Question

explain the order of melting points for the solids you tested using what u know about the molecular structures of ionic molecular and covalent network solids.

Answer 1

I don’t know which specific solids you tested, so here’s a clear, general explanation you can use to interpret your results. The key idea is how strongly the particles (ions, molecules, or atoms) are held together — i.e., what you must break when the solid melts.

Summary order (typical):
covalent network solids > ionic solids > molecular solids

Why, in plain terms:

- Covalent network solids
- What holds them together: a continuous 3-D network of strong covalent bonds (every atom bonded to neighbors across the whole solid).
- Consequence: melting requires breaking many strong covalent bonds throughout the structure, so melting points are extremely high.
- Examples: diamond, SiC, SiO2 (quartz).
- Typical behavior: very high melting points, very hard.

- Ionic solids
- What holds them together: strong electrostatic attraction between oppositely charged ions (lattice energy).
- Consequence: you have to overcome these strong but non-directional ionic attractions to melt, so melting points are high — usually lower than covalent networks but much higher than molecular solids.
- Important factors: ion charges and sizes (higher charges and smaller radii → higher lattice energy → higher mp). Example: MgO (very high mp) vs NaCl (moderate high mp).
- Examples: NaCl, KBr, MgO.

- Molecular solids
- What holds them together: relatively weak intermolecular forces (London dispersion, dipole–dipole, hydrogen bonding) between discrete molecules; the covalent bonds within each molecule are not broken on melting.
- Consequence: melting only requires overcoming these weak forces, so melting points are low.
- Variations: molecules with stronger intermolecular forces (polar molecules, H-bonding, large polarizable molecules) have higher melting points than small nonpolar molecules.
- Examples: I2 (sublimes at ~113 °C), dry ice (CO2 sublimes at −78 °C), ice (H2O) — ice is a molecular solid but has relatively high mp among molecular solids because of hydrogen bonding.

Useful details to interpret exceptions and trends:
- Ionic mp depends on lattice energy: multiply-charged ions (Mg2+, Al3+) and small ions → much higher mp than singly charged large ions.
- Molecular mp rises with molecular size (more electrons → stronger London forces) and with stronger specific interactions (H-bonds).
- Some materials can blur categories (e.g., graphite is a covalent network in two dimensions with weak interlayer forces so it has a much lower mp than diamond).
- Always compare the type of bonding you must disrupt on melting: breaking covalent network bonds (very costly) vs. overcoming ionic lattice (very costly but usually less than breaking covalent network) vs. overcoming intermolecular forces (relatively small cost).

If you tell me which solids you actually tested, I can rank them specifically and explain each comparison using their molecular structures and bonding.