Question

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

Short answer

Chemists usually ignore London forces between methanol molecules because they are significantly weaker than the other intermolecular forces present in methanol, namely hydrogen bonding and dipole-dipole forces. The polar hydroxyl group and hydrogen bonding between methanol molecules make these stronger forces more dominant in determining methanol's physical and chemical properties, and thus London forces are disregarded for a more accurate representation of the interactions between the molecules.

Step-by-step solution

Identify the types of intermolecular forces

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.

Identify the intermolecular forces present in methanol

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.

Compare the strength of London forces to dipole-dipole forces and hydrogen bonding in methanol

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.

Explain why London forces are ignored in methanol

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.

Key concepts

London Dispersion Forces

London dispersion forces, also known as van der Waals forces, are one type of intermolecular force that exist between all molecules. These forces occur because of temporary shifts in electron density that create moments of induced polarity.

Even though they are always present, London dispersion forces are usually the weakest of the intermolecular forces. They are important when considering non-polar molecules, as they can be the only types of interaction available. These forces become stronger when:

• Molecules have larger electron clouds, which allows for greater fluctuations in electron density.
• Molecules are larger and heavier, as this increases their surface area for interactions.

In the context of methanol, while London forces do exist, they are less significant because methanol has much stronger interactions due to other types of intermolecular forces, specifically hydrogen bonding and dipole-dipole forces.

Hydrogen Bonding

Hydrogen bonding is a special type of dipole-dipole force, and it's particularly strong compared to London dispersion forces. This bond occurs when a hydrogen atom is directly bonded to a highly electronegative atom like nitrogen, oxygen, or fluorine, and then it is attracted to another electronegative atom nearby.

In methanol, the hydroxyl group (OH) facilitates hydrogen bonding. The hydrogen atom of one methanol molecule aims to bond with the oxygen atom of another, creating a strong and specific attraction.

• Hydrogen bonds lead to higher boiling and melting points because of the extra energy needed to break these bonds during phase changes.
• They are crucial for determining the structure and properties of water and many biological molecules, like proteins and DNA.

Because of its impactful role, hydrogen bonding is often the primary focus when analyzing methanol's properties in comparison to other intermolecular forces.

Dipole-Dipole Forces

Dipole-dipole forces come into play between molecules that have permanent dipoles—situations where there is an uneven charge distribution within the molecule. These occur in polar molecules where parts of the molecule have partial positive and negative charges.

Methanol is a polar molecule due to its molecular structure, particularly because of the electronegativity difference between the oxygen and hydrogen in the hydroxyl group. This leads to permanent dipole-dipole interactions.

• These forces are stronger than London dispersion forces but typically not as strong as hydrogen bonds.
• They significantly impact the boiling and melting points of substances by creating additional resistance to breaking apart.

In methanol, dipole-dipole interactions add to the stability and structure of the liquid, but they are still considered less influential compared to the pervasive hydrogen bonding that dominates its intermolecular interactions.