NEET UG Inorganic Chemistry Syllabus

Inorganic Chemistry syllabus — full chapter and topic list

Expand any chapter to see its topics and sub-topics. This is the whole examinable outline for Inorganic Chemistry in NEET UG, not a summary of it.

1. 1

   UNIT 9: Classification of Elements and Periodicity in Properties

   3 topics

   • Modern periodic law and the current form of the periodic table
   • s, p, d, and f-block elements
   • Periodic trends in properties such as atomic and ionic radii, ionization enthalpy, electron gain enthalpy, valence, oxidation states, and chemical reactivity
2. 2

   UNIT 10: P-Block Elements

   5 topics

   • General Introduction
   • Electronic configuration
   • General trends in physical and chemical properties
   • Unique behavior of the first element of each group
   • Group 13 to Group 18 elements
3. 3

   UNIT 11: d- and f-Block Elements

   6 topics

   • General introduction
   • Electronic configurations
   • Occurrence and characteristics of transition elements
   • General trends in properties
     • Physical properties
     • Ionization enthalpy
     • Oxidation states
     • Atomic radii
     • Color
     • Catalytic behavior
     • Magnetic properties
     • Complex formation
   • Preparation, properties, and uses of compounds like K₂Cr₂O₇ and KMnO₄
   • Lanthanoids and actinoids
     • Electronic configurations
     • Oxidation states
     • Lanthanoid contraction
4. 4

   UNIT 12: Coordination Compounds

   8 topics

   • Introduction to coordination compounds
   • Werner’s theory
   • Ligands, coordination number, denticity, chelation
   • IUPAC nomenclature of mononuclear coordination compounds
   • Isomerism
   • Bonding: Valence bond approach and basic ideas of Crystal field theory
   • Color and magnetic properties
   • Importance of coordination compounds in qualitative analysis, extraction of metals, and in biological systems

Inorganic Chemistry flashcards for NEET UG

22 of 51 cards from the Inorganic Chemistry deck — real questions with worked answers.

1. What is an ionic (electrovalent) bond, and what type of elements typically form it?

   An electrostatic attraction between oppositely charged ions formed by complete transfer of electrons, typically between a metal (low ionization energy) and a non-metal (high electron affinity).

2. State Fajans' rules for the covalent character of an ionic bond.

   Covalent character increases with: small cation, large anion, high charge on either ion, and a cation with pseudo-noble-gas (18-electron, d10) configuration.

3. What does VSEPR theory predict, and what is the shape and bond angle of a molecule with 4 bonding pairs and no lone pairs (e.g., CH4)?

   VSEPR predicts molecular geometry from minimizing electron-pair repulsion. With 4 bond pairs and 0 lone pairs the shape is tetrahedral with a bond angle of 109.5 degrees.

4. According to VSEPR, in what order does repulsion between electron pairs decrease?

   Lone pair–lone pair > lone pair–bond pair > bond pair–bond pair.

5. Give the hybridization, shape, and bond angle of NH3 and explain why its angle differs from CH4.

   NH3 is sp3 hybridized, pyramidal, bond angle ~107 degrees. The lone pair on N exerts greater repulsion than bond pairs, compressing the H-N-H angle below 109.5 degrees.

6. What hybridization and geometry correspond to SF6, PCl5, and BF3?

   SF6: sp3d2, octahedral; PCl5: sp3d, trigonal bipyramidal; BF3: sp2, trigonal planar.

7. State the postulates of valence bond theory (VBT) for covalent bond formation.

   A covalent bond forms by overlap of half-filled atomic orbitals of two atoms with opposite spins; greater overlap gives a stronger bond; sigma bonds form by axial overlap and pi bonds by sideways overlap.

8. In molecular orbital theory, give the bond order formula and the bond order of O2.

   Bond order = (Nb - Na)/2, where Nb and Na are electrons in bonding and antibonding orbitals. O2 has bond order 2.

9. Why is O2 paramagnetic according to molecular orbital theory?

   O2 has two unpaired electrons in its degenerate pi* (antibonding) molecular orbitals, making it paramagnetic.

10. What is the MO energy-level ordering difference for molecules up to N2 versus O2 and F2?

    For Li2 to N2 the sigma2pz lies above pi2px=pi2py (due to s-p mixing); for O2 and F2 the sigma2pz lies below the pi orbitals.

11. List the main types of intermolecular forces in increasing order of strength.

    London dispersion (Van der Waals) < dipole–dipole < hydrogen bonding (dipole–dipole is generally weaker than H-bonding).

12. What is hydrogen bonding and what are its two types?

    An attractive force between an H atom bonded to a highly electronegative atom (F, O, N) and a lone pair on another electronegative atom. Types: intermolecular (e.g., water) and intramolecular (e.g., o-nitrophenol).

13. Why does water have an unusually high boiling point compared to H2S?

    Water forms strong intermolecular hydrogen bonds due to O's high electronegativity and small size, whereas H2S only has weak Van der Waals forces, so water boils much higher.

14. Why is hydrogen placed in Group 1 but also resembles Group 17 elements?

    Like Group 1, it has one valence electron and forms H+; like Group 17 (halogens) it is one electron short of a stable noble-gas configuration, can gain an electron to form H-, and is a diatomic gas.

15. What are the three isotopes of hydrogen and their compositions?

    Protium (1H, 1 proton, 0 neutron), Deuterium (2H or D, 1 proton, 1 neutron), Tritium (3H or T, 1 proton, 2 neutrons, radioactive).

16. What are interstitial (non-stoichiometric) hydrides, and which elements form them?

    Hydrides in which H atoms occupy interstitial spaces in a metal lattice, often non-stoichiometric. Formed by transition metals (d-block) and lanthanoids/actinoids; the d-block elements of groups 7, 8, 9 generally do not form them (hydride gap).

17. Define temporary and permanent hardness of water and how each is removed.

    Temporary hardness is due to bicarbonates of Ca and Mg, removed by boiling or Clark's method (adding Ca(OH)2). Permanent hardness is due to chlorides/sulphates of Ca and Mg, removed by washing soda (Na2CO3), Calgon, or ion-exchange resins.

18. Why are alkali metals (Group 1) the most reactive metals, and how does reactivity change down the group?

    They have a single, loosely held valence electron and low ionization energy, so they readily lose it. Reactivity increases down the group as ionization energy decreases.

19. What flame colours do Li, Na, K, Rb, and Cs impart?

    Li: crimson red; Na: golden yellow; K: lilac/violet; Rb: red-violet; Cs: blue/violet.

20. What is the diagonal relationship, and which Group 1 element shows it with which Group 2 element?

    A similarity in properties between an element and the one diagonally placed in the next group/period due to similar charge/radius ratio. Lithium resembles magnesium.

21. 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.

22. 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.

See more Inorganic Chemistry flashcards →

Planning Inorganic Chemistry for NEET UG

Inorganic Chemistry (NEET UG) FAQ