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This question already has an answer here:

A substance that sublimates doesn't come to a liquid state at room temperature. But when we melt the substance it melts! Why does it skip the liquid state at room temperature?

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marked as duplicate by Mithoron, andselisk, A.K., Todd Minehardt, Nilay Ghosh Feb 28 at 6:25

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  • $\begingroup$ Who said it melts? Some don't. $\endgroup$ – Ivan Neretin Feb 27 at 5:06
  • $\begingroup$ But some do. Why? $\endgroup$ – Arifa Akhtar Feb 27 at 5:11
  • $\begingroup$ Also which don't $\endgroup$ – Arifa Akhtar Feb 27 at 5:11
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    $\begingroup$ Also solids have a vapour pressure. And if you increase the outside pressure, all substances melt before evaporating. $\endgroup$ – Karl Feb 27 at 6:22
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    $\begingroup$ Your Q is also awkward. Does it sublimate or does it melt? Whatever happens before heating the sample is evaporation. As another said solids have vapour pressure, too. $\endgroup$ – Alchimista Feb 27 at 8:57
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From a thermodynamic point of view there are two competing possibilities for condensation of a vapor:

(1) vapor $\ce{->}$ liquid, with some negative $\Delta H$

(2) vapor $\ce{->}$ solid, with an absolutely larger negative $\Delta H$

In the absence of large quantum mechanical effects (see here), the second process with a more negative enthalpy change must become favorable at a low enough temperature above absolute zero; the actual temperature where this happens depends on the material. So when the condensation is done at a low enough temperature, it goes directly from vapor to solid, and the reverse process of evaporation then also misses the liquid.

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    $\begingroup$ Spot on. If the material does not sublime then below the melting point the solid is evaporating, and above the melting point the liquid is evaporating. If material sublimes there is no liquid phase (at that pressure) and the solid will sublime (appears the same as evaporation but no chance of liquification as temperature increases, just faster sublimation). Temperature of solid is held at the Solid/gas phase line by natural cooling due to the evaporation. $\endgroup$ – KalleMP Feb 27 at 19:30
  • $\begingroup$ Sorry, this explains nothing. At any low enough temperature (at given pressure), the material simply either does not evaporate, or is still in thermal equillibrium with it's gas phase. Or it has a melting point. ;-) $\endgroup$ – Karl Feb 27 at 20:46
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Polar or flexible molecules tend to have a liquid phase also at low pressures, so do metals. Rigid, unpolar particles not (naphtalene, helium, iodine). The difference makes wether a not perfectly ordered phase still can have enough intermolecular forces to keep it in a condensed (=liquid) state.

A molten metal is likely still metallic. Hydrogen bonds in liquid water are still very strong. Dodecane is flexible, and methyl end groups can take up a lot of thermal energy without greatly changing the geometry. Etc. Whereas naphtalene is too large to rotate the whole molecule in the condensed phase, but once you have enough thermal energy to break the pi-stacking, there is no other interaction present that could keep the molecule from flying away.

The better the interaction in the liquid phase, the less pressure is needed, resp. the larger the temperature range in which the liquid is the thermodynamically preferred state.

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