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Boiling points of hydrides of groups iv to vii
Boiling Point (°C)
C H 4 -164
N H 3 -33
H 2 O 100
H F 20
Si H 4 -111.8
P H 3 -87.7
H 2 S -60.7
H Cl -85
Ge H 4 -88.5
As H 3 -55
H 2 Se -41.5
H Br -67
Sn H 4 -52
Sb H 3 -17.1
H 2 Te -2.2
H I -35

In tabular form, there are no obvious trends here, and therefore no obvious connection to the structure orbonding in the molecules. The data in the table are displayed in a suggestive form, however, in , the boiling point of each hydride is plotted according to which period (row) of the periodictable the main group element belongs. For example, the Period 2 hydrides ( C H 4 , N H 3 , H 2 O , and H F ) are grouped in a column to the left of the figure, followed by a column for the Period 3 hydrides( Si H 4 , P H 3 , H 2 S , H Cl ), etc.

Now a few trends are more apparent. First, the lowest boiling points in each period are associated with the GroupIV hydrides ( C H 4 , Si H 4 , Ge H 4 , Sn H 4 ), and the highest boiling points in each period belong to the Group VI hydrides ( H 2 O , H 2 S , H 2 Se , H 2 Te ). For this reason, the hydrides belonging to a single group have been connected in .

Boiling points of main group hydrides

Second, we notice that, with the exceptions of N H 3 , H 2 O , and H F , the boiling points of the hydrides alwaysincrease in a single group as we go down the periodic table: for example, in Group IV, the boiling points increase in the order C H 4 Si H 4 Ge H 4 Sn H 4 . Third, we can also say that the hydrides from Period 2 appear to have unusually high boiling points except for C H 4 , which as noted has the lowest boiling point of all.

We begin our analysis of these trends by assuming that there is a relationship between the boiling points of these compounds and the structure and bonding in their molecules.Recalling our kinetic molecular model of gases and liquids, we recognize that a primary difference between these two phases isthat the strength of the interaction between the molecules in the liquid is much greater than that in the gas, due to the proximityof the molecules in the liquid. In order for a molecule to leave the liquid phase and enter into the gas phase, it must possesssufficient energy to overcome the interactions it has with other molecules in the liquid. Also recalling the kinetic moleculardescription, we recognize that, on average, the energies of molecules increase with increasing temperature. We can concludefrom these two statements that a high boiling point implies that significant energy is required to overcome intermolecularinteractions. Conversely, a substance with a low boiling point must have weak intermolecular interactions, surmountable even at lowtemperature.

In light of these conclusions, we can now look at as directly (though qualitatively) revealing the comparative strengths of intermolecular interactions of the varioushydrides. For example, we can conclude that, amongst the hydrides considered here, the intermolecular interactions are greatestbetween H 2 O molecules and weakest between C H 4 molecules. We examine the three trends in this figure, described above, in light of thestrength of intermolecular forces.

First, the most dominant trend in the boiling points is that, within a single group, the boiling points of thehydrides increase as we move down the periodic table.This is true in all four groups in ; the only exceptions to this trend are N H 3 , H 2 O , and H F . We can conclude that, with notableexceptions, intermolecular interactions increase with increasing atomic number of the central atom in the molecule. This is truewhether the molecules of the group considered have dipole moments (as in Groups V, VI, and VII) or not (as in Group IV). We can inferthat the large intermolecular attractions for molecules with large central atoms arises from the large number of charged particles inthese molecules.

This type of interaction arises from forces referred to as London forces or dispersion forces . These forcesare believed to arise from the instantaneous interactions of the charged particles from one molecule with the charged particles inan adjacent molecule. Although these molecules may not be polar individually, the nuclei in one molecule may attract the electronsin a second molecule, thus inducing an instantaneous dipole in the second molecule. In turn, the second molecule induces a dipole inthe first. Thus, two non-polar molecules can interact as if there were dipole-dipole attractions between them, with positive andnegative charges interacting and attracting. The tendency of a molecule to have an induced dipole is called the polarizability of the molecule. The more charged particles there are in a molecule, the more polarizable a molecule is and the greater the attractions arising from dispersion forces will be.

Second, we note that, without exception, the Group IV hydrides must have the weakest intermolecular interactionsin each period. As noted above, these are the only hydrides that have no dipole moment. Consequently, in general, molecules withoutdipole moments have weaker interactions than molecules which are polar. We must qualify this carefully, however, by noting that thenonpolar Sn H 4 has a higher boiling point than the polar P H 3 and H Cl . We can conclude from these comparisons that the increased polarizability of molecules with heavier atoms can offset the lackof a molecular dipole.

Third, and most importantly, we note that the intermolecular attractions involving N H 3 , H 2 O , and H F must be uniquely and unexpectedly large, since their boiling points aremarkedly out of line with those of the rest of their groups. The common feature of these molecules is that they contain small atomicnumber atoms which are strongly electronegative, which have lone pairs, and which are bonded to hydrogen atoms. Molecules withoutthese features do not have unexpectedly high boiling points. We can deduce from these observations that the hydrogen atoms in eachmolecule are unusually strongly attracted to the lone pair electrons on the strongly electronegative atoms with the sameproperties in other molecules. This intermolecular attraction of a hydrogen atom to an electronegative atom is referred to as hydrogen bonding . It is clear from our boiling point data that hydrogen bonding interactions are much stronger than eitherdispersion forces or dipole-dipole attractions.

Review and discussion questions

Explain the significance to the developmentof the kinetic molecular model of the observation that the ideal gas law works well only at low pressure.

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Explain the significance to the development of the kinetic molecular model of the observation that the pressurepredicted by the ideal gas law is independent of the type of gas.

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Sketch the value of P V n R T as a function of density for two gases, one with strong intermolecular attractionsand one with weak intermolecular attractions but strong repulsions.

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Give a brief molecular explanation for the observation that the pressure of a gas at fixed temperatureincreases proportionally with the density of the gas.

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Give a brief molecular explanation for the observation that the pressure of a gas confined to a fixed volumeincreases proportionally with the temperature of the gas.

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Give a brief molecular explanation for the observation that the volume of a balloon increases roughlyproportionally with the temperature of the gas inside the balloon.

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Explain why there is a correlation between high boiling point and strong deviation from the Ideal Gas Law .

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Referring to , explain why the hydride of the Group 4 element always has the lowest boiling pointin each period.

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Explain why the Period 2 hydrides except C H 4 all have high boiling points, and explain why C H 4 is an exception.

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Source:  OpenStax, General chemistry ii. OpenStax CNX. Mar 25, 2005 Download for free at http://cnx.org/content/col10262/1.2
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