Short Answer Questions — Chapter 8: Periodic Classification of Elements
Mendeleev's Periodic Law (1869): The physical and chemical properties of elements are periodic functions of their atomic masses. When elements are arranged in order of increasing atomic mass, elements with similar properties recur at regular intervals.
Merits of Mendeleev's Periodic Table:
1. Arrangement of all known elements: Successfully arranged all 63 elements known at that time.
2. Prediction of undiscovered elements: Left gaps in the table and predicted properties of elements yet to be discovered:
All showed properties very close to Mendeleev's predictions.
- Eka-boron → later discovered as Scandium (Sc)
- Eka-aluminium → Gallium (Ga)
- Eka-silicon → Germanium (Ge)
3. Correction of atomic masses: Revised some atomic masses (e.g., Be, In, La) to fit the periodic table, which were later confirmed.
4. Foundation for modern chemistry: The basis for the modern periodic table.
1. Position of hydrogen: Hydrogen resembles both alkali metals (Group 1) and halogens (Group 17) but fits perfectly in neither.
2. Anomalous pairs (wrong order): Some elements placed in wrong order of atomic mass to maintain property grouping:
- Ar (40) before K (39): Ar should come after K by mass, but placed before
- Co (59) before Ni (58.7)
- Te (128) before I (127)
3. Position of isotopes: Isotopes of same element have different masses. By Mendeleev's law, they should occupy different positions. No place was given to isotopes.
4. Position of lanthanides and actinides: 14 lanthanides (rare earth elements) were grouped together as one position — not ideal.
5. Cannot explain why periodicity exists — no explanation in terms of atomic structure.
Modern Periodic Law (Henry Moseley, 1913): The physical and chemical properties of elements are periodic functions of their atomic numbers (number of protons).
How it improved on Mendeleev:
| Problem | Mendeleev's table | Modern table |
|---|---|---|
| Anomalous pairs (Ar-K, Co-Ni, Te-I) | Wrong order | Correct order when arranged by atomic number |
| Isotopes | No place | Same atomic number → same position (same element) |
| Why periodicity? | No explanation | Electron configuration explains it |
| New elements | Some gaps | Can predict by atomic number |
Electronic basis of modern periodic law:
- Period number = number of electron shells
- Group number relates to number of valence electrons
- Elements in same group have same valence electron configuration → same chemical properties
Atomic radius: Distance from nucleus to outermost electron.
Across a period (left to right):
- Atomic number increases → nuclear charge increases
- Number of electron shells remains the same
- Higher nuclear charge pulls electrons closer to nucleus
- Atomic radius decreases across a period
Example (Period 3): Na > Mg > Al > Si > P > S > Cl > Ar
Down a group (top to bottom):
- Atomic number increases → one new shell added each time
- New outer shell is farther from nucleus
- Shielding effect of inner shells reduces effective nuclear charge on outer electrons
- Atomic radius increases down a group
Example (Group 1): Li < Na < K < Rb < Cs
| Property | Metals | Non-metals | Metalloids |
|---|---|---|---|
| Position in table | Left side | Right side | Staircase boundary |
| Physical state | Mostly solid | Mostly solid/gas | Solid |
| Luster | Shiny | Dull | Semi-lustrous |
| Conductivity | Good conductors | Poor conductors | Semi-conductors |
| Malleability | Malleable | Brittle | Brittle |
| Electronegativity | Low | High | Intermediate |
| Ion formation | Form positive ions (cations) | Form negative ions (anions) | Both |
| Oxides | Basic (Na₂O, MgO) | Acidic (CO₂, SO₂) | Amphoteric (SiO₂) |
| Examples | Na, Fe, Cu, Al | C, N, O, Cl, S | B, Si, Ge, As, Sb |
Metalloids (also called semi-metals) show properties of both metals and non-metals. They are very important as semiconductors in electronics (silicon, germanium).
Ionisation Energy (IE): The minimum energy required to remove an electron from the outermost orbit of a neutral gaseous atom.
M(g) → M⁺(g) + e⁻ ; energy needed = first ionisation energy
Unit: kJ/mol or eV
Trends:
Across a period (left to right):
- Atomic radius decreases → electrons held more tightly
- Nuclear charge increases
- Ionisation energy increases across a period
- Exceptions: slight dips at Group 13 (e.g., Al < Mg) and Group 16 (e.g., S < P) due to electron pairing and orbital effects
Down a group (top to bottom):
- Atomic radius increases → outer electrons farther from nucleus
- Shielding effect increases
- Ionisation energy decreases down a group
Why noble gases have highest IE: Completely filled shells — very stable, hard to remove electrons.
Why alkali metals have lowest IE: 1 valence electron, far from nucleus — easily removed.
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