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Comparison of the physical properties of alcohols with those of hydrocarbons of comparable molecular weight shows several striking differences, especially for those with just a few carbons. Alcohols are substantially less volatile, have higher melting points, and greater water solubility than the corresponding hydrocarbons (see Table 15-1), although the differences become progressively smaller as molecular weight increases.

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Table 15-1: Comparison of Physical Properties of Alcohols and Hydrocarbons

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The reason for these differences in physical properties is related to the high polarity of the hydroxyl group which, when substituted on a hydrocarbon chain, confers a measure of polar character to the molecule. As a result, there is a significant attraction of one molecule for another that is particularly pronounced in the solid and liquid states. This polar character leads to association of alcohol molecules through the rather positive hydrogen of one hydroxyl group with a correspondingly negative oxygen of another hydroxyl group:

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This type of association is called “hydrogen bonding,” and, although the strengths of such bonds are much less than those of most conventional les-grizzlys-catalans.orgical bonds, they are still significant (about (5) to (10 : extkcal) per bond). Clearly then, the reason alcohols have higher boiling points than corresponding alkyl halides, ethers, or hydrocarbons is because, for the molecules to vaporize, additional energy is required to break the hydrogen bonds. Alternatively, association through hydrogen bonds may be regarded as effectively raising the molecular weight, thereby reducing volatility (also see Section 1-3).

The water solubility of the lower-molecular-weight alcohols is pronounced and is understood readily as the result of hydrogen bonding with water molecules:

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In methanol, the hydroxyl group accounts for almost half of the weight of the molecule, and it is not surprising that the substance is completely soluble in water. As the size of the hydrocarbon groups of alcohols increases, the hydroxyl group accounts for progressively less of the molecular weight, hence water solubility decreases (Figure 15-1). Indeed, the physical properties of higher-molecular-weight alcohols are very similar to those of the corresponding hydrocarbons (Table 15-1). The importance of hydrogen bonding in the solvation of ions was discussed in Section 8-7F.

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The chart below shows the boiling points of the following simple primary alcohols with up to 4 carbon atoms:

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These boiling points are compared with those of the equivalent alkanes (methane to butane) with the same number of carbon atoms.

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Notice that:

The boiling point of an alcohol is always significantly higher than that of the analogous alkane. The boiling points of the alcohols increase as the number of carbon atoms increases.

The patterns in boiling point reflect the patterns in intermolecular attractions.



Hydrogen bonding

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Hydrogen bonding occurs between molecules in which a hydrogen atom is attached to a strongly electronegative element: fluorine, oxygen or nitrogen. In the case of alcohols, hydrogen bonds occur between the partially-positive hydrogen atoms and lone pairs on oxygen atoms of other molecules.

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The hydrogen atoms are slightly positive because the bonding electrons are pulled toward the very electronegative oxygen atoms. In alkanes, the only intermolecular forces are van der Waals dispersion forces. Hydrogen bonds are much stronger than these, and therefore it takes more energy to separate alcohol molecules than it does to separate alkane molecules. This the main reason for higher boiling points in alcohols.


The effect of van der Waals forces

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Boiling points of the alcohols: Hydrogen bonding is not the only intermolecular force alcohols experience. There are also van der Waals dispersion forces and dipole-dipole interactions. The hydrogen bonding and dipole-dipole interactions are much the same for all alcohols, but dispersion forces increase as the alcohols get bigger. These attractions get stronger as the molecules get longer and have more electrons. This increases the sizes of the temporary dipoles formed. This is why the boiling points increase as the number of carbon atoms in the chains increases. It takes more energy to overcome the dispersion forces, and thus the boiling points rise. Comparison between alkanes and alcohols: Even without any hydrogen bonding or dipole-dipole interactions, the boiling point of the alcohol would be higher than the corresponding alkane with the same number of carbon atoms.

Compare ethane and ethanol:

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Small alcohols are completely soluble in water; mixing the two in any proportion generates a single solution. However, solubility decreases as the length of the hydrocarbon chain in the alcohol increases. At four carbon atoms and beyond, the decrease in solubility is noticeable; a two-layered substance may appear in a test tube when the two are mixed.

Consider ethanol as a typical small alcohol. In both pure water and pure ethanol the main intermolecular attractions are hydrogen bonds.

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In order to mix the two, the hydrogen bonds between water molecules and the hydrogen bonds between ethanol molecules must be broken. Energy is required for both of these processes. However, when the molecules are mixed, new hydrogen bonds are formed between water molecules and ethanol molecules.

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The energy released when these new hydrogen bonds form approximately compensates for the energy needed to break the original interactions. In addition, there is an increase in the disorder of the system, an increase in entropy. This is another factor in deciding whether les-grizzlys-catalans.orgical processes occur. Consider a hypothetical situation involving 5-carbon alcohol molecules.

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The hydrocarbon chains are forced between water molecules, breaking hydrogen bonds between those water molecules. The -OH ends of the alcohol molecules can form new hydrogen bonds with water molecules, but the hydrocarbon "tail" does not form hydrogen bonds. This means that many of the original hydrogen bonds being broken are never replaced by new ones.

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In place of those original hydrogen bonds are merely van der Waals dispersion forces between the water and the hydrocarbon "tails." These attractions are much weaker, and unable to furnish enough energy to compensate for the broken hydrogen bonds. Even allowing for the increase in disorder, the process becomes less feasible. As the length of the alcohol increases, this situation becomes more pronounced, and thus the solubility decreases.


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