WBJEE 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 WBJEE, not a summary of it.

1. 1

   Periodicity and Chemical Bonding

   3 topics

   • Periodic Table and Periodicity
     • Modern periodic law and classification
     • Periodic trends in atomic and physical properties
   • Chemical Bonding and Molecular Structure
     • Ionic and covalent bonding
     • VSEPR theory and hybridization
     • Molecular orbital theory
     • Hydrogen bonding and dipole moment
   • Hydrogen and its Compounds
     • Position of hydrogen and isotopes
     • Water, hydrogen peroxide and hydrides
2. 2

   s and p Block Elements

   3 topics

   • s-Block Elements
     • Alkali metals and their compounds
     • Alkaline earth metals and their compounds
   • p-Block Elements (Groups 13-15)
     • Boron and carbon families
     • Nitrogen family and compounds
   • p-Block Elements (Groups 16-18)
     • Oxygen and sulphur families
     • Halogens and their compounds
     • Noble gases
3. 3

   Transition Elements and Coordination Chemistry

   3 topics

   • d and f Block Elements
     • Transition elements and their properties
     • Lanthanides and actinides
   • Coordination Compounds
     • Nomenclature and isomerism
     • Werner's theory and valence bond theory
     • Crystal field theory
   • Metallurgy
     • Principles and methods of extraction
     • Thermodynamic and electrochemical principles
4. 4

   Applied Inorganic Chemistry

   3 topics

   • Qualitative Analysis
     • Detection of cations and anions
     • Group separation of salts
   • Environmental Chemistry
     • Atmospheric and water pollution
     • Green chemistry
   • Chemistry in Everyday Life
     • Drugs and their classification
     • Chemicals in food and cleansing agents

Inorganic Chemistry flashcards for WBJEE

22 of 53 cards from the Inorganic Chemistry deck — real questions with worked answers.

1. State the modern periodic law.

   The physical and chemical properties of elements are a periodic function of their atomic numbers ($Z$), i.e. when elements are arranged in order of increasing atomic number, elements with similar properties recur at regular intervals.

2. On what basis did Mendeleev arrange elements, and how does this differ from the modern periodic table?

   Mendeleev arranged elements in order of increasing atomic mass (periodic function of atomic mass). The modern table arranges them by increasing atomic number ($Z$), which resolved anomalies like the $\ce{Ar}$/$\ce{K}$ and $\ce{Co}$/$\ce{Ni}$ pair orderings.

3. How are the four blocks (s, p, d, f) of the periodic table defined?

   By the subshell into which the last (differentiating) electron enters: s-block (groups 1-2), p-block (groups 13-18), d-block (groups 3-12, transition metals), f-block (lanthanoids and actinoids, inner-transition metals).

4. Define atomic radius and state how it varies across a period and down a group.

   Atomic radius is roughly the distance from the nucleus to the outermost electron shell. It decreases across a period (increasing effective nuclear charge, $Z_{eff}$) and increases down a group (new shells added).

5. Why is the cation smaller and the anion larger than the parent atom?

   A cation has fewer electrons than protons, so greater effective nuclear charge per electron and often one fewer shell contracts it. An anion has more electrons than protons, increasing electron-electron repulsion and shielding, expanding it. e.g. $\ce{Na+} < \ce{Na}$ and $\ce{Cl-} > \ce{Cl}$.

6. Define first ionization enthalpy and give its general trend across a period and down a group.

   It is the energy required to remove the most loosely bound electron from one mole of gaseous atoms: $\ce{M(g) -> M+(g) + e-}$. It generally increases across a period and decreases down a group.

7. Why is the first ionization enthalpy of boron lower than that of beryllium?

   $\ce{Be}$ ($1s^{2}2s^{2}$) has a stable, fully filled $2s$ subshell, whereas $\ce{B}$ ($1s^{2}2s^{2}2p^{1}$) loses an electron from the higher-energy, less-penetrating $2p$ orbital, which is easier to remove.

8. Why is the first ionization enthalpy of oxygen lower than that of nitrogen?

   $\ce{N}$ has a stable half-filled $2p^{3}$ configuration. In $\ce{O}$ ($2p^{4}$), the fourth $p$ electron pairs up, and the resulting electron-electron repulsion makes its removal easier.

9. Define electron gain enthalpy. Why is that of chlorine more negative than that of fluorine?

   Electron gain enthalpy is the enthalpy change when an electron is added to a gaseous atom: $\ce{X(g) + e- -> X^-(g)}$. $\ce{Cl}$ is more negative than $\ce{F}$ because the small size of $\ce{F}$ causes high electron-electron repulsion in its compact $2p$ subshell, reducing the energy released.

10. Define electronegativity and name the most electronegative element. State the trend across a period and down a group.

    Electronegativity is the tendency of an atom to attract the shared electron pair in a bond toward itself. The most electronegative element is fluorine ($\ce{F}$). It increases across a period and decreases down a group.

11. What are diagonal relationships? Give two examples.

    Certain second-period elements resemble the third-period element diagonally placed (down-right) due to similar charge/size ratios. Examples: $\ce{Li}$ resembles $\ce{Mg}$, and $\ce{Be}$ resembles $\ce{Al}$; also $\ce{B}$ resembles $\ce{Si}$.

12. State the octet rule and name one common exception.

    Atoms tend to gain, lose, or share electrons to attain a stable noble-gas configuration of eight electrons in their valence shell. Exceptions include incomplete octets ($\ce{BF3}$, $\ce{BeCl2}$) and expanded octets ($\ce{SF6}$, $\ce{PCl5}$).

13. Define lattice enthalpy of an ionic compound.

    The energy required to completely separate one mole of a solid ionic compound into its gaseous ions: $\ce{NaCl(s) -> Na+(g) + Cl-(g)}$. Higher charge and smaller ionic size give larger lattice enthalpy.

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

    Covalent character increases with: (1) small cation size, (2) large anion size, (3) high charge on either ion, and (4) cations with a non-noble-gas (pseudo) configuration ($18$-electron). Greater polarization of the anion by the cation increases covalent character.

15. What does VSEPR theory predict, and what is the basic geometric assumption?

    Valence Shell Electron Pair Repulsion theory predicts molecular shape based on minimizing repulsion between valence electron pairs around the central atom. Repulsion order: lone pair-lone pair $>$ lone pair-bond pair $>$ bond pair-bond pair.

16. Give the shape and bond angle of $\ce{CH4}$, $\ce{NH3}$, and $\ce{H2O}$, and explain the angle trend.

    $\ce{CH4}$: tetrahedral, $109.5^{\circ}$ (0 lone pairs). $\ce{NH3}$: trigonal pyramidal, $\sim 107^{\circ}$ (1 lone pair). $\ce{H2O}$: bent, $\sim 104.5^{\circ}$ (2 lone pairs). Increasing lone pairs increases repulsion and compresses the bond angle.

17. What hybridization and shape correspond to $\ce{sp}$, $\ce{sp^2}$, $\ce{sp^3}$, $\ce{sp^3d}$, and $\ce{sp^3d^2}$?

    $\ce{sp}$: linear ($180^{\circ}$); $\ce{sp^2}$: trigonal planar ($120^{\circ}$); $\ce{sp^3}$: tetrahedral ($109.5^{\circ}$); $\ce{sp^3d}$: trigonal bipyramidal; $\ce{sp^3d^2}$: octahedral ($90^{\circ}$).

18. Determine the hybridization of the central atom in $\ce{PCl5}$ and $\ce{SF6}$.

    $\ce{PCl5}$: $\ce{sp^3d}$ (5 bond pairs, trigonal bipyramidal). $\ce{SF6}$: $\ce{sp^3d^2}$ (6 bond pairs, octahedral).

19. What is the steric-number formula used to predict hybridization?

    Steric number = (number of sigma bonds / bonded atoms) + (number of lone pairs on the central atom). For a main-group species, $2 \to sp$, $3 \to sp^{2}$, $4 \to sp^{3}$, $5 \to sp^{3}d$, $6 \to sp^{3}d^{2}$.

20. State the formula for bond order in molecular orbital theory.

    $$\text{Bond order} = \frac{1}{2}\left(N_{b} - N_{a}\right)$$ where $N_{b}$ and $N_{a}$ are the numbers of electrons in bonding and antibonding molecular orbitals, respectively.

21. Why does the $\ce{He2}$ molecule not exist according to MO theory?

    $\ce{He2}$ has 4 electrons: configuration $\sigma_{1s}^{2}\,\sigma^{*}_{1s}^{2}$. Bond order $=\frac{1}{2}(2-2)=0$, so no net bond forms and the molecule is unstable.

22. Use MO theory to explain why $\ce{O2}$ is paramagnetic.

    $\ce{O2}$ (16 electrons) has the configuration with two unpaired electrons in the degenerate $\pi^{*}_{2p}$ antibonding orbitals: $\pi^{*}_{2p_x}^{1}\pi^{*}_{2p_y}^{1}$. Unpaired electrons make it paramagnetic; bond order $=2$.

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