Boiling Points of Gasses ```Name: Tina A. Status: educator Age: 30s Location: N/A Country: N/A Date: Sunday, December 08, 2002 ``` Question: I would like to know why oxygen gas has a higher boiling point than nitrogen and fluorine gases. Replies: First, it is not easy to pinpoint a "cause" of why certain substances have a certain boiling point (or melting point). Evaporation is a complicated process. Second, the boiling point is not the correct measure to use if one is comparing volatility. The "normal" boiling point is the temperature at which the vapor pressure is 1 atm (760mm of Hg). This selection of the temperature is arbitrary. You might choose to compare the temperature where the vapor pressure is 100 mm of Hg, or 1000mm of Hg. If you would do so you would find that some of the ordering of these "boiling points" change. So boiling point is not a good characteristic parameter of the volatility. Third, consider the process that is occurring. molecules of liquid have sufficient volatility that a certain number of them can have sufficient kinetic energy to escape the liquid phase -- of course the reverse is also happening and some vapor molecules are striking the surface of the liquid and "sticking". It should be evident that the higher the number of molecules that have this energy the more this process is occurring -- so evaporation has something to do with the vapor density. This would suggest that the critical temperature (the temperature at which the density of the liquid and vapor are equal) would be a better parameter to compare.The critical temperature of N2, O2, and F2 is 126.2, 154.6, and 144.3 (all temperatures in K). Molecular size, shape, volume, mass, and intermolecular forces all play a role in determining the volatility of a substance. The average is 141.7 K, so the critical temperature of all three gases is the same within 10%, and that's about all the sensitivity you can expect. The critical temperature of the elements Ne, Ar, Kr, and Xe is 44.4, 150.8, 209.4 and 289.7. These follow a mass trend, but that may be coincidental, because the atomic size of these molecules (all with the same spherical shape) should also play a role. The atoms have very different polarizability (you can think of this as how "squishy" the atoms are), and this plays an important role in the formation of induced dipoles, which plays a role in the volatility. The bottom line is: It is OK to make these property comparisons to help you develop an intuition about how atoms and molecules behave, but do not take them too seriously, and expect some deviations. Most all the property comparisons you make depend upon the interplay of several parameters, so it is not possible (usually) to say thus and so is the "cause" of the trend. Nonetheless, chemists do it all the time and come up with the right answer or trend for the wrong reason!!! Vince Calder Dear Tina, Rules are meant to be broken!! F2 has more electrons than O2, which has more electrons than N2. If more electrons ALWAYS meant greater dispersion forces then the boiling point ordering would be N2 < O2 < F2. But the REAL property which makes dispersion forces large or small is POLARIZABILITY (this means how easy or hard it for electrons to be pushed around within the molecule by outside forces). Since polarizability often increases with the number of electrons you might guess the above ordering to be correct. However, something else is also changing; the effective nuclear charge is increasing as you go from N2 to O2 to F2, which tends to DECREASE the polarizability. O2 probably has the optimum combination; it has more electrons than N2, but it is more polarizable than F2 (F is the most electronegative element, after all). It is generally safer to make comparisons within the same group (like comparing H2S to H2Se to H2Te) or within a series of otherwise similar compounds (like methane, ethane, propane, butane, pentane...). Then the increasing number of electrons correlates nicely with the polariability, which almost always correlates with the boiling point (but please notice that I said "almost"). prof. topper Click here to return to the Chemistry Archives

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