Question

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

Answer 1

Short answer: in most cases the melting points follow
covalent network solids >> ionic solids >> molecular (discrete‑molecule) solids.
That ordering comes directly from the types of bonding and the structure of each class.

Why — by class

- Covalent network solids
- Structure: atoms linked by an extended, essentially infinite network of strong directional covalent bonds (examples: diamond, SiC, quartz/SiO2).
- Consequence: to melt you must break many strong covalent bonds throughout the structure, which needs very large input of energy → very high melting points (often >1000 °C).
- Exceptions/notes: layered covalent materials (e.g., graphite) have very strong in‑plane bonds but weak interlayer forces, so their behavior differs; some require high energy to sublimate rather than melt.

- Ionic solids
- Structure: a regular 3D lattice of oppositely charged ions held together by electrostatic (Coulombic) attraction.
- Consequence: melting requires overcoming strong ionic lattice forces (lattice energy). Typical melting points are high (hundreds to over a thousand °C), but lower than for many covalent networks.
- Key factors: lattice energy ∝ (charge product)/(ionic radii). So high charges and small radii (e.g., MgO, CaO) give very high melting points; singly charged ions with larger radii (e.g., KCl) melt at lower temperatures. Ionic solids with highly polarizable ions or lower coordination can have lower MPs.

- Molecular (discrete-molecule) solids
- Structure: separate molecules arranged in a lattice, held together by intermolecular forces (van der Waals/dispersion, dipole–dipole, hydrogen bonding).
- Consequence: because only intermolecular forces must be overcome, melting points are usually low (often <200 °C). The strength of the relevant intermolecular forces controls the MP:
- London dispersion increases with molar mass and polarizability (iodine higher than oxygen).
- Permanent dipole–dipole interactions increase MP for polar molecules.
- Hydrogen bonding (H–bond donors/acceptors) raises MP substantially (e.g., water, ammonia higher than similar-size non‑H‑bonding molecules).
- Result: wide range but generally much lower than ionic or network solids.

Quick examples (typical MPs)
- Diamond ~ 3500 °C (network — extremely high)
- SiO2 (quartz) ~ 1600–1700 °C (network)
- NaCl ~ 801 °C (ionic)
- KCl ~ 770 °C (ionic — similar but lower than NaCl because larger ions)
- MgO ~ 2850 °C (ionic with high charge density → very high)
- Ice 0 °C, iodine ~ 114 °C (sublimes), dry CO2 sublimes at −78 °C (molecular)
- Ammonia melts −78 °C, methanol −98 °C (molecular, H‑bonding changes values)

How to apply this to your measurements
- If a solid you tested had a very high MP, check whether it’s a network covalent or an ionic compound with high lattice energy (high charge/small ions).
- Low MPs indicate molecular solids; compare their molar mass, polarity, and hydrogen bonding ability to explain differences among them.
- Outliers are explained by structure (e.g., layered solids, highly polarizable ions, or strong hydrogen bonding).

If you tell me which specific solids you tested and their measured melting points, I can explain the order of those particular samples in terms of these principles.