1. Why is Classification of Elements Necessary?
- More than 100 elements are known.
- Studying each element separately is difficult.
- Classification helps:
- Understand similarities and differences.
- Predict properties of unknown elements.
- Make learning easier.
2. Development of the Periodic Table
(A) Dobereiner’s Triads (1829)
Definition
Groups of three elements having similar properties.
Features
- Atomic mass of middle element ≈ average of the other two.
- Chemical properties were similar.
Examples
| Triad | Elements |
|---|---|
| Alkali metals | Li, Na, K |
| Alkaline earth metals | Ca, Sr, Ba |
| Halogens | Cl, Br, I |
Limitation
Worked only for a few elements.
(B) Newlands’ Law of Octaves (1865)
Statement
When elements are arranged in increasing atomic mass, every eighth element has similar properties.
Limitation
- Valid only up to Calcium.
- Failed for heavier elements.
(C) Mendeleev’s Periodic Table
Periodic Law
The physical and chemical properties of elements are periodic functions of their atomic masses.
Merits
✔ Classified most elements correctly.
✔ Left gaps for undiscovered elements.
✔ Predicted properties of unknown elements like:
- Gallium
- Germanium
✔ Corrected doubtful atomic masses.
Limitations
- Position of Hydrogen uncertain.
- Isotopes could not be explained.
- Increasing atomic mass order was violated in some cases.
3. Modern Periodic Law
Given by
Henry Moseley (1913)
Statement
The physical and chemical properties of elements are periodic functions of their atomic numbers.
Importance
Atomic number is the basis of the modern periodic table.
4. Modern Periodic Table
- 7 Periods (horizontal rows)
- 18 Groups (vertical columns)
Number of elements in each period
| Period | Elements |
|---|---|
| 1st | 2 |
| 2nd | 8 |
| 3rd | 8 |
| 4th | 18 |
| 5th | 18 |
| 6th | 32 |
| 7th | Incomplete (up to 32 possible) |
5. IUPAC Naming of Elements
Temporary names are formed using number roots.
| Number | Root |
|---|---|
| 0 | Nil |
| 1 | Un |
| 2 | Bi |
| 3 | Tri |
| 4 | Quad |
| 5 | Pent |
| 6 | Hex |
| 7 | Sept |
| 8 | Oct |
| 9 | Enn |
Example
Atomic number 120
Name → Unbinilium
Symbol → Ubn
6. Electronic Configuration and Periodic Table
Period Number
Represents the highest principal quantum number (n).
Example
Na → 3rd Period
Cl → 3rd Period
Group Number
Determined by valence electrons.
Elements in the same group have:
- Similar electronic configuration
- Similar chemical properties
7. Blocks of the Periodic Table
(A) s-block
Groups: 1 and 2
Outer configuration:
ns¹ – ns²
Properties
- Highly reactive metals
- Electropositive
- Form ionic compounds
Examples
Li, Na, Mg, Ca
(B) p-block
Groups: 13–18
Outer configuration
ns² np¹–np⁶
Properties
- Metals
- Non-metals
- Metalloids
Contains:
- Halogens
- Noble gases
(C) d-block
Groups: 3–12
Outer configuration
(n−1)d¹–¹⁰ ns⁰–²
Also called
Transition Elements
Properties
- Variable oxidation state
- Coloured compounds
- Good catalysts
(D) f-block
Contains
- Lanthanoids
- Actinoids
Properties
- Inner transition elements
- Mostly radioactive (Actinoids)
8. Metals, Non-metals and Metalloids
Metals
Properties
- Good conductor
- Malleable
- Ductile
- Lustrous
- Lose electrons
Examples
Na, Fe, Cu
Non-metals
Properties
- Poor conductor
- Brittle
- Gain electrons
Examples
O, Cl, N
Metalloids
Show properties of both metals and non-metals.
Examples
- Silicon
- Germanium
- Arsenic
- Antimony
- Tellurium
9. Periodic Trends
(A) Atomic Radius
Across a Period
➡ Decreases
Reason
Effective nuclear charge increases.
Down a Group
⬇ Increases
Reason
New electron shells are added.
(B) Ionic Radius
Cations
Smaller than parent atom.
Example
Na > Na⁺
Anions
Larger than parent atom.
Example
F⁻ > F
10. Ionization Enthalpy
Definition
Energy required to remove one electron from an isolated gaseous atom.
Trend
Across a period
⬆ Increases
Down a group
⬇ Decreases
Reason
- Increasing nuclear charge across a period.
- Increasing atomic size down a group.
11. Electron Gain Enthalpy
Definition
Energy change when an electron is added to an atom.
Trend
Across a period
Becomes more negative.
Down a group
Becomes less negative.
Highest negative value
Chlorine
12. Electronegativity
Definition
Ability of an atom to attract shared electrons.
Highest
Fluorine (4.0)
Lowest
Cesium / Francium
Trend
Across a period
⬆ Increases
Down a group
⬇ Decreases
13. Metallic Character
Across a Period
⬇ Decreases
Down a Group
⬆ Increases
14. Non-metallic Character
Across a Period
⬆ Increases
Down a Group
⬇ Decreases
15. Oxidation State
For representative elements
| Group | Common Oxidation State |
|---|---|
| 1 | +1 |
| 2 | +2 |
| 13 | +3 |
| 14 | ±4 |
| 15 | -3, +3, +5 |
| 16 | -2, +4, +6 |
| 17 | -1, +1, +3, +5, +7 |
| 18 | 0 |
16. Anomalous Behaviour of Second Period Elements
Elements
Li, Be, B, C, N, O, F
Reasons
- Small size
- High charge density
- High electronegativity
- Absence of d-orbitals
17. Chemical Reactivity
Alkali Metals
Reactivity increases down the group.
Example
Li < Na < K < Rb < Cs
Halogens
Reactivity decreases down the group.
Example
F > Cl > Br > I
Important NCERT Definitions
Modern Periodic Law: The physical and chemical properties of elements are periodic functions of their atomic numbers.
Ionization Enthalpy: Energy required to remove the outermost electron from an isolated gaseous atom.
Electron Gain Enthalpy: Energy change when an electron is added to a gaseous atom.
Electronegativity: Ability of an atom to attract the shared pair of electrons.
One-Line Revision
- Modern periodic table → Based on Atomic Number
- Periods → 7
- Groups → 18
- s-block → Groups 1–2
- p-block → Groups 13–18
- d-block → Groups 3–12
- f-block → Lanthanoids & Actinoids
- Atomic radius → ↓ across, ↑ down
- Ionization enthalpy → ↑ across, ↓ down
- Electronegativity → ↑ across, ↓ down
- Metallic character → ↓ across, ↑ down
- Non-metallic character → ↑ across, ↓ down
- Most electronegative element → Fluorine
- Largest atomic size → Francium (approx.)
- Most reactive metal → Francium/Cesium
- Most reactive non-metal → Fluorine