The Boiling Point [EXCLUSIVE]

The boiling point of a substance is the temperature at which it changes state from liquid to gas throughout the bulk of the liquid. At the boiling point molecules anywhere in the liquid may be vaporized.

The Boiling Point

In this post, we will talk about the melting and boiling points of organic compounds and their correlation with intermolecular forces such as dipole-dipole, London dispersion (also known as Van der Waals) interactions, and hydrogen bonding. We discussed these infractions in the previous post and today, the focus will be more from the perspective of physical properties.

We mentioned in the previous post that stronger intermolecular interactions increase the boiling and melting points, but how exactly they affect the physical properties, might be your next question.

As expected, hydrogen bonding affects the physical properties which we can see, for example, by comparing the boiling point of ethanol shown above and dimethyl ether. These two are constitutional isomers meaning they have the same chemical formula and therefore, molecular mass.

Similar behavior can be seen when comparing the boiling point of three isomeric amines: Trimethylamine, Ethylmethylamine, and Propylamine. These are examples of tertiary, secondary, and primary amines which are defined based on the number of alkyl groups connected to the nitrogen.

Notice that the boiling point increases as we are going from the tertiary to the primary amine as a result of increasing the number of hydrogen bonding per nitrogen. Propylamine has two hydrogens connected to the nitrogen and each can make a hydrogen bond with a neighboring nitrogen atom from another molecule. This makes a stronger intermolecular interaction and therefore, more energy is needed to break it pushes the molecules to the gas phase. The tertiary amine, on the other hand, has no hydrogens and boils at a lot lower temperature.

Another factor that influences the boiling point is the surface of the molecule. The larger this surface, the stronger the intermolecular interactions, and thus, the higher the boiling point. This can be seen by comparing the boiling points of pentane, 2-methylbutane, and 2,2-dimethylpropane:

However, the boiling point decreases quite significantly as we move towards the more branched isomers. And this is a demonstration of a direct relationship between the surface area and the boiling point. Pentane is unbranched and provides a large surface for intermolecular interactions. The 2-methylbutane has one substituent so it is a little more branched than pentane. This reduces the surface for intermolecular interactions and lowers the boiling point by about 8 oC. The highly branched 2,2-dimethylpropane, on the other hand, lacks this surface interaction and has the lowest boiling point.

All these examples demonstrate the importance of the molecular surface in intermolecular interactions which directly affect the boiling point of a compound. However, you might have noticed one important thing missing here.

We just said that the molecules are nonpolar and therefore lack dipole-dipole interaction so what type of interactions increase the boiling point with a larger surface?

For example, pentane has a very low melting point compared to butanal since it only relies on London dispersion forces, while butanal contains a polar C=O bond and therefore exhibits dipole-dipole interactions:

Potassium tert-butoxide being an ionic compound has the highest melting point. 1-butanol is the second because of the OH group and thus, hydrogen bonding. Remember, the order of increasing intramolecular interactions in covalent compounds:

Aside from the intermolecular interactions, however, the melting point depends also on how the molecules are packed or arranged in the solid form. The more symmetrical they are, the better they pack and form a perfect crystal lattice which results in a higher melting point. So, as the molecules fit tighter, more energy is required to break the lattice and melt them apart.

Interestingly, the pattern is not observed for the melting points. 2,2-dimethylpropane has a higher melting point () since it is more symmetrical than pentane and when in solid phase (before melting) its molecules are better packed.

We can also see the effect of symmetry by comparing the melting point temperatures of the butanol isomers. All of these are constitutional isomers capable of hydrogen bonding. However, compared to the other isomeric alcohols, the tert-Butyl alcohol has a much higher melting point because of its symmetrical structure and therefore, compact packing in the solid phase:

In contrast, the cis isomer is a polar molecule with a higher boiling point (60 oC vs 48 oC ) because of the net molecular dipole moment and intermolecular dipole-dipole interactions.

Summarizing it, remember that given the same functional groups, the boiling and melting points would naturally be expected to increase with the molecular mass (size) of the molecule. Also, stronger intermolecular interactions presume higher boiling and melting point. However, for the melting point, you need to also consider the factor of symmetry. More symmetry means tighter packing in the solid phases and therefore, a higher melting point.

Pure, crystalline solids have a characteristic melting point, thetemperature at which the solid melts to become a liquid. The transition between the solidand the liquid is so sharp for small samples of a pure substance that melting points canbe measured to 0.1oC. The melting point of solid oxygen, for example, is-218.4oC.

Liquids have a characteristic temperature at which they turn into solids, known astheir freezing point. In theory, the melting point of a solid should bethe same as the freezing point of the liquid. In practice, small differences between thesequantities can be observed.

It is difficult, if not impossible, to heat a solid above its melting point because theheat that enters the solid at its melting point is used to convert the solid into aliquid. It is possible, however, to cool some liquids to temperatures below their freezingpoints without forming a solid. When this is done, the liquid is said to be supercooled.

Because it is difficult to heat solids to temperatures above their melting points, andbecause pure solids tend to melt over a very small temperature range, melting points areoften used to help identify compounds. We can distinguish between the three sugars knownas glucose (MP = 150oC), fructose (MP =103-105oC), and sucrose (MP = 185-186oC), forexample, by determining the melting point of a small sample.

Measurements of the melting point of a solid can also provide information about thepurity of the substance. Pure, crystalline solids melt over a very narrow range oftemperatures, whereas mixtures melt over a broad temperature range. Mixtures also tend tomelt at temperatures below the melting points of the pure solids.

When a liquid is heated, it eventually reaches a temperature at which the vaporpressure is large enough that bubbles form inside the body of the liquid. This temperatureis called the boiling point. Once the liquid starts to boil, thetemperature remains constant until all of the liquid has been converted to a gas.

The normal boiling point of water is 100oC. But if you try to cook an egg inboiling water while camping in the Rocky Mountains at an elevation of 10,000 feet, youwill find that it takes longer for the egg to cook because water boils at only 90oCat this elevation.

In theory, you shouldn't be able to heat a liquid to temperatures above its normalboiling point. Before microwave ovens became popular, however, pressure cookers were usedto decrease the amount of time it took to cook food. In a typical pressure cooker, watercan remain a liquid at temperatures as high as 120oC, and food cooks in aslittle as one-third the normal time.

To explain why water boils at 90oC in the mountains and 120oC ina pressure cooker, even though the normal boiling point of water is 100oC, wehave to understand why a liquid boils. By definition, a liquid boils when the vaporpressure of the gas escaping from the liquid is equal to the pressure exerted on theliquid by its surroundings, as shown in the figure below.

The normal boiling point of water is 100oC because this is the temperatureat which the vapor pressure of water is 760 mmHg, or 1 atm. Under normal conditions, whenthe pressure of the atmosphere is approximately 760 mmHg, water boils at 100oC.At 10,000 feet above sea level, the pressure of the atmosphere is only 526 mmHg. At theseelevations, water boils when its vapor pressure is 526 mmHg, which occurs at a temperatureof 90oC.

Liquids often boil in an uneven fashion, or bump. They tend to bump when therearen't any scratches on the walls of the container where bubbles can form. Bumping iseasily prevented by adding a few boiling chips to the liquid, which provide a roughsurface upon which bubbles can form. When boiling chips are used, essentially all of thebubbles that rise through the solution form on the surface of these chips.

Figuring out the order of boiling points is all about understanding trends. The key thing to consider here is that boiling points reflect the strength of forces between molecules. The more they stick together, the more energy it will take to blast them into the atmosphere as gases.