Why is the 1st ionization enthalpy of Nitrogen higher than Oxygen despite Oxygen having a larger atomic number? Why does Chlorine release more energy than Fluorine when gaining an electron? How do ionic radii vary among isoelectronic species? Classification of Elements and Periodicity in Properties forms Chapter 3 of the CBSE Class 11 Chemistry syllabus and is a crucial foundation for JEE Main, JEE Advanced, and NEET.
This comprehensive guide covers the evolution of periodic law from Mendeleev to Henry Moseley, IUPAC nomenclature for superheavy elements (Z > 100), $s, p, d, f$ block configurations, periodic trends (atomic/ionic radii, ionization enthalpy, electron gain enthalpy, electronegativity), famous chemical exceptions, diagonal relationships, and five step-by-step solved entrance exam questions.
- 1. Evolution of Periodic Law & Moseley's Experiment
- 2. IUPAC Nomenclature for Elements with Z > 100
- 3. Block Classification: s, p, d, and f Blocks
- 4. Atomic Radii & Isoelectronic Series Ionic Radii
- 5. Ionization Enthalpy (ΔiH) & Key Exceptions (Be vs B, N vs O)
- 6. Electron Gain Enthalpy (ΔegH) & Electronegativity Trends
- 7. Diagonal Relationship & Chemical Reactivity Trends
- 8. Solved Entrance Exam Questions (JEE / NEET)
- 9. Frequently Asked Questions (FAQ)
1. Evolution of Periodic Law & Moseley's Experiment
Dmitri Mendeleev (1869) arranged elements based on atomic mass. However, in 1913, Henry Moseley studied X-ray emission spectra of elements and proved that atomic number (Z) is a more fundamental property than atomic mass:
"The physical and chemical properties of elements are periodic functions of their atomic numbers."
Cause of Periodicity: Recurrence of similar outer valence electronic configurations at regular intervals of atomic numbers ($2, 8, 8, 18, 18, 32$).
2. IUPAC Nomenclature for Elements with Z > 100
To avoid naming disputes, IUPAC devised a systematic nomenclature derived directly from atomic number digits:
| Digit | Root | Abbreviation | Atomic Number (Z) | IUPAC Systematic Name | Official Symbol |
|---|---|---|---|---|---|
| 0 | nil | n | Z = 101 | Unnilunium | Unu (Mendelevium, Md) |
| 1 | un | u | Z = 104 | Unnilquadium | Unq (Rutherfordium, Rf) |
| 2 | bi | b | Z = 111 | Unununium | Uuu (Roentgenium, Rg) |
| 3 | tri | t | Z = 114 | Ununquadium | Unq (Flerovium, Fl) |
| 4 | quad | q | Z = 118 | Ununoctium | Uuo (Oganesson, Og) |
3. Block Classification: s, p, d, and f Blocks
• p-Block (Groups 13 to 18): General outer config: $ns^2 np^{1-6}$. Includes metals, metalloids, non-metals, halogens (Grp 17) & noble gases (Grp 18). $s + p$ blocks together are called Representative / Main Group Elements.
• f-Block (Inner Transition Elements): General config: $(n-2)f^{1-14} (n-1)d^{0-1} ns^2$. Includes 4f-Lanthanoids (Ce to Lu) & 5f-Actinoids (Th to Lr).
4. Atomic Radii & Isoelectronic Series Ionic Radii
• Down a Group (Top to Bottom): Atomic radius increases due to addition of new principal energy shells ($n$).
• Cation vs Anion Radius: Cation radius < Parent atom radius (loss of $e^-$ increases $Z_{\text{eff}}$); Anion radius > Parent atom radius (gain of $e^-$ increases inter-electronic repulsion).
Isoelectronic Series Ionic Radii Rule
Isoelectronic Species are ions/atoms having the same number of electrons (e.g. 10 electrons in $\text{N}^{3-}, \text{O}^{2-}, \text{F}^-, \text{Na}^+, \text{Mg}^{2+}, \text{Al}^{3+}$):
| Isoelectronic Ion | Atomic Number (Z) | Number of Electrons | Nuclear Charge (Protons) | Ionic Radius (Å) & Trend Order |
|---|---|---|---|---|
| N³¯ (Nitride) | Z = 7 | 10 e¯ | +7 | 1.71 Å (Largest Size) |
| O²¯ (Oxide) | Z = 8 | 10 e¯ | +8 | 1.40 Å |
| F¯ (Fluoride) | Z = 9 | 10 e¯ | +9 | 1.36 Å |
| Na± (Sodium ion) | Z = 11 | 10 e¯ | +11 | 1.02 Å |
| Mg²± (Magnesium ion) | Z = 12 | 10 e¯ | +12 | 0.72 Å |
| Al³± (Aluminium ion) | Z = 13 | 10 e¯ | +13 | 0.54 Å (Smallest Size) |
Rule: For isoelectronic species, higher positive nuclear charge (Z) = smaller ionic radius!
5. Ionization Enthalpy (ΔiH) & Key Exceptions (Be vs B, N vs O)
Ionization Enthalpy ($\Delta_i H$) is the minimum energy required to remove the most loosely bound electron from an isolated gaseous atom in its ground state: $X(g) + \Delta_i H \rightarrow X^+(g) + e^-$.
• Reason: Beryllium has fully-filled stable $2s^2$ subshell ($1s^2 2s^2$) with high penetration power. Boron ($1s^2 2s^2 2p^1$) requires removing a less tightly held $2p$ electron!
• Reason: Nitrogen has extra stable half-filled $2p^3$ subshell ($1s^2 2s^2 2p_x^1 2p_y^1 2p_z^1$). Oxygen ($1s^2 2s^2 2p^4$) has one paired $2p$ orbital experiencing inter-electronic repulsion!
6. Electron Gain Enthalpy (ΔegH) & Electronegativity Trends
• Halogen Anomaly: $\Delta_{eg} H (\text{Chlorine}) = -349\text{ kJ/mol}$ is MORE NEGATIVE than $\Delta_{eg} H (\text{Fluorine}) = -328\text{ kJ/mol}$!
Reason: Fluorine atom has a very small $2n=2$ shell ($2p$ orbital). Adding an incoming electron into this compact $2p$ subshell experiences strong inter-electronic repulsions, lowering energy release compared to the larger $3p$ orbital of Chlorine.
• Pauling Scale Values: Fluorine ($F = 4.0$, Most Electronegative!) > Oxygen ($O = 3.5$) > Nitrogen ($N = 3.0$) = Chlorine ($Cl = 3.0$) > Bromine ($Br = 2.8$) > Carbon ($C = 2.5$) = Sulfur ($S = 2.5$) > Hydrogen ($H = 2.1$).
• Trend: Increases across a period (left to right) and decreases down a group.
7. Diagonal Relationship & Chemical Reactivity Trends
Certain elements of the 2nd period show striking chemical similarities with elements of the 3rd period lying diagonally opposite to them:
• Beryllium (Be) & Aluminium (Al)
• Boron (B) & Silicon (Si)
2. Nearly identical polarising power ($\text{Charge} / \text{Radius}^2$).
3. Similar electronegativities ($\text{Be} = 1.5, \text{Al} = 1.5$).
8. Solved Entrance Exam Questions (JEE / NEET)
Nuclear charge (protons $Z$): $\text{Al}^{3+} (Z=13) > \text{Mg}^{2+} (Z=12) > \text{Na}^+ (Z=11) > \text{F}^- (Z=9) > \text{O}^{2-} (Z=8) > \text{N}^{3-} (Z=7)$.
Higher nuclear charge pulls electrons closer, shrinking ionic radius.
Increasing order of radius: $\mathbf{\text{Al}^{3+} < \text{Mg}^{2+} < \text{Na}^+ < \text{F}^- < \text{O}^{2-} < \text{N}^{3-}}$.
• Nitrogen ($Z=7$): $1s^2 2s^2 2p^3$ (Exactly half-filled $2p^3$ subshell with maximum exchange energy).
• Oxygen ($Z=8$): $1s^2 2s^2 2p^4$ (Has one paired orbital $2p_x^2$ experiencing inter-electronic repulsions).
Removing an electron from Nitrogen requires breaking half-filled stability, while removing an electron from Oxygen relieves paired repulsion to yield a stable $2p^3$ configuration!
• 1 = un
• 1 = un
• 7 = sept
Systematic IUPAC Name: Ununseptium.
Official IUPAC Symbol: Uus (Official element name: Tennessine, Ts).
Reason: Fluorine atom ($2p$) is extremely small in size. Adding an 8th electron into Fluorine's compact $2p$ subshell creates strong electron-electron repulsion, which partially offsets energy release. In Chlorine ($3p$), the larger volume accommodates the electron with minimal repulsion.
2. Oxides ($\text{BeO}, \text{Al}_2\text{O}_3$) and hydroxides ($\text{Be(OH)}_2, \text{Al(OH)}_3$) of both elements are amphoteric in nature.
3. Polymeric hydrides and chlorides ($\text{BeCl}_2, \text{Al}_3\text{Cl}_6$) have electron-deficient halogen-bridged structure.
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9. Frequently Asked Questions (FAQ)
The Modern Periodic Law states that the physical and chemical properties of elements are periodic functions of their atomic numbers ($Z$). Henry Moseley discovered this by plotting $\sqrt{\nu}$ of X-ray spectra against atomic number.
Beryllium has a fully-filled stable $2s^2$ electronic configuration ($1s^2 2s^2$) with strong penetration power, whereas Boron ($1s^2 2s^2 2p^1$) loses a less tightly bound $2p^1$ electron easily.
Because Fluorine's small $2p$ orbital suffers strong inter-electronic repulsions when gaining an electron. Chlorine's larger $3p$ orbital accommodates the incoming electron with less repulsion, releasing more energy (-349 kJ/mol vs -328 kJ/mol).
In an isoelectronic series, ionic radius decreases as nuclear charge ($Z$, number of protons) increases, because a stronger positive nucleus pulls the same number of electrons tighter.
Diagonal relationship occurs between 2nd and 3rd period diagonal pairs (Li-Mg, Be-Al, B-Si) due to their nearly equal atomic/ionic radii, similar polarising power ($\text{charge}/\text{radius}^2$), and equal electronegativity values.
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