Historical Development
Dobereiner's Triads (1817)
Elements grouped in triads where atomic mass of middle element was the average of the other two. Example: Li (7), Na (23), K (39). Limited success.
Newlands' Law of Octaves (1864)
Every 8th element had similar properties when arranged by atomic mass. Worked only up to calcium.
Mendeleev's Periodic Table (1869)
Arranged elements by atomic mass. Left gaps for undiscovered elements. Predicted properties of eka-aluminium (Ga), eka-silicon (Ge).
Merits: Predicted new elements, corrected atomic masses. Demerits: Position of isotopes, anomalous pairs (Ar-K, Co-Ni), no place for hydrogen.
Modern Periodic Law
Physical and chemical properties of elements are periodic functions of their atomic numbers.
Long Form Periodic Table
Based on electronic configuration. 18 groups, 7 periods.
Periods
- Period 1: 2 elements (H, He) —
1s - Period 2: 8 elements (Li to Ne) —
2s 2p - Period 3: 8 elements (Na to Ar) —
3s 3p - Period 4: 18 elements (K to Kr) —
4s 3d 4p - Period 5: 18 elements (Rb to Xe) —
5s 4d 5p - Period 6: 32 elements (Cs to Rn) —
6s 4f 5d 6p - Period 7: Incomplete
s, p, d, and f Blocks
s-Block Elements
Groups 1 and 2. Valence electron in s-orbital.
- Group 1: Alkali metals (ns^1). Highly reactive.
- Group 2: Alkaline earth metals (ns^2).
p-Block Elements
Groups 13 to 18. Last electron enters p-orbital.
- Includes metals, non-metals, and metalloids.
- Noble gases in Group 18.
d-Block Elements (Transition Elements)
Groups 3 to 12. Last electron enters d-orbital.
- All are metals.
- Variable oxidation states, coloured compounds, catalytic properties.
f-Block Elements (Inner Transition Elements)
- Lanthanoids (Ce to Lu): 4f subshell filling.
- Actinoids (Th to Lr): 5f subshell filling.
- Radioactive, mostly synthetic.
Periodic Trends
Atomic Radius
Decreases across a period (increased nuclear charge pulls electrons closer). Increases down a group (new shells added).
Covalent radius: Half the distance between nuclei of two covalently bonded atoms. Metallic radius: Half the distance between nuclei in a metallic crystal. Van der Waals radius: Half the distance between non-bonded atoms.
Order: Van der Waals > Metallic > Covalent.
Ionization Enthalpy (IE)
Energy required to remove an electron from a gaseous atom. Increases across a period, decreases down a group.
Successive IEs: IE_1 < IE_2 < IE_3 (removing from increasingly positive species).
Large jump occurs when removing from a stable configuration (e.g., noble gas core).
Electron Gain Enthalpy (Delta_eg H)
Energy change when an electron is added to a gaseous atom.
- Generally negative (energy released).
- Becomes more negative across a period.
- Becomes less negative down a group.
- Noble gases have positive values (very stable).
- Group 17 has most negative values.
Electronegativity
Tendency of an atom to attract shared electrons in a chemical bond.
- Increases across a period, decreases down a group.
- F is the most electronegative element (4.0 on Pauling scale).
- Cs and Fr are the least.
Pauling Scale
Based on bond energies. Difference: X_A - X_B = 0.208 sqrt(Delta), where Delta = E_(AB) - (E_(AA) + E_(BB))/2.
Periodic Trends Summary
| Property | Across Period | Down Group |
|---|---|---|
| Atomic radius | Decreases | Increases |
| Ionization enthalpy | Increases | Decreases |
| Electron gain enthalpy | More negative | Less negative |
| Electronegativity | Increases | Decreases |
| Metallic character | Decreases | Increases |
Worked Examples
Example 1: Arrange O, F, N in increasing order of atomic radius.
Solution: Across period: N > O > F. So F < O < N.
Example 2: Explain why IE_1 of N is higher than O.
Solution: N has half-filled 2p subshell (2p^3), which is more stable. Removing one electron disrupts this stability. O has 2p^4, and removing one gives stable 2p^3 configuration.
Common Mistakes
- Electronegativity vs EA: EN is tendency to attract shared electrons. EA is energy change when adding an electron.
- IE variations: IE of Mg > Al, and N > O (due to subshell stability).
- Anomalous pairs: Ar (atomic mass 40) comes before K (atomic mass 39) in the modern table because Ar has Z=18, K has Z=19.
- Period number: Not equal to number of valence electrons. Period number = highest n value.
ISC Exam Focus
- Theory (70%): Periodic law, block classification, periodic trends, reasons for anomalies.
- Application (30%): Comparing properties of elements, predicting trends.
- ISC frequently asks: "Arrange in order of increasing ..." and "Explain why first IE of Mg is greater than Al."
- Questions on periodicity trends carry 2-4 marks.
Self-Test Questions
Q1: State the modern periodic law. Answer: Properties of elements are periodic functions of their atomic numbers.
Q2: Arrange Na, Mg, Al in increasing order of ionization enthalpy.
Answer: Na < Mg < Al (Na has lowest IE, Al highest among these).
Q3: Why does atomic radius decrease across a period? Answer: Nuclear charge increases while electrons are added to the same shell, pulling the electron cloud inward.
Q4: Write the block, group, and period of Fe (Z=26).
Answer: Fe: [Ar] 3d^6 4s^2. d-block, Group 8, Period 4.
Q5: Which element has the highest electronegativity and why? Answer: Fluorine (4.0). Small size and high nuclear charge attracts electrons strongly.
Q6: Explain why noble gases have positive electron gain enthalpy. Answer: Noble gases have stable ns^2 np^6 configurations. Adding an electron requires energy to place it in a higher energy orbital.
