Classification of Elements and Periodicity
'The periodic table is a map of the elements — a guide to the chemical universe.' — Chemistry
1. Chapter Overview
The PERIODIC TABLE is one of the GREATEST intellectual achievements — a systematic ARRANGEMENT of all known elements that reveals PERIODIC TRENDS in properties. This chapter covers the HISTORICAL development of the periodic table (Dobereiner, Newlands, Mendeleev, Moseley), the MODERN PERIODIC LAW based on atomic number, ELECTRONIC CONFIGURATIONS of elements, and PERIODIC TRENDS in atomic radius, ionisation energy, electron gain enthalpy, and electronegativity.
2. Historical Development
Dobereiner's Triads (1817)
- Groups of THREE elements with SIMILAR properties
- Atomic weight of middle element ≈ AVERAGE of the other two
- Example: Li (7), Na (23), K (39) → (7+39)/2 = 23 ✓
Newlands' Law of Octaves (1864)
- Every EIGHTH element has similar properties (like musical notes)
- Worked only up to calcium (failed for heavier elements)
Mendeleev's Periodic Table (1869)
- Elements arranged by INCREASING ATOMIC MASS
- Periodic Law: Properties of elements are PERIODIC function of atomic mass
- Left GAPS for undiscovered elements (predicted Ga, Ge, Sc)
- Limitations: Position of isotopes, anomalous pairs (Ar-K, Co-Ni)
Moseley's Modern Periodic Law (1913)
- Modern Law: Properties are PERIODIC function of ATOMIC NUMBER (not mass)
- Based on X-ray spectra studies
- Eliminated the anomalies in Mendeleev's table
3. Modern Periodic Table
Blocks Based on Electronic Configuration
| Block | Last electron enters | Groups | Elements |
|---|---|---|---|
| s | ns¹-² | 1, 2 | Alkali metals, alkaline earth metals, H, He |
| p | np¹-⁶ | 13-18 | Non-metals, metalloids, noble gases |
| d | (n-1)d¹-¹⁰ | 3-12 | Transition metals |
| f | (n-2)f¹-¹⁴ | Inner transition | Lanthanoids and Actinoids |
Periods
- Period 1: 2 elements (H, He)
- Period 2,3: 8 elements each
- Period 4,5: 18 elements each
- Period 6: 32 elements (includes lanthanoids)
- Period 7: Incomplete
Groups
- 18 groups (1-18, IUPAC numbering)
- Elements in SAME group have SIMILAR chemical properties (same number of valence electrons)
4. Periodic Trends in Physical Properties
Atomic Radius
- Decreases ACROSS a period (↑ Z, ↑ nuclear charge, electrons added to same shell)
- Increases DOWN a group (↑ new shells, ↑ shielding)
- Types: Covalent radius, metallic radius, van der Waals radius
Ionic Radius
- Cations are SMALLER than parent atoms (loss of outer shell)
- Anions are LARGER than parent atoms (gained electrons, more repulsion)
Ionisation Enthalpy (ΔᵢH)
- Definition: Energy required to remove the MOST LOOSELY bound electron from an isolated atom
- Increases ACROSS a period (↑ nuclear charge, ↓ size)
- Decreases DOWN a group (↑ size, ↑ shielding)
- Exceptions: Group 2 > Group 3 (ns² stability), Group 15 > Group 16 (np³ half-filled stability)
- IE₁ < IE₂ < IE₃ (each removal gets harder)
Electron Gain Enthalpy (Δₑ₉H)
- Definition: Energy change when an electron is added to an isolated atom
- Most negative (favourable) for Group 17 (halogens want one electron)
- Noble gases have POSITIVE Δₑ₉H (already stable)
- Variation: Becomes more negative ACROSS a period (↑ Z, ↓ size)
Electronegativity (χ)
- Definition: Tendency of an atom to ATTRACT shared electrons in a bond
- Increases ACROSS a period, decreases DOWN a group
- Pauling scale: F (4.0) > O (3.5) > N (3.0) > Cl (3.0)
- Most electronegative: Fluorine
- Least electronegative: Francium / Caesium
5. Periodic Trends Comparison Table
| Property | Across Period (L→R) | Down Group (T→B) |
|---|---|---|
| Atomic Radius | DECREASES | INCREASES |
| Ionisation Enthalpy | INCREASES* | DECREASES |
| Electron Gain Enthalpy | More NEGATIVE | Less negative** |
| Electronegativity | INCREASES | DECREASES |
| Metallic Character | DECREASES | INCREASES |
| Non-metallic Character | INCREASES | DECREASES |
*With exceptions (Be-B, N-O)
**With exceptions (Cl > F in some cases due to small size of F)
6. s, p, d, f Block Elements — Key Features
s-Block Elements
- Group 1 (Alkali metals): ns¹, VERY reactive, strong reducing agents
- Group 2 (Alkaline earth metals): ns², less reactive than Group 1
p-Block Elements
- Include metals, non-metals, and metalloids
- Group 17 (Halogens): HIGHEST electronegativity
- Group 18 (Noble gases): COMPLETELY filled p-subshell (inert)
d-Block Elements (Transition Metals)
- Variable oxidation states
- Form COLOURED compounds
- Catalytic activity
- Paramagnetism (unpaired electrons)
f-Block Elements (Inner Transition)
- Lanthanoids: 4f, similar properties across series (lanthanoid contraction)
- Actinoids: 5f, RADIOACTIVE
7. Common Mistakes
- IE of N > O: N has HALF-FILLED 2p³ (stable), O has 2p⁴ (electron-electron repulsion makes removal easier)
- EG of Cl is more negative than F: F's small size causes electron-electron repulsion, making electron gain less favourable despite higher EN
- Atomic radius of noble gases is van der Waals radius, not covalent: Noble gas radii appear larger than halogen radii in the same period
- IE₂ >> IE₁: Removing an electron from a CATION requires much more energy than from a neutral atom
- Metallic character ≠ number of metals: It describes tendency to lose electrons (form cations)
8. CBSE Exam Focus
- Modern periodic law — grouping into s, p, d, f blocks (3-mark)
- Periodic trends — atomic radius, IE, EG, EN (5-mark)
- Reasons for exceptions in IE (N > O, Be > B)
- Comparison of properties across period and down group (5-mark)
- Anomalous properties of second-period elements
- Mendeleev's table — merits and limitations
9. Key Formulas (Conceptual)
- IE₁ < IE₂ < IE₃... (successive IEs increase)
- Pauling EN: χ_A — χ_B = 0.208√(Δ) where Δ = bond energy deviation
- Screening (shielding) constant: σ (Slater's rules)
- Z_eff = Z — σ
10. Self-Test (5+ Q&A)
Q1: Arrange Na, Mg, Al, Si in increasing order of IE₁. A: Na < Mg < Al < Si. Actually: Na < Al < Mg < Si (Mg has stable 3s², so higher than Al).
Q2: Why is the ionic radius of Cl⁻ larger than that of Cl? A: Cl⁻ has gained an electron → increased electron-electron repulsion → EXPANDED electron cloud.
Q3: What is the lanthanoid contraction? What is its consequence? A: Gradual DECREASE in atomic/ionic radius across lanthanoid series (poor shielding of 4f electrons). Consequence: Zr and Hf have NEARLY IDENTICAL radii and very similar properties.
Q4: Which element has the highest electronegativity and why? A: FLUORINE (F = 4.0 on Pauling scale). It has small size, high nuclear charge, and requires only ONE electron to achieve noble gas configuration.
Q5: Explain the anomalous trend in IE: Boron (B) has lower IE than Beryllium (Be). A: Be has [He]2s² configuration (FILLED s-subshell, extra stability). B has [He]2s²2p¹ — the 2p electron is EASIER to remove than a 2s electron (higher energy, better shielded).
11. Conclusion
The periodic table is not just a chart — it is the ORGANISING PRINCIPLE of chemistry. Understanding the trends in atomic radius, ionisation energy, electron gain enthalpy, and electronegativity allows you to PREDICT chemical behaviour across elements. The block classification (s, p, d, f) connects electronic configuration to position in the table. Mastering these periodic relationships is ESSENTIAL for understanding chemical bonding, reactions, and compound properties.
