Inorganic Chemistry Exceptions for JEE Main 2027: Revision Guide

Topic

General Trend

Exception

Electronic Configuration

Orbitals fill according to the Aufbau principle.

Chromium (Cr): $[\mathrm{Ar}] 3 \mathrm{~d}^5 4 \mathrm{~s}^1$ instead of $3 \mathrm{~d}^4 4 \mathrm{~s}^2$; Copper (Cu): $[\mathrm{Ar}] 3 \mathrm{~d}^{10} 4 \mathrm{~s}^1$ instead of $3 \mathrm{~d}^9 4 \mathrm{~s}^2$ due to extra stability of half-filled and completely filled d-orbitals.

First Ionisation Enthalpy

Increases across a period.

Be > B because the electron removed from B is a 2p electron, which is easier to remove than Be's 2s electron.

First Ionisation Enthalpy

Increases across a period.

N > O because oxygen has paired 2p electrons, causing electron-electron repulsion and making electron removal easier.

Electron Gain Enthalpy

Becomes more negative across a period.

Be, Mg, and Noble Gases have positive or nearly zero electron gain enthalpy due to stable electronic configurations.

Electron Gain Enthalpy

More negative down a group.

Chlorine has a more negative electron gain enthalpy than Fluorine, although fluorine is more electronegative. The very small size of fluorine causes strong electron-electron repulsion.

Atomic Radius

Increases down a group.

Gallium (Ga) has a slightly smaller atomic radius than Aluminium (Al) due to poor shielding by d-electrons (d-block contraction).

Electronegativity

Generally decreases down a group.

Noble gases usually have no standard electronegativity values because they rarely form compounds.

Metallic Character

Increases down a group.

Hydrogen is a unique element and cannot be classified strictly as a metal or non-metal.

Oxidation States

Group number often indicates the highest oxidation state.

Fluorine always shows -1 oxidation state in its compounds because it is the most electronegative element.

Position in the Periodic Table

Elements show similar properties within a group.

Hydrogen resembles both Group 1 and Group 17, but it is placed separately due to its unique properties.

Diagonal Relationship

Not common in the periodic table.

Li–Mg, Be–Al, and B–Si show similar properties because of comparable ionic size, charge density, and electronegativity.

Topic

General Rule

Exception

Octet Rule

Atoms complete their octet.

Incomplete octet: $\mathrm{BeCl}_2, \mathrm{BF}_3$, and $\mathrm{AlCl}_3$ have less than 8 electrons around the central atom.

Expanded Octet

Period 2 elements cannot expand their octet.

Period 3 and beyond can expand the octet, e.g., $\mathrm{PCl}_5, \mathrm{SF}_6, \mathrm{ClF}_3, \mathrm{XeF}_4$.

Bond Angle

A lone pair always decreases the bond angle.

H₂O (104.5°) < NH₃ (107°) < CH₄ (109.5°) due to increasing lone pair repulsion.

Bond Order

Bond order is always an integer.

Molecular orbital theory predicts fractional bond orders, e.g., $\mathrm{O}_2^{+}(2.5), \mathrm{O}_2^{-}(1.5)$

Paramagnetism

Molecules with all paired electrons are diamagnetic.

O₂ is paramagnetic because it contains two unpaired electrons in π* antibonding orbitals.

Topic

General Rule

Exception

Reaction with Oxygen

Alkali metals form normal oxides (M₂O).

Li forms oxide (Li₂O), Na mainly forms peroxide $\left(\mathrm{Na}_2 \mathrm{O}_2\right)$, while K, Rb, and Cs form superoxides ($\left(\mathrm{KO}_2, \mathrm{RbO}_2, \mathrm{C} s \mathrm{O}_2\right)$..

Nature of Hydroxides

All Group 2 hydroxides have similar solubility.

$\mathrm{Mg}(\mathrm{OH})_2$ is sparingly soluble, whereas $\mathrm{Ba}(\mathrm{OH})_2$ is highly soluble. Solubility increases down the group.

Thermal Stability of Carbonates

Carbonates are generally stable on heating.

$\mathrm{Li}_2 \mathrm{CO}_3$ decomposes on heating, whereas other alkali metal carbonates are thermally stable.

Thermal Stability of Nitrates

Alkali metal nitrates decompose similarly.

$\mathrm{LiNO}_3$ decomposes to $\mathrm{Li}_2 \mathrm{O}$, while other alkali metal nitrates decompose to nitrites (MNO₂).

Carbonates

All Group 2 carbonates decompose equally.

Thermal stability of Group 2 carbonates increases down the group $\left(\mathrm{BeCO}_3<\mathrm{MgCO}_3<\mathrm{CaCO}_3<\mathrm{SrCO}_3<\mathrm{BaCO}_3\right)$..

Hydroxides

Basic strength remains similar.

Basic strength of hydroxides increases down Group 2: $\mathrm{Be}(\mathrm{OH})_2<\mathrm{Mg}(\mathrm{OH})_2<\mathrm{Ca}(\mathrm{OH})_2<\mathrm{Sr}(\mathrm{OH})_2<\mathrm{Ba}(\mathrm{OH})_2$.

Diagonal Relationship

Elements in different groups show different properties.

Lithium resembles Magnesium more than other alkali metals due to the diagonal relationship.

Topic

General Rule

Exception

Oxidation State

The group oxidation state remains constant.

Due to the inert pair effect, heavier p-block elements prefer lower oxidation states

$\left(\mathrm{Tl}^{+}>\mathrm{Tl}^{3+}, \mathrm{Pb}^{2+}>\mathrm{Pb}^{4+}, \mathrm{Bi}^{3+}>\mathrm{Bi}^{5+}\right)$.

Oxides of Nitrogen

All nitrogen oxides are coloured.

$\mathrm{N}_2 \mathrm{O}$ and NO are colourless, while NO₂ is brown.

Acid Strength

Hydracid strength follows electronegativity.

HF is the weakest acid among hydrogen halides despite fluorine being the most electronegative element because of its very strong H–F bond.

Reducing Power of Hydrogen Halides

Higher electronegativity means a stronger reducing agent.

Reducing power increases as

HI > HBr > HCl > HF.

Oxyacids of Halogens

All halogens form all oxyacids.

Fluorine does not form oxyacids because it cannot exhibit positive oxidation states.

Halogen Oxidation States

All halogens show positive oxidation states.

Fluorine shows only −1 oxidation state in its compounds.

Maximum Covalency

All elements can expand their octet.

Second-period elements (C, N, O, F) cannot expand their octet, while heavier p-block elements can.

Oxides

All oxides are acidic.

CO, NO and $\mathrm{N}_2 \mathrm{O}$ are neutral oxides.

Acidic Character

The acidity of oxides increases across a period.

$\mathrm{Al}_2 \mathrm{O}_3$ is amphoteric, not purely acidic.

Noble Gases

Noble gases are completely inert.

Xe forms compounds such as $\mathrm{XeF}_2, \mathrm{XeF}_4$, $\mathrm{XeF}_6, \mathrm{XeO}_3$ and $\mathrm{XeOF}_4$.. Krypton forms only a few compounds $\left(\mathrm{KrF}_2\right)$.

Topic

General Rule

Exception

Electronic Configuration

Electrons fill orbitals according to the Aufbau principle.

$C r:[A r] 3 d^5 4 s^1$ and $C u:[A r] 3 d^{10} 4 s^1$ show exceptional electronic configurations due to the extra stability of half-filled and filled d-orbitals.

Oxidation States

Elements show only one or two oxidation states.

Mn exhibits the maximum oxidation state of +7, while most transition elements show multiple oxidation states.

Colour of Ions

All transition metal ions are coloured.

$\mathrm{Sc}^{3+}, \mathrm{Ti}^{4+}, \mathrm{Zn}^{2+}, \mathrm{Cu}^{+}$and $\mathrm{Cd}^{2+}$ are colourless because they have d⁰ or d¹⁰ electronic configurations.

Magnetic Properties

All transition metal ions are paramagnetic.

$\mathrm{Zn}^{2+}, \mathrm{Cd}^{2+}, \mathrm{Hg}^{2+}, \mathrm{Sc}^{3+}$ and $\mathrm{Ti}^{4+}$ are diamagnetic due to the absence of unpaired electrons.

Lanthanoid Contraction

Atomic size increases down a group.

Lanthanoid contraction causes almost identical sizes of Zr and Hf, leading to very similar chemical properties.

Oxidation States of Lanthanoids

All lanthanoids show a +3 oxidation state only.

Ce also shows +4, while Eu and Yb commonly show +2 oxidation states due to extra electronic stability.

Transition Elements

All d-block elements are transition elements.

Zn, Cd and Hg are d-block elements but not transition elements because their atoms and common ions have filled d-orbitals ($\left(d^{10}\right)$.

Ion Formation

3d electrons are removed first during ionisation.

4s electrons are removed before 3d electrons when transition metals form cations.

Topic

General Rule

Exception

Ligand Name

Ligand names usually end with -o.

Neutral ligands have special names:

$\mathrm{H}_2 \mathrm{O} \rightarrow$ aqua, $\mathrm{NH}_3 \rightarrow$ ammine, $\mathrm{CO} \rightarrow$ carbonyl, NO → nitrosyl.

Anionic Complex

The metal name remains unchanged.

In anionic complexes, the metal name ends with -ate (Fe → ferrate, Cu → cuprate, Ag → argentate, Au → aurate).

Oxidation Number

Ligands are neutral.

NO (nitrosyl) can behave as $\mathrm{NO}^{+}$ (+1), unlike most neutral ligands.

Chelating Ligands

All ligands are monodentate.

en, oxalate $\left(\mathrm{C}_2 \mathrm{O}_4^{2-}\right)$), and EDTA are multidentate (chelating) ligands.

Magnetic Behaviour

All complexes of a metal have the same number of unpaired electrons.

The same metal ion can form high-spin or low-spin complexes depending on ligand strength (e.g., $\mathrm{Fe}^{2+}, \mathrm{Co}^{3+}$ ).

Hybridization

Hybridisation depends only on the coordination number.

Strong-field ligands $(\mathrm{CN}, \mathrm{CO})$ cause electron pairing, whereas weak-field ligands $\left(\mathrm{F}^{-}, \mathrm{Cl}^{-}, \mathrm{H}_2 \mathrm{O}\right)$ usually do not.

Coordination Number

Ligands contribute one donor atom.

Polydentate ligands contribute multiple donor atoms; e.g., EDTA has a coordination number of 6.

Isomerism

All coordination compounds show isomerism.

Tetrahedral complexes generally do not show geometrical isomerism, whereas square planar and octahedral complexes often do.

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