Metals & Non‑Metals
CBSE Class 10 – Science
(MIT‑Level, First‑Principles View)
This note is written for a motivated student who already knows the basics of chemistry but wants a deeper, rigorous treatment. It blends the standard Class 10 curriculum with concepts from solid‑state physics, quantum chemistry, and materials science.
1. Overview & Classification
| Category | Representative Elements | Typical Uses |
|---|---|---|
| Metals | Fe, Cu, Al, Ag, Pt | Conductors, structural materials |
| Non‑metals | C, H, O, N, P, S | Biomolecules, polymers, fuels |
| Metalloids | Si, Ge, As, Sb | Semiconductors, electronic devices |
Key Idea: The divide is not absolute – many “metalloids” display both metallic and non‑metallic traits.
2. Atomic & Electronic Foundations
2.1 Atomic Structure Recap
- Nucleus: protons + neutrons (mass number (A = Z + N)).
- Electrons: occupy quantum‑mechanical orbitals described by shell (n), angular momentum (l), magnetic (m_l), and spin (s).
2.2 Electron Configuration & Periodic Trend
| Property | Trend Across a Period | Trend Down a Group |
|---|---|---|
| Ionization Energy (IE) | ↑ (except for 3d/4d blocks due to shielding) | ↓ |
| Electron Affinity (EA) | ↓ (most negative at group 17) | ↓ |
| Electronegativity (Pauling) | ↑ | ↓ |
| Metallic Radius | ↓ | ↑ |
Why?
- Shielding: Inner s/p electrons reduce nuclear pull on valence electrons.
- Effective Nuclear Charge (Z_\text{eff}): (Z_\text{eff} = Z - \sigma) (Slater’s rules).
First‑Principles Insight:
The energy required to remove an electron is the difference between the ground‑state energy of the atom and the ion. Calculated via ab initio quantum‑chemical methods (Hartree–Fock, DFT).
2.3 Band Theory – The Bridge Between Atoms & Solids
- Periodic Potential: Electrons in a crystal experience a periodic lattice potential (V(\mathbf{r})).
- Bloch’s Theorem: Wavefunctions (\psi_{\mathbf{k}}(\mathbf{r}) = e^{i\mathbf{k}\cdot\mathbf{r}}u_{\mathbf{k}}(\mathbf{r})).
- Energy Bands: Allowed energies form continuous bands separated by band gaps.
Metals: partially filled valence band → no band gap at (E_F).
Non‑Metals: wide band gap → insulators or semiconductors (if the gap is small).
Metalloids: band gap comparable to thermal energy at room temperature → semiconducting.
Edge Case: Transition metals have partially filled d‑bands, leading to complex band structures and magnetic phenomena (Hund’s rule, crystal field splitting).
3. Physical Properties
| Property | Metals | Non‑Metals |
|---|---|---|
| Electrical Conductivity | High (delocalised electrons) | Low or negligible |
| Thermal Conductivity | High (electron + phonon transport) | Low |
| Malleability & Ductility | High (covalent metallic bonding) | Poor (covalent/ionic bonding) |
| Luster | Metallic shine | Varies (often dull) |
| Melting/Boiling Point | High (strong metallic bonds) | Variable |
| Density | High | Low (often gases) |
| Magnetic Behaviour | Diamagnetic, paramagnetic, ferromagnetic | Diamagnetic |
3.1 Metallic Bonding – Quantum View
- Electron Gas Model: Nuclei fixed in lattice; valence electrons move freely.
- Coulombic Attraction: Provides cohesive force.
- Pauli Exclusion & Fermi–Dirac Statistics: Determine electrical/thermal conductivity.
Rigorous Reasoning:
The band formed from overlapping s/p/d orbitals leads to a continuous density of states. The Fermi level (E_F) lies within this band, allowing electrons to be excited with infinitesimal energy → high conductivity.
3.2 Non‑Metallic Bonding
- Covalent (e.g., C, Si): directional bonding, electrons are shared.
- Ionic (e.g., NaCl‑type non‑metals in alloys): electron transfer, lattice of ions.
- Van der Waals (e.g., noble gases, graphite layers): weak interactions, leading to low melting points.
4. Chemical Properties
4.1 Reactivity Trends
- Metals: Tend to lose electrons → form cations.
- Non‑Metals: Tend to gain or share electrons → form anions or covalent bonds.
Reactivity Series (simplified):
Al > Mg > Ca > Zn > Fe > Sn > Pb > H₂ > Cu > Hg > Ag
Edge Case: Alkali earth metals (Group 2) are less reactive