Melting Points

Melting Points in Organic Chemistry

The melting point of a compound is the temperature at which the solid phase transitions into the liquid phase at a standard pressure of 1 atmosphere. The melting point of a compound is a physical property, like solubility, density, color, and electronegativity that can be used to identify a compound. Determining the exact temperature at which a compound begins to melt is a challenging task; because of this, the melting point of compounds is reported as a range. The lower limit of the melting point range is the temperature at which the first drops of liquid are observed. The upper limit of the range is the temperature at which all of the solid phase has transitioned to the liquid phase. Reference guides with accepted values exist in the literature, which are used to identify compounds.

The Effect of Intermolecular Forces on Melting Points

One major factor that impacts the melting point of the compound is the type of intermolecular forces that exist within the compound. Intermolecular forces are either attractive or repulsive between the molecules of a compound. In the solid phase, the molecules of a compound will form an organized lattice structure as the molecules are packed close together. There are three major types of intermolecular forces:

1. Hydrogen bonding – Hydrogen bonding is a type of intermolecular force that occurs due to the attraction forces between an electronegative oxygen and a hydrogen atom. Therefore, for this type of intermolecular force to be present, the compound must contain oxygen and hydrogen. Therefore, compounds that contain hydroxyl groups, such as alcohols, readily form hydrogen bonds. Within the hydroxyl group, a dipole forms as the more electronegative oxygen pulls the electron density towards it, making the oxygen have a partial negative charge. This also leaves the hydrogen with a partial positive charge. Nearby electronegative oxygens are attracted to the partial positive charge, forming a hydrogen bond. Of the three types of intermolecular forces, hydrogen bonding is the strongest.
2. Dipole-Dipole Interactions – The second-strongest type of intermolecular force, dipole-dipole interactions forms in molecules that contain electronegative atoms such as oxygen, nitrogen, and any of the halides such as chlorine and fluorine. For example, a hydrocarbon molecule containing fluorine will form dipole-dipole interactions. How? The electronegative fluorine atom will pull the electron density towards it, making it have a partial negative charge. The connecting atom, carbon, loses some of that electron density and thus gains a partial negative charge. This forms a temporary dipole at the fluorine-carbon bond. As opposite charges attract, the partially negative fluorine is attracted to the partially positive carbon of another neighboring molecule, forming a dipole-dipole interaction.
3. London dispersion forces – This type of interaction is a form of van der Waals forces and is present in all compounds. London dispersion forces are the weakest type of intermolecular forces. Like dipole-dipole interactions, there is a redistribution of electron density around the molecule, causing the formation of temporary charges. Unlike dipole-dipole interactions, the dipoles formed in London dispersion forces are very weak and minimal. For example, nonpolar compounds like methane, ethane, pentane, and octane interact via London dispersion forces. The surface area and length of the molecule determines the strength of the attractive forces, such that compounds with more surface area have greater London dispersion forces than smaller compounds. Therefore, octane would have stronger London dispersion forces than methane.

Each type of intermolecular force has a different strength of attraction. Therefore, compounds containing hydrogen bonds require more energy to break the attraction between molecules than a nonpolar compound that only has London dispersion forces. Thus, the presence of hydrogen bonds increases the melting point of a compound.

The Effect of Impurities on Melting Points

Reported literature values of melting points assume that you have a pure sample of the compound in question. Often in the lab or in unknown samples, the samples being tested are not pure compounds. Impurities cause the observed melting point of a mixture to be lower than the actual melting temperature of the pure compound. The observable range is greater than that of the pure substance.

In a pure compound, the solid is composed of a uniform and ordered structure and requires a certain amount of temperature to break the structure apart for the compound to transition into the liquid phase. In a mixture containing impurities, the solid phase is composed of a disorganized structure. This requires much less energy to transition into the liquid phase, thus lowering the melting point. This phenomenon is known as melting point depression. The more impurities in the sample, the broader the melting point range, and the lower the melting temperature.

The melting point of a substance is the temperature at which that substance starts changing from the solid phase to the liquid phase. At this temperature, the liquid and solid phases are in equilibrium. With additional heat, the substance will melt completely. But what determines a substance’s melting point? Let’s think about solids and liquids. A solid’s molecules hold each other in a rigid, ordered structure called a lattice, while a liquid’s molecules have weaker interactions and move around.

Heating a solid transfers energy to the molecules. With enough energy, the molecules overcome the forces keeping them in the lattice and start moving around. In other words, if we heat a solid enough, it melts into a liquid. So, the melting point depends on the energy it takes to overcome the forces between the molecules, or the intermolecular forces, holding them in the lattice. The stronger the intermolecular forces are, the more energy is required, so the higher the melting point is.

Many intermolecular forces depend on how strongly atoms in the molecule attract electrons — or their electronegativity. Nitrogen, oxygen, fluorine, and chlorine are highly electronegative, while carbon, hydrogen, and sulfur are only moderately electronegative. Bonds between atoms with significantly different electronegativities are polar. For instance, a typical carbon-oxygen bond is polar, but a typical carbon-hydrogen bond is not.

A molecule’s electrons spend more time around its most electronegative atoms, giving it a slight negative charge on that side and a slight positive charge on the other side. This is called a dipole. If the dipole isn’t canceled out by an equal and opposite dipole in the same molecule, the molecule has a permanent dipole and is polar.

Now, let’s discuss three important intermolecular forces: hydrogen bonding, dipole-dipole interactions, and London dispersion forces. Hydrogen bonding occurs between an electron-withdrawing atom with a lone pair of electrons and a hydrogen bound to a more electronegative atom. Hydrogen bonds are among the strongest intermolecular forces.

Dipole-dipole interactions occur between polar molecules. In an attractive dipole-dipole interaction, the negative side of one dipole aligns with the positive side of another dipole. Dipole-dipole interactions are generally weaker than hydrogen bonds.

London dispersion forces come from brief, random shifts in a molecule’s electron distribution, which cause corresponding shifts in nearby molecules. These random shifts happen in every molecule, so this is one of the few interactions available to nonpolar molecules. London dispersion forces are among the weakest intermolecular forces.

Earlier, we predicted that stronger intermolecular forces corresponded to higher melting points. We can see this in action with hexadecane, 2-hexadecanone, and hexadecanoic acid. As the strength of the intermolecular interactions available to each molecule increases, so does the melting point.

Intermolecular forces aren’t the only factor that determines the melting point of a substance. Its purity significantly affects its melting and freezing points in an effect called ‘freezing-point depression’. This effect means that a solution has a lower freezing point than the pure solvent does. That’s why streets are sprinkled with salt when it gets very cold. If any water collects on the street, the salt quickly dissolves to make a solution with a much lower freezing point than pure water.

In a solid, impurities are incorporated into the lattice structure. These areas often have weaker intermolecular interactions, making parts of the structure easier to disrupt. So, compared to a pure solid, melting starts at a lower temperature and occurs over a wider temperature range.

In this lab, you will measure the melting points of two known organic compounds and then analyze a mixture to explore how impurities affect the melting point range.