When discussing the intermolecular forces between methanol molecules, chemists usually ignore any London forces between them. Why are they justified in doing this?

There are three main types of intermolecular forces (IMFs) that can occur between molecules: 1. London dispersion forces (also known as van der Waals forces or London forces) - these are temporary attractive forces that occur due to fluctuations in electron density. These forces exist between all molecules, regardless of polarity, and are typically the weakest of the IMFs. 2. Dipole-dipole forces - these forces occur when the positive end of one polar molecule is attracted to the negative end of another polar molecule. These forces are stronger than London dispersion forces. 3. Hydrogen bonding - this is a specific type of dipole-dipole force that occurs when a hydrogen atom is bonded to a highly electronegative atom (usually nitrogen, oxygen, or fluorine) and interacts with an electronegative atom on another molecule. Hydrogen bonding is the strongest among these three types of IMFs.

First, let's identify the molecular structure of methanol. Methanol has the chemical formula CH3OH and consists of a methyl group (CH3) bonded to a hydroxyl group (OH). Due to the electronegativity difference between hydrogen and oxygen atoms in the hydroxyl group, methanol is a polar molecule. Therefore, dipole-dipole forces will be present between methanol molecules. Additionally, because there is a hydrogen atom bonded directly to an oxygen atom (which is an electronegative atom), hydrogen bonding can also occur between methanol molecules. As previously mentioned, London forces are present between all molecules. However, we need to determine whether it is justified to ignore these forces when discussing methanol molecules' intermolecular forces.

When comparing the strength of intermolecular forces, hydrogen bonding is the strongest, followed by dipole-dipole forces, and London forces being the weakest. In the case of methanol, hydrogen bonding and dipole-dipole forces are significantly stronger than the London forces. The presence of the polar hydroxyl group and the hydrogen bonding between methanol molecules makes these forces more dominant in determining methanol's physical and chemical properties.

London forces are weaker than both dipole-dipole forces and hydrogen bonding in methanol molecules. Due to this significant difference in strength, chemists often choose to disregard London forces when discussing methanol's intermolecular forces. Focusing on the stronger intermolecular forces (hydrogen bonding and dipole-dipole forces) provides a more accurate representation of the interactions between methanol molecules and better explains methanol's properties compared to only considering the much weaker London forces.