JEE Main 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 JEE Main, not a summary of it.

1. 1

   Classification of Elements and Periodicity in Properties

   3 topics

   • Modem periodic law and present form of the periodic table
   • s, p, d and f block elements
   • Periodic trends in properties of elements
     • Atomic and ionic radii
     • Ionization enthalpy
     • Electron gain enthalpy
     • Valence
     • Oxidation states and chemical reactivity
2. 2

   P-Block Elements

   1 topic

   • Group -13 to Group 18 Elements
     • Electronic configuration and general trends
     • Unique behaviour of the first element in each group
3. 3

   d - and f- Block Elements

   2 topics

   • Transition Elements
     • General introduction, electronic configuration, occurrence and characteristics
     • General trends in properties of the first-row transition elements
     • Physical properties, ionization enthalpy, oxidation states, atomic radii, colour, catalytic behaviour, magnetic properties, complex formation, interstitial compounds, alloy formation
     • Preparation, properties, and uses of K2Cr2O7, and KMnO4
   • Inner Transition Elements
     • Lanthanoids - Electronic configuration, oxidation states, and lanthanoid contraction
     • Actinoids - Electronic configuration and oxidation states
4. 4

   Co-ordination Compounds

   4 topics

   • Introduction to coordination compounds
     • Werner's theory
     • Ligands, coordination number, denticity, chelation
   • IUPAC nomenclature of mononuclear co-ordination compounds, isomerism
   • Bonding-Valence bond approach and basic ideas of Crystal field theory, colour and magnetic properties
   • Importance of co-ordination compounds
     • In qualitative analysis
     • Extraction of metals
     • In biological systems

Inorganic Chemistry flashcards for JEE Main

22 of 50 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 (number of protons), i.e. properties repeat at regular intervals when elements are arranged in order of increasing atomic number $Z$.

2. How is the long form (modern) periodic table organized in terms of periods and groups?

   It has 7 horizontal periods and 18 vertical groups. The period number equals the principal quantum number $n$ of the outermost shell being filled; elements in a group have similar valence-shell electronic configurations and hence similar properties.

3. On what basis are the s, p, d and f blocks defined?

   They are named after the subshell into which the differentiating (last) electron enters: s-block (last electron in $ns$), p-block ($np$), d-block ($(n-1)d$), and f-block ($(n-2)f$).

4. Which groups make up the s-block, and what is its general valence-shell configuration?

   Groups 1 (alkali metals) and 2 (alkaline earth metals). General configuration $ns^{1\text{--}2}$.

5. Which groups make up the p-block and what is its general valence configuration?

   Groups 13 to 18. General valence-shell configuration $ns^{2}np^{1\text{--}6}$ (for He it is $1s^{2}$).

6. What is the general valence configuration of d-block (transition) elements?

   $(n-1)d^{1\text{--}10}\,ns^{0\text{--}2}$, where $n$ is the outermost principal shell. They lie in groups 3 to 12.

7. What is the general valence configuration of f-block elements and what are the two series called?

   $(n-2)f^{1\text{--}14}(n-1)d^{0\text{--}1}ns^{2}$. The two series are the lanthanoids ($\ce{^{4f}}$, Ce to Lu) and actinoids ($\ce{^{5f}}$, Th to Lr); they are the inner transition elements.

8. 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 with same shell) and increases down a group (new shells added).

9. Why is the ionic radius of a cation smaller and an anion larger than the parent atom?

   A cation loses electrons, increasing effective nuclear charge per electron and often removing a shell, so it shrinks. An anion gains electrons, increasing electron-electron repulsion and effective size, so it grows larger than the neutral atom.

10. For an isoelectronic series such as $\ce{O^{2-}, F^-, Na^+, Mg^{2+}, Al^{3+}}$, how does ionic radius vary?

    All have the same number of electrons (10). Ionic radius decreases as nuclear charge $Z$ increases, so $\ce{O^{2-}} > \ce{F^-} > \ce{Na^+} > \ce{Mg^{2+}} > \ce{Al^{3+}}$.

11. Define first ionization enthalpy ($\Delta_i H_1$).

    The energy required to remove the most loosely bound electron from one mole of isolated gaseous atoms in their ground state: $\ce{X(g) -> X^+(g) + e^-}$. It is always positive (endothermic).

12. How does ionization enthalpy generally vary across a period and down a group?

    It increases across a period (rising $Z_{eff}$, smaller size) and decreases down a group (larger size and increased shielding make outer electron easier to remove).

13. Why is the first ionization enthalpy of nitrogen higher than that of oxygen?

    $\ce{N}$ has a stable half-filled $2p^{3}$ configuration which is hard to disturb, whereas removing an electron from $\ce{O}$ ($2p^{4}$) relieves electron-electron repulsion in a paired orbital, so O loses its electron more readily.

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

    In $\ce{Be}$ ($2s^{2}$) the electron is removed from a stable fully filled 2s orbital, whereas in $\ce{B}$ ($2s^{2}2p^{1}$) the 2p electron is higher in energy and better shielded, so it is removed more easily.

15. Define electron gain enthalpy ($\Delta_{eg}H$) and give its sign convention.

    The enthalpy change when an electron is added to a neutral gaseous atom: $\ce{X(g) + e^- -> X^-(g)}$. It is negative if energy is released (favourable) and positive if energy must be supplied.

16. Why is the electron gain enthalpy of chlorine more negative than that of fluorine?

    Fluorine's small 2p subshell causes strong electron-electron repulsion on adding an electron, reducing the energy released. Chlorine's larger 3p orbital accommodates the extra electron with less repulsion, giving a more negative $\Delta_{eg}H$.

17. Why do noble gases have large positive electron gain enthalpies?

    They have completely filled stable octet (or duplet) configurations; the incoming electron must enter a higher-energy new shell, so energy is absorbed rather than released.

18. Define valence in terms of electronic configuration and state how it varies for s- and p-block elements across a period.

    Valence is the combining capacity, related to the number of valence electrons or electrons needed to complete the octet. For representative elements it rises from 1 to 4 and then falls to 0 across a period (e.g. Na=1, ..., C=4, ..., Ne=0).

19. What is the general trend in metallic and non-metallic character across a period and down a group?

    Metallic (electropositive) character decreases across a period and increases down a group; non-metallic character increases across a period and decreases down a group.

20. Why do elements show variable oxidation states, and how does this relate to chemical reactivity of metals vs non-metals?

    Variable oxidation states arise when an element can lose/share different numbers of electrons (common in d-block due to close $(n-1)d$ and $ns$ energies). Metallic reactivity (tendency to lose electrons) increases down a group; non-metal reactivity (tendency to gain electrons) increases up a group toward fluorine.

21. Name the Group 13 to Group 18 element families.

    Group 13: Boron family; Group 14: Carbon family; Group 15: Pnictogens (nitrogen family); Group 16: Chalcogens (oxygen family); Group 17: Halogens; Group 18: Noble gases.

22. What is the inert pair effect and where is it most prominent?

    The reluctance of the outer $ns^{2}$ electrons to participate in bonding due to poor shielding by intervening d and f electrons. It increases down groups 13-16, stabilizing lower oxidation states (e.g. $\ce{Pb^{2+}} > \ce{Pb^{4+}}$, $\ce{Tl^+} > \ce{Tl^{3+}}$).

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