CSIR NET Chemical Sciences 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 CSIR NET Chemical Sciences, not a summary of it.

1. 1

   Chemical Periodicity

   5 topics

   • Periodic Table and Periodic Properties
   • Shapes of Orbitals
   • Ionization Energy
   • Electron Affinity
   • Electronegativity
2. 2

   Structure and Bonding

   4 topics

   • VSEPR Theory
   • Molecular Orbital Theory
   • Hybridization
   • Metal-Ligand Bonding in Transition Metal Complexes
3. 3

   s-Block Elements

   3 topics

   • Alkali Metals
   • Alkaline Earth Metals
   • Compounds of s-Block Elements
4. 4

   p-Block Elements

   2 topics

   • Group 13 to Group 18 Elements
   • Oxides, Halides and Hydrides of p-Block Elements
5. 5

   d-Block Elements

   4 topics

   • General Characteristics
   • Complex Formation
   • Coloured Ions
   • Catalytic Properties
6. 6

   f-Block Elements

   3 topics

   • Lanthanides
   • Actinides
   • Separation of Lanthanides
7. 7

   Coordination Compounds

   4 topics

   • Nomenclature
   • Isomerism
   • Bonding in Coordination Compounds
   • Applications of Coordination Compounds
8. 8

   Organometallics

   2 topics

   • Synthesis and Reactivity
   • Applications in Catalysis
9. 9

   Bioinorganic Chemistry

   2 topics

   • Metalloproteins and Metalloenzymes
   • Role of Metal Ions in Biological Systems
10. 10

    Instrumental Methods of Analysis

    4 topics

    • UV-Visible Spectroscopy
    • IR Spectroscopy
    • NMR Spectroscopy
    • Mass Spectrometry
11. 11

    Analytical Chemistry

    3 topics

    • Volumetric Analysis
    • Gravimetric Analysis
    • Chromatographic Techniques

Inorganic Chemistry flashcards for CSIR NET Chemical Sciences

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

1. What is the modern periodic law, and on what property is the modern periodic table arranged?

   The modern periodic law states that the properties of elements are a periodic function of their atomic numbers. The table is arranged in order of increasing atomic number (number of protons), not atomic mass.

2. How do atomic radius and metallic character vary across a period and down a group?

   Across a period (left to right): atomic radius decreases, metallic character decreases. Down a group: atomic radius increases, metallic character increases.

3. What is the difference in shape between s, p, and d orbitals?

   s orbitals are spherical; p orbitals are dumbbell-shaped (two lobes along an axis); d orbitals are mostly double-dumbbell (cloverleaf) shaped, except d(z^2) which has a dumbbell with a central torus (ring).

4. How many radial and angular nodes does an orbital have in terms of n and l?

   Angular nodes = l; radial nodes = n - l - 1; total nodes = n - 1.

5. Define first ionization energy.

   The minimum energy required to remove the most loosely bound electron from one mole of isolated gaseous atoms to form one mole of gaseous cations (M(g) -> M+(g) + e-).

6. Why is the first ionization energy of nitrogen higher than that of oxygen?

   Nitrogen has a stable half-filled 2p^3 configuration, which is extra stable; removing an electron from oxygen's 2p^4 relieves electron-electron repulsion in the paired orbital, so oxygen loses an electron more easily, giving it a lower IE.

7. Define electron affinity and state its sign convention.

   Electron affinity is the energy change when an electron is added to a neutral gaseous atom to form a gaseous anion (X(g) + e- -> X-(g)). It is usually negative (energy released) for the first electron and positive for adding to an anion.

8. Why is the electron affinity of chlorine greater (more negative) than that of fluorine?

   Fluorine's small size causes high electron-electron repulsion in its compact 2p subshell, reducing the energy released. Chlorine's larger size accommodates the added electron with less repulsion, giving a more negative EA. Order: Cl > F > Br > I.

9. Define electronegativity and name two common scales.

   Electronegativity is the tendency of an atom in a molecule to attract the shared bonding electron pair toward itself. Common scales: Pauling scale and Mulliken scale (Mulliken EN = average of ionization energy and electron affinity).

10. State the basic postulate of VSEPR theory.

    Electron pairs (bonding and lone) around a central atom arrange themselves to minimize repulsion, adopting positions as far apart as possible. Repulsion order: lone pair-lone pair > lone pair-bond pair > bond pair-bond pair.

11. Give the molecular shapes for AX3E (3 bond pairs, 1 lone pair) and AX2E2 (2 bond pairs, 2 lone pairs).

    AX3E is trigonal pyramidal (e.g., NH3); AX2E2 is bent/angular (e.g., H2O). Both are based on a tetrahedral electron geometry.

12. In Molecular Orbital Theory, what is bond order and how is it calculated?

    Bond order = (number of electrons in bonding MOs - number in antibonding MOs) / 2. A higher bond order means a stronger, shorter bond; bond order of zero means the molecule does not exist.

13. Why is O2 paramagnetic according to MOT?

    In O2, the two highest-energy electrons occupy the two degenerate antibonding pi*(2p) orbitals singly with parallel spins (Hund's rule), giving two unpaired electrons, which makes O2 paramagnetic. Bond order = 2.

14. What hybridization and geometry correspond to sp, sp2, sp3, sp3d, and sp3d2?

    sp - linear (180°); sp2 - trigonal planar (120°); sp3 - tetrahedral (109.5°); sp3d - trigonal bipyramidal; sp3d2 - octahedral (90°).

15. In crystal field theory, distinguish strong-field and weak-field ligands and their effect on spin.

    Strong-field ligands cause large d-orbital splitting (large Δ), favoring low-spin (electrons pair up) complexes. Weak-field ligands cause small Δ, favoring high-spin complexes. Spin depends on whether Δ exceeds the pairing energy.

16. Write the spectrochemical series order from weak-field to strong-field for common ligands.

    I- < Br- < S^2- < Cl- < F- < OH- < H2O < NH3 < en < NO2- (N-bonded) < CN- < CO (weak field to strong field).

17. Why are alkali metals stored under kerosene and what is their characteristic flame color?

    Alkali metals are highly reactive with air and water, so they are stored under kerosene to prevent oxidation. Flame colors: Li - crimson red, Na - golden yellow, K - lilac/violet, Rb - red-violet, Cs - blue.

18. What is diagonal relationship, and name one pair showing it in the s-block.

    A diagonal relationship is the similarity in properties between an element and the element diagonally placed in the next group/period, due to similar charge/radius ratios. Examples: Li-Mg, Be-Al, B-Si.

19. Why is beryllium anomalous among alkaline earth metals?

    Be is anomalous due to its small size, high charge density, high ionization energy, and absence of d-orbitals: it forms predominantly covalent compounds, does not react readily with water, and shows a diagonal relationship with aluminium.

20. What products form when sodium reacts with excess oxygen, and what are the analogous products for K, Rb, Cs?

    Lithium forms the oxide (Li2O), sodium forms the peroxide (Na2O2), and K, Rb, Cs form superoxides (e.g., KO2). This trend reflects stabilization of larger anions by larger cations.

21. Why does the stability of carbonates and the solubility of hydroxides change down Group 2?

    Down Group 2, thermal stability of carbonates increases (larger cation stabilizes the large carbonate anion) and solubility of hydroxides increases (lattice energy decreases faster than hydration energy), while sulphate solubility decreases.

22. What is the inert pair effect and where is it most prominent in the p-block?

    The inert pair effect is the reluctance of the ns^2 electron pair to participate in bonding in heavier p-block elements, making lower oxidation states more stable down a group (e.g., Tl+ > Tl3+, Pb2+ > Pb4+, Bi3+ > Bi5+).

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