GATE Chemistry 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 GATE Chemistry, not a summary of it.

1. 1

   Main Group Elements

   21 topics

   • Hydrides
   • Halides
   • Oxides
   • Oxoacids
   • Nitrides
   • Sulfides
   • Shapes and reactivity
   • Structure and bonding of boranes
   • Carboranes
   • Silicones
   • Silicates
   • Boron nitride
   • Borazines
   • Phosphazenes
   • Allotropes of carbon
   • Phosphorous and sulphur
   • Industrial synthesis of compounds of main group elements
   • Chemistry of noble gases
   • Pseudohalogens
   • Interhalogen compounds
   • Acid-base concepts and principles
     • Lewis
     • Brønsted
     • HSAB
     • Acid-base catalysis
2. 2

   Transition Elements

   6 topics

   • Coordination chemistry
     • Structure and isomerism
   • Theories of bonding
     • VBT
     • CFT
     • MOT
   • Energy level diagrams in various crystal fields
     • CFSE
     • Applications of CFT
     • Jahn-Teller distortion
   • Electronic spectra of transition metal complexes
     • Spectroscopic term symbols
     • Selection rules
     • Orgel and Tanabe-Sugano diagrams
     • Nephelauxetic effect
     • Racah parameter
     • Charge-transfer spectra
   • Magnetic properties of transition metal complexes
     • Ray-Dutt
     • Bailar twists
   • Reaction mechanisms
     • Kinetic and thermodynamic stability
     • Substitution and redox reactions
     • Metal-metal multiple bond
3. 3

   Lanthanides and Actinides

   1 topic

   • Recovery
4. 4

   Organometallics

   7 topics

   • 18-Electron rule
   • Metal-alkyl, metal-carbonyl, metal-olefin and metal-carbene complexes
   • Metallocenes
   • Fluxionality in organometallic complexes
   • Types of organometallic reactions
   • Homogeneous catalysis
     • Hydrogenation
     • Hydroformylation
     • Acetic acid synthesis
     • Metathesis
     • Olefin oxidation
   • Heterogeneous catalysis
     • Fischer-Tropsch reaction
     • Ziegler-Natta polymerization
5. 5

   Radioactivity

   4 topics

   • Detection of radioactivity
   • Decay processes
   • Half-life of radioactive elements
   • Fission and fusion processes
6. 6

   Bioinorganic Chemistry

   5 topics

   • Ion transport
     • Na+
     • K+
   • Oxygen binding, transport and utilization
   • Electron transfer reactions
   • Nitrogen fixation
   • Metalloenzymes containing magnesium, molybdenum, iron, cobalt, copper and zinc
7. 7

   Solids

   10 topics

   • Crystal systems and lattices
   • Miller planes
   • Crystal packing
   • Crystal defects
   • Bragg’s law
   • Ionic crystals
   • Structures of AX, AX2, ABX3 type compounds
   • Spinels
   • Band theory
   • Metals and semiconductors
8. 8

   Instrumental Methods of Analysis

   9 topics

   • UV-visible, fluorescence and FTIR spectrophotometry
   • NMR and ESR spectroscopy
   • Mass spectrometry
   • Atomic absorption spectroscopy
   • Mössbauer spectroscopy
     • Fe
     • Sn
   • X-ray crystallography
   • Chromatography
     • GC
     • HPLC
   • Electroanalytical methods
     • Polarography
     • Cyclic voltammetry
     • Ion-selective electrodes
   • Thermoanalytical methods

Inorganic Chemistry flashcards for GATE Chemistry

22 of 48 cards from the Inorganic Chemistry deck — real questions with worked answers.

1. What is the general trend in the thermal stability of Group 14 hydrides $\ce{EH4}$ down the group?

   Thermal stability decreases down the group: $\ce{CH4} > \ce{SiH4} > \ce{GeH4} > \ce{SnH4} > \ce{PbH4}$, because the $\ce{E-H}$ bond strength decreases as $E$ becomes larger and the orbital overlap with H worsens.

2. Classify hydrides into their three main types based on bonding, with an example of each.

   Ionic (saline) hydrides — e.g. $\ce{NaH}$, $\ce{CaH2}$ (s-block, contain $\ce{H^-}$); Covalent (molecular) hydrides — e.g. $\ce{CH4}$, $\ce{NH3}$, $\ce{H2O}$ (p-block); Metallic (interstitial) hydrides — e.g. $\ce{PdH_{x}}$, $\ce{TiH2}$ (many d/f-block metals).

3. Why does $\ce{NH3}$ have a much higher boiling point than $\ce{PH3}$ despite N being smaller?

   $\ce{NH3}$ engages in strong intermolecular hydrogen bonding (N is highly electronegative), whereas $\ce{PH3}$ shows only weak van der Waals interactions, so $\ce{NH3}$ boils much higher (−33 °C vs −88 °C).

4. How does the acidic strength of the hydrogen halides $\ce{HX}$ vary, and why?

   Acid strength increases $\ce{HF} < \ce{HCl} < \ce{HBr} < \ce{HI}$. The dominant factor is the decreasing $\ce{H-X}$ bond dissociation energy down the group; $\ce{HF}$ is the weakest because of its strong bond and H-bonding.

5. What are the two general methods of classifying halides, and give an example of an electron-deficient halide.

   Halides are classified as ionic (e.g. $\ce{NaCl}$), covalent (e.g. $\ce{CCl4}$), and polymeric/bridged. Electron-deficient halides include $\ce{BCl3}$ and $\ce{AlCl3}$ (which dimerizes to $\ce{Al2Cl6}$ with chlorine bridges).

6. Why is $\ce{CCl4}$ not hydrolyzed by water whereas $\ce{SiCl4}$ is readily hydrolyzed?

   Carbon has no accessible d-orbitals and is sterically protected, so there is no low-energy pathway for nucleophilic attack. Silicon can expand its coordination number using vacant d-orbitals, allowing $\ce{H2O}$ attack: $\ce{SiCl4 + 2H2O -> SiO2 + 4HCl}$.

7. Distinguish between acidic, basic, amphoteric and neutral oxides with one example each.

   Acidic — $\ce{SO3}$, $\ce{CO2}$ (form acids with water/react with bases); Basic — $\ce{Na2O}$, $\ce{CaO}$ (react with acids); Amphoteric — $\ce{Al2O3}$, $\ce{ZnO}$ (react with both); Neutral — $\ce{CO}$, $\ce{N2O}$, $\ce{NO}$.

8. How does the acidic/basic character of oxides change across a period and down a group?

   Across a period (left to right) oxides change from basic → amphoteric → acidic as electronegativity rises. Down a group oxides become more basic (or less acidic) as metallic character increases.

9. Write the structures (shapes) of $\ce{SO2}$ and $\ce{SO3}$ in the gas phase.

   $\ce{SO2}$ is bent (angular), V-shaped with a lone pair on S, bond angle $\approx 119^\circ$. Gaseous $\ce{SO3}$ is trigonal planar ($D_{3h}$), bond angle $120^\circ$.

10. For oxoacids of the same central element, how does acid strength vary with oxidation state? Illustrate with chlorine oxoacids.

    Acid strength increases with the oxidation state of the central atom (more terminal O atoms stabilize the conjugate base): $\ce{HOCl} < \ce{HOClO} < \ce{HOClO2} < \ce{HOClO3}$, i.e. $\ce{HClO} < \ce{HClO2} < \ce{HClO3} < \ce{HClO4}$.

11. State Pauling's rule for estimating the strength of an oxoacid $\ce{(HO)_pXO_q}$.

    For an oxoacid $\ce{(HO)_pXO_q}$, $pK_a \approx 8 - 5q$ where $q$ is the number of non-hydroxyl (terminal) oxygens. Each additional terminal O lowers $pK_a$ by about 5 units (stronger acid).

12. Compare the structures of phosphorous acid and phosphoric acid, and give their basicity.

    Phosphorous acid $\ce{H3PO3}$ has structure $\ce{HP(O)(OH)2}$ — one P–H bond — and is dibasic (2 ionizable H). Phosphoric acid $\ce{H3PO4}$ is $\ce{OP(OH)3}$ — no P–H — and is tribasic.

13. What distinguishes covalent, ionic and interstitial nitrides? Give an example of each.

    Ionic nitrides contain $\ce{N^{3-}}$, e.g. $\ce{Li3N}$, $\ce{Mg3N2}$ (s-block). Covalent nitrides, e.g. $\ce{BN}$, $\ce{Si3N4}$. Interstitial nitrides — hard, refractory d-block compounds, e.g. $\ce{TiN}$, $\ce{VN}$.

14. What is produced when an ionic nitride such as $\ce{Mg3N2}$ reacts with water?

    Ionic nitrides hydrolyze to give ammonia and the metal hydroxide: $\ce{Mg3N2 + 6H2O -> 3Mg(OH)2 + 2NH3}$.

15. Distinguish ionic, covalent (molecular) and metallic sulfides with an example of each.

    Ionic sulfides contain $\ce{S^{2-}}$, e.g. $\ce{Na2S}$. Covalent/molecular sulfides, e.g. $\ce{CS2}$, $\ce{P4S3}$. Metallic/non-stoichiometric sulfides with metallic conduction, e.g. $\ce{FeS}$, $\ce{CoS}$. Many heavy-metal sulfides (e.g. $\ce{CuS}$, $\ce{HgS}$) are highly insoluble and covalent.

16. Using VSEPR, predict the shapes of $\ce{XeF2}$, $\ce{XeF4}$ and $\ce{XeF6}$.

    $\ce{XeF2}$: linear (trigonal bipyramidal electron geometry, 3 lone pairs equatorial). $\ce{XeF4}$: square planar (octahedral, 2 axial lone pairs). $\ce{XeF6}$: distorted octahedral (capped octahedron) due to one lone pair in 7 electron domains.

17. Predict the shape and bond angle of $\ce{ClF3}$ using VSEPR.

    $\ce{ClF3}$ is T-shaped: 5 electron domains (3 bonding + 2 lone pairs in a trigonal bipyramid, lone pairs equatorial). The F–Cl–F angles are slightly less than $90^\circ$ ($\approx 87.5^\circ$).

18. State Wade's rules for the structures of boranes ($n$ skeletal atoms) for closo, nido and arachno clusters.

    Count skeletal electron pairs (SEP). $closo$: $n+1$ pairs (complete polyhedron). $nido$: $n+2$ pairs (one vertex missing). $arachno$: $n+3$ pairs (two vertices missing). For boranes, $\ce{B_nH_n^{2-}}$ closo, $\ce{B_nH_{n+4}}$ nido, $\ce{B_nH_{n+6}}$ arachno.

19. Describe the bonding in diborane $\ce{B2H6}$ and the nature of its bridging bonds.

    $\ce{B2H6}$ has 4 terminal 2c–2e $\ce{B-H}$ bonds and 2 bridging three-center two-electron (3c–2e) $\ce{B-H-B}$ 'banana' bonds. Each B is roughly $sp^3$; the molecule is electron-deficient (only 12 valence electrons).

20. Using Wade's rules, classify $\ce{B5H9}$ and predict its structure.

    $\ce{B5H9}$: skeletal electron pairs $= \tfrac{1}{2}(5\times2 + 9 - 5\times... )$. Standard count gives $n=5$, SEP $= n+2 = 7$ → nido. Structure is a square pyramid (octahedron missing one vertex).

21. What are carboranes, and what is the formula and shape of the most famous icosahedral carborane?

    Carboranes are cluster compounds containing both carbon and boron (and hydrogen) in the cage. The classic example is closo-$\ce{C2B10H12}$ (dicarba-closo-dodecaborane), an icosahedral cage with 12 vertices; its $ortho$, $meta$, $para$ isomers differ in C positions.

22. How are the two carbon atoms of $\ce{C2B10H12}$ counted as skeletal electron donors in Wade's rules?

    Each $\ce{CH}$ unit contributes 3 skeletal electrons (C has 4 valence e, minus 1 for the terminal C–H bond, giving 3), versus 2 for each $\ce{BH}$ unit. For $\ce{C2B10H12}$: $2(3)+10(2)=26 = 13$ pairs $= n+1$ ($n=12$) → closo.

See more Inorganic Chemistry flashcards →

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Inorganic Chemistry (GATE Chemistry) FAQ