Phase Changes and Intermolecular Forces

Greetings,

This post explores the nature of phase changes a bit more deeply by explaining intermolecular forces as they relate to the more dense phases; liquid and solid.  The forces involved within a solid molecular substance are the strongest of the three states of matter as the molecules are essentially locked in place within a crystalline (usually) structure.  The molecules within a liquid are free to move about, but there are still significant intermolecular forces at work within this state of matter.  However, the intermolecular forces of a liquid are not strong enough to keep molecules in place long enough for a crystalline structure to develop.

Types of Intermolecular Attractions

The types of attractions holding molecules together depend on the existence of molecules with permanent dipoles.  The existence of permanent molecular dipoles is in turn dependent on the particular geometry of the molecule.  A molecule with an asymmetric structure, such as water (bent) and ammonia (trigonal pyramidal), will acquire an overall partial separation of electron charge density across the length of the structure.  We say that such a structure has a dipole moment and is deemed to be a polar molecule.  A molecule with perfect symmetry, such as carbon tetrachloride, cannot have an overall separation of electron density across its structure and therefore cannot develop a permanent dipole moment: We call that molecule nonpolar.  But, in terms of a statistical analysis, there is a slight probability that, at any given moment, there will exist a partial separation of electron density within atoms of a group of nonpolar molecules.  This results in relatively weak and very short-lived regions of polarity within nonpolar molecules which provides enough overall intermolecular attractions for a nonpolar compound to exist as a liquid within a certain temperature range (usually much lower than the temperatures involved with polar molecular compounds).  This attraction between nonpolar molecules, just described, is called london-dispersion forces.

The diagram below provides a summary of this discussion.

Notice the hexagonal arrangement of water molecules in ice.  The hexagonal geometry is duplicated until we can literally observe it; as a six sided snowflake!  The structure shown for carbon dioxide is a unit cell of dry ice.  The geometry of the dry ice crystal is face-centered-cubic.  The view of ethylene glycol molecules shows dipole-dipole (H-bonding) interactions between hydroxyl groups.  Overall, the collective strength of H-bonding in ethylene glycol is quite high, which results in the compound's very high boiling point (197.3 deg C)!  The london-dispersion forces, which dominant for carbon tetrachloride, are indicated by the negative and positive charge symbols dispersed across interacting chlorine atoms.

That's all for now.
Have a good one!


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