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Name: Amy S.
Status: educator
Age: 30s
Location: N/A
Country: N/A
Date: Sunday, December 08, 2002

Based on intermolecular forces, carbon monoxide should boil at a higher temperature than methane. However, it does not. Why?

Carbon monoxide does have a dipole moment but it is very small. C and O are similar in electronegativity so the charge difference is quite small.

Methane has no dipole moment but has more electrons than CO does and is more polarizable.

So even though CO molecules exhibit dipole-dipole and dispersion forces and CH4 molecules only exhibit dispersion forces, CH4's dispersion forces are apparently attractive enough to win the game.

A more extreme example can be seen if you compare n-octane, a nonpolar hydrocarbon, with water, which exhibits hydrogen bonding, dipole-dipole and dispersion forces. However, water's boiling point is 100 deg.C while octane's is 125.7 deg.C . Dispersion forces can be quite strong, and the bigger the molecule the stronger the dispersion force.

R. Topper

Correlating boiling points (or melting points) with "intermolecular forces" is only approximate. Both boiling and melting are much more complicated phenomena. In addition, what experimental data was used to determine the "intermolecular forces" can change both the absolute and relative values assigned to a given molecule. The molecular volume and shape also play a subtle but important role. Remember that the boiling point is rather arbitrary. It is the temperature at which the vapor pressure equals 1 atmosphere. Choose another vapor pressure and the order of the list of boiling points may change. The value of the critical constants is frequently a better parameter to compare. The critical constants Tc, Pc, and Vc are the value of temperature, pressure, and molar volume where the density of the liquid equals the density of the vapor. In fact, the parameter Zc = PcVc/RTc is often used to compare the aggregate effect of the various molecular factors that measure things like volatility. However, the critical constants are known for only a small number of molecules. The bottom line is: Do not take correlations too seriously; they are only estimates (read that "guesses").

Vince Calder

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