Question

explin the order of melting points for the solids you tested using what you know about the molecular and covalent network solid

Answer 1

I can explain the trends you should see. If you want a specific order for the solids you tested, tell me which solids they were and I’ll order them and justify it. Meanwhile, here is a concise framework you can use to explain the melting-point order.

Basic idea
- Melting a solid means overcoming whatever holds its particles in the solid lattice.
- Different classes of solids are held together by very different forces. In general:
- Covalent network solids (diamond, quartz, SiC, etc.) have very high melting points because melting requires breaking a continuous 3‑D network of strong covalent bonds.
- Molecular solids (I2, naphthalene, sucrose, water ice, etc.) have much lower melting points because melting only requires overcoming intermolecular forces between intact molecules (dispersion, dipole–dipole, hydrogen bonding).
- So for a mixed set you will usually find: molecular solids << ionic/metallic << covalent network (with ionic and metallic solids often in the middle depending on lattice/metal bonding strength).

How to order solids within each class
- Molecular solids: order depends on the strength of intermolecular forces
- London dispersion forces increase with molar mass and polarizability → larger, heavier molecules usually have higher melting points.
- Permanent dipole–dipole interactions raise melting points vs. nonpolar molecules of similar size.
- Hydrogen bonding (O–H, N–H) gives the strongest intermolecular attraction among typical molecular solids and raises mp substantially (e.g., water, sugars, amides).
- Crystal packing and symmetry also matter (tighter packing → higher mp).
- Ionic solids: melting point relates to lattice energy (charge magnitude, ionic radii, lattice geometry). Higher charges and smaller ions → larger lattice energy → higher mp.
- Metallic solids: melting point depends on the strength of the metallic bonding (number of delocalized electrons, atomic size, and structure).
- Covalent network solids: melting requires breaking directional covalent bonds across the network → very high mp (often much higher than ionic or metallic solids).

Typical qualitative ordering (mixed example)
- Small nonpolar molecular solids (e.g., O2, N2, CH4) — very low mp
- Larger nonpolar molecular solids (e.g., I2, naphthalene) — higher mp (due to stronger dispersion)
- Polar molecular solids with H-bonding (e.g., H2O ice, sucrose) — higher still
- Ionic solids (e.g., NaCl, MgO) — much higher mp (depends on lattice energy; MgO >> NaCl)
- Covalent network solids (e.g., diamond, SiO2/quartz) — highest mp

Short explanation you can use in a lab report
- “The measured melting points reflect the type and strength of the forces holding the solids together. Molecular solids melted at the lowest temperatures because only intermolecular forces had to be overcome; within molecular solids, heavier molecules and those that hydrogen-bond melted at higher temperatures. Ionic/metallic solids melted at much higher temperatures because strong ionic or metallic bonding must be disrupted. Covalent network solids had the highest melting points because melting requires breaking covalent bonds throughout the extended network.”

If you give me the exact list of solids you tested, I’ll put them in order and provide a sentence-by-sentence justification for each comparison.