Q.What are the major differences between metals and non-metals?

The key idea is that metallic character is a periodic trend: it decreases across a period and increases down a group. This trend arises from differences in atomic structure — specifically, how tightly the nucleus holds onto its valence electrons.

Essential reasoning:

1. Electron loss tendency: Metals have low ionisation enthalpy and low electronegativity, so they readily lose electrons to form positive ions (cations). Non-metals have high ionisation enthalpy and high electronegativity, so they gain electrons to form negative ions (anions).

2. Physical properties: Metals are lustrous, malleable, ductile, and good conductors of heat and electricity due to the presence of free (delocalised) electrons in their metallic lattice. Non-metals are typically dull, brittle, and poor conductors (except graphite, a non-metal that conducts electricity). …

The Core Idea: It’s All About Electrons

The entire distinction between metals and non-metals boils down to how tightly an atom holds its outermost electrons. Metals have few valence electrons (typically 1–3) that are loosely held — the nucleus doesn’t pull them strongly. Non-metals have more valence electrons (4–7) that are held much more tightly. This single difference explains every property you’ll ever be asked about.

Think of it this way: a metal atom is like a generous friend who easily lends money (electrons). A non-metal is like a miser who guards every rupee and even tries to take from others.

Step-by-Step Comparison

1. Physical State at Room Temperature

• Metals: Almost all are solids (except mercury, which is liquid). The metallic bond — a "sea" of delocalised electrons holding positive ions together — is strong enough to keep atoms in fixed positions.
• Non-metals: Exist in all three states. Oxygen and nitrogen are gases, bromine is a liquid, carbon and sulphur are solids. Their bonding (covalent or molecular) is weaker, so many are gases.

2. Lustre (Shininess)

• Metals: Have a characteristic metallic lustre — they shine when polished. The free electrons on the surface absorb and re-emit light of all wavelengths, giving a mirror-like finish.
• Non-metals: Generally dull. They lack free electrons, so light is absorbed or scattered irregularly. (Exception: iodine and diamond have some lustre, but this is rare.)

3. Malleability and Ductility

• Metals: Highly malleable (can be hammered into sheets) and ductile (can be drawn into wires). The layers of metal ions can slide over each other without breaking the metallic bond — the electron sea simply flows with them.
• Non-metals: Brittle in solid form. If you hammer sulphur or carbon, they shatter. Their covalent bonds are directional and rigid — they cannot slide without breaking.

4. Electrical and Thermal Conductivity

• Metals: Excellent conductors. The delocalised electrons move freely through the metal, carrying charge (electricity) and transferring kinetic energy (heat).
• Non-metals: Poor conductors (insulators). Electrons are localised in covalent bonds and cannot move. (Exception: graphite conducts electricity because of delocalised electrons in its layers — but it’s a non-metal.)

5. Melting and Boiling Points

• Metals: Generally high melting and boiling points (tungsten melts at 3422°C). The metallic bond requires a lot of energy to break. (Exceptions: gallium and caesium melt in your hand.)
• Non-metals: Generally low melting and boiling points. Weak intermolecular forces (van der Waals) hold molecules together, so little energy is needed. (Exception: diamond and silicon have extremely high melting points due to giant covalent networks.)

6. Chemical Nature: Formation of Ions

• Metals: Tend to lose electrons and form positive ions (cations). Example: $\text{Na} \rightarrow \text{Na}^+ + e^-$. They are electropositive.
• Non-metals: Tend to gain electrons and form negative ions (anions). Example: $\text{Cl} + e^- \rightarrow \text{Cl}^-$. They are electronegative. …