF2, XeF4, XeF6).

Give the hybridization and shapes of XeF2, XeF4, and XeF6.

XeF2: sp3d, linear; XeF4: sp3d2, square planar; XeF6: sp3d3, distorted octahedral.

Why do transition (d-block) elements show variable oxidation states?

Their (n-1)d and ns electrons have similar energies, so a variable number of electrons can participate in bonding, giving multiple oxidation states (e.g., Mn from +2 to +7).

Why are most transition metal compounds coloured?

Due to d-d electronic transitions: in a ligand field the d-orbitals split, and absorption of visible light promotes electrons between split d-levels; the complementary colour is observed.

Why do transition metals and their compounds often act as good catalysts?

Because of variable oxidation states (allowing them to form intermediates and provide reaction surfaces) and the ability to form complexes, lowering activation energy.

What are lanthanoid contraction and one of its consequences?

The steady decrease in atomic and ionic size across the lanthanoid series due to poor shielding by 4f electrons. Consequence: Zr and Hf (and other 4d/5d pairs) have nearly identical sizes and very similar properties, making their separation difficult.

How do the common oxidation states of lanthanoids and actinoids differ?

Lanthanoids show mainly +3 (with occasional +2, +4); actinoids show a wider range of oxidation states (+3 to +6 or +7) because 5f, 6d, and 7s orbitals have comparable energies.

State Werner's coordination theory: primary vs secondary valency.

Primary valency is ionizable and corresponds to the oxidation state (satisfied by anions); secondary valency is non-ionizable and equals the coordination number (satisfied by ligands), giving the complex a definite geometry.

What is a chelate ligand and a denticity? Give an example of a hexadentate ligand.

A chelate is a ring formed when a polydentate ligand binds a metal through more than one donor atom; denticity is the number of donor atoms it uses. EDTA4- is a hexadentate ligand (binds through 4 O and 2 N donor atoms).

Using crystal field theory, what is the spectrochemical series and how does it relate to high-spin vs low-spin complexes?

The spectrochemical series ranks ligands by field strength (e.g., I- < Br- < Cl- < F- < OH- < H2O < NH3 < en < NO2- < CN- < CO). Strong-field ligands (large splitting, delta-o) give low-spin complexes; weak-field ligands give high-spin complexes.