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Often in chemistry, students are asked to predict trends in chemical properties. Is compound A more acidic than compound B? Which compound has the strongest covalent bond? Or perhaps, which compound has the highest boiling point? Let's focus on the last question and explore three properties that affect the boiling points of complexes.
Predicting boiling points as well as many other properties is all about understanding chemical trends. Here are three important trends to consider regarding boiling points:
1. The relative strength of the four intermolecular forces
Ionic > Hydrogen Bonding > Dipole-Dipole > Van der Waals
The number one factor when comparing boiling points is considering the strength of the intermolecular forces (IMFs). The stronger the molecules of a liquid are stuck together, the more energy it will take to break them apart so they can enter the gas phase. The functional groups on molecules determine what types of IMFs exist. Pentane, with no functional groups present, has a boiling point of 36ºC. The only IMFs present are the very week Van der Waal interactions. Adding an –OH group to pentane to make pentanol increases the boiling point by over 100ºC because of the added hydrogen bonding interactions. Among molecules with roughly similar molecular weights, the boiling points will be determined by the functional groups present.
2. For molecules with a given functional group, boiling point increases with increasing molecular weight.
The key to this trend is the Van der Waal dispersion forces. These forces are proportional to surface area. As we increase the size or length of a molecular chain we are also increasing the surface area. The increase in surface area increases the ability of individual molecules to attract each other through Van der Waal interactions.
On a non-scientific level, consider a giant plate of spaghetti. The longer the noodles are the harder it is to pull them apart. Increasing the size also increases the chance that two noodles will get stuck to each other. Remember Van der Waals interactions are not strong. It's easy to pull apart two noodles, but that doesn't mean they can't get stuck to each other.
3. Shape In a nutshell, the more cylindrical or rod-like a molecule is, the greater the surface area. We already learned that increasing the surface area also increases the Van der Waal interactions, which, in turn, increases the boiling point. Consider a pile of pencils versus a pile of rocks:
The pencils can stack in a nice ordered pile, with each pencil touching or interacting with another pencil at many different points. Rod-like molecules do the same thing and thus have higher boiling points. The pile of rocks is more random and there are less interactions between rocks. The more spherical a molecule is, the lower its boiling point will be.
Mercury (Hg) is the only metal that is a liquid at room temperature and pressure. What makes mercury so special? It's the only-child of the liquid metal family and it hates sharing electrons.
Image from here.
Most metal atoms are solids because they readily share their valence electrons with surrounding metal atoms. As we discussed before, a metallic crystal is one in which metal atoms sit at the lattice points of the cell and the valence electrons swarm around them, acting as an electronic "glue." This is the reason metals conduct electricity—the shared electrons are free to move around all of the metal atoms.
Mercury is different. Mercury hangs on to its valence 6s electrons very tightly, just like this baby monkey riding on a pig. Because each individual mercury atom has a tighter grip on its own individual valence electrons and does not share these electrons very much with its neighbors, the mercury-mercury bond is very weak. Heat easily overcomes the weak binding and mercury boils and melts at lower temperatures than any other metal. This means it is liquid at room temperature.
Why is mercury so greedy? It turns out that these electrons are able to bury themselves very close to the positive core of the mercury atom. The electrons wiz around the very massive mercury nucleus at speeds close to the speed of light, which causes relativistic effects. We'll leave Einstein's theory of relativity for another day, but basically, these fancy relativistic effects cause the electrons to behave as if they are more massive than electrons moving at slower speeds. Because of the added "fake" mass, the electrons hang out closer to the nucleus.