Chemistry | Intermolecular Forces

Welcome to Intermolecular Forces!

Hi there! This chapter might seem a little abstract, but it is absolutely fundamental to understanding why different substances—from water to fuels to medicines—have the physical properties they do. Everything you learn here about forces between molecules will explain boiling points, solubility, and even the structure of ice.
Why is this important? In later chapters on Halogenoalkanes and Alcohols, we use intermolecular forces constantly to predict reactivity and explain trends in boiling points. Get this right now, and the organic chemistry will be much easier!

Key Distinction: Intermolecular vs. Intramolecular Forces

Before diving in, let's clear up the biggest source of confusion:

• Intramolecular Forces ("Intra" means inside): These are the strong chemical bonds within a molecule (covalent, ionic, metallic). These forces are broken during a chemical reaction.
• Intermolecular Forces (IMFs) ("Inter" means between): These are the weak forces of attraction between separate molecules. These are the forces we break when melting or boiling a substance (changing state).

Analogy: Think of a brick wall. The intramolecular forces are the cement holding the individual bricks together (very strong). The intermolecular forces are the weak forces holding one wall next to another (easy to separate).

The Three Main Types of Intermolecular Forces (IMFs)

IMFs are generally much weaker than covalent or ionic bonds, but they stack up! We will look at three types, going from weakest to strongest:

1. Van der Waals forces (Temporary Dipole-Induced Dipole)
2. Permanent Dipole-Dipole forces
3. Hydrogen Bonding

1. Van der Waals Forces (VdWs) / London Dispersion Forces

Don't worry, this long name just refers to the weakest, most universal type of force. Every single molecule (polar or non-polar) experiences Van der Waals forces.

The Origin of VdWs: Temporary Dipoles

Electrons are always moving randomly around the nucleus. At any given instant, the electrons might briefly bunch up on one side of the molecule, creating a temporary imbalance:

• This momentary imbalance creates a temporary dipole (a small, temporary positive side and a negative side).
• This temporary dipole can then instantly induce a dipole in a neighboring molecule, causing attraction.
• This attraction is the Van der Waals force.

Analogy: Imagine a crowd of people walking randomly. Occasionally, just by chance, all the people momentarily gather on the right side of the room. This brief crowding is like the temporary dipole. It causes the people in the next room to shift slightly, creating a slight attraction.

Factors Affecting VdWs Strength

The strength of VdWs depends on two main things:

A. Number of Electrons (Size/Relative Molecular Mass, \(M_r\)):

• A larger molecule has more electrons.
• More electrons mean the electron cloud is larger and more polarisable (easier to distort).
• This leads to stronger temporary dipoles, and therefore stronger VdWs.

Example: The boiling points of alkanes increase dramatically as you go from methane \(CH_4\) to octane \(C_8H_{18}\) because the molecules get much bigger, increasing VdWs.

B. Molecular Shape (Surface Area):

• Long, straight molecules (like n-butane) have a large surface area for contact with neighboring molecules. This maximizes the attraction.
• Branched molecules (like iso-butane) are more spherical (compact) and have less surface area for contact.
• Key Takeaway: Increased surface area = Stronger VdWs = Higher boiling point.

2. Permanent Dipole-Dipole Forces

These forces occur only between molecules that are polar (they have a permanent overall dipole).

Prerequisite: Electronegativity and Polarity

A dipole occurs when two atoms in a covalent bond have different electronegativities (the ability of an atom to attract the electron pair in a covalent bond).

• The atom with higher electronegativity (e.g., Chlorine, Oxygen) gains a partial negative charge (\(\delta-\)).
• The atom with lower electronegativity (e.g., Carbon, Hydrogen) gains a partial positive charge (\(\delta+\)).

In a molecule with a permanent dipole (like Hydrogen Chloride, \(HCl\)), the \(\delta+\) end of one molecule is permanently attracted to the \(\delta-\) end of a neighboring molecule.

Crucial Point: Molecules with permanent dipoles have both VdWs forces and permanent dipole-dipole forces. This means they are generally held together more strongly than non-polar molecules of a similar size.

Example Application (Halogenoalkanes): A halogenoalkane like chloroethane \(CH_3CH_2Cl\) is polar because the C-Cl bond is polar. The permanent dipole-dipole forces contribute to its boiling point being higher than that of the non-polar alkane, propane (\(CH_3CH_2CH_3\)), even though they have similar \(M_r\).

3. Hydrogen Bonding (The Special Case)

Hydrogen bonding is the strongest type of IMF. It is so strong that we often give it its own separate category. However, it is fundamentally just a very powerful type of permanent dipole-dipole attraction.

Criteria for Hydrogen Bonding

Hydrogen bonding can only occur when a Hydrogen atom (H) is covalently bonded to one of the three most electronegative, small atoms:

N, O, or F (Nitrogen, Oxygen, or Fluorine)

The resulting bond (\(H-N\), \(H-O\), or \(H-F\)) is highly polarized because these small atoms pull the electrons very strongly, leaving the hydrogen atom with a large \(\delta+\) charge and no electron shielding (because H has no inner electron shells).

This exposed, highly positive H atom is then powerfully attracted to a lone pair of electrons on an N, O, or F atom of a neighboring molecule.

Mnemonic: Think of the NOF club—only molecules with Hydrogen bonded directly to N, O, or F can form H-bonds.

The Impact of H-Bonds:

• Substances capable of H-bonding have drastically higher boiling points than molecules of similar size that only have VdWs or standard dipole-dipole forces.

Example: Water (\(H_2O\))

Water has a very low \(M_r\), but its boiling point is exceptionally high (100 °C) compared to \(H_2S\) (-60 °C) or \(CH_4\) (-161 °C). This is because water molecules form extensive networks of hydrogen bonds, requiring a huge amount of energy to break them apart for boiling.

Example: Alcohols

Alcohols contain the -OH functional group. Because of the H-O bond, they can form hydrogen bonds with each other. This is why small alcohols (like ethanol) have significantly higher boiling points than alkanes or halogenoalkanes of similar \(M_r\).

Linking Intermolecular Forces to Physical Properties

The strength of IMFs determines how easily you can separate the molecules. The stronger the IMFs, the more energy (heat) is required to break them, leading to higher boiling and melting points.

1. Boiling Points (B.P.)

Step-by-Step Analysis for Comparing B.P.

When asked to explain the difference in boiling points between two substances, follow this process:

Step 1: Identify the primary IMFs present in Substance A and Substance B.

Step 2: If one substance can form H-bonds (e.g., Alcohols), it will have the highest B.P. (H-bonds > Dipole-Dipole > VdWs).

Step 3: If neither has H-bonds, compare molecular size (\(M_r\)). If \(M_r\) is very different, VdWs will dominate. Larger \(M_r\) = stronger VdWs = higher B.P.

Step 4: If \(M_r\) is similar, check for polarity. The polar molecule (Dipole-Dipole + VdWs) will have a higher B.P. than the non-polar molecule (VdWs only).

Step 5 (Crucial Conclusion): State that the substance with stronger IMFs requires more energy to overcome these forces during boiling.

Example (Comparing Propane, Chloroethane, and Ethanol):

1. Propane (\(C_3H_8\)): Non-polar. Only VdWs. (Lowest B.P.)
2. Chloroethane (\(C_2H_5Cl\)): Polar. VdWs + Permanent Dipole-Dipole. (Higher B.P.)
3. Ethanol (\(C_2H_5OH\)): Has O-H bond. VdWs + Dipole-Dipole + Hydrogen Bonding. (Highest B.P.)

2. Solubility (Like Dissolves Like)

Solubility is often determined by whether the solute can form the same type of IMFs as the solvent.

• Non-polar solvents (like hexane, \(C_6H_{14}\)): Best dissolve non-polar solutes (which rely only on VdWs).
• Polar solvents (like water, \(H_2O\)): Best dissolve polar or ionic solutes.

Solubility of Alcohols:

Small alcohols (e.g., methanol, ethanol) are miscible (mix completely) with water because they can form H-bonds with the water molecules. They are also soluble in non-polar organic solvents due to their large non-polar hydrocarbon tail.

As the carbon chain of the alcohol increases, the molecule becomes more hydrophobic (water-hating). The long, non-polar alkyl chain dominates, meaning that large alcohols (e.g., hexanol) become poorly soluble in water, as the VdWs forces of the long chain cannot be overcome by the formation of new H-bonds with water.

Solubility of Halogenoalkanes:

Halogenoalkanes are polar, but they generally do not form hydrogen bonds. Therefore, they cannot disrupt the strong network of H-bonds in water. This is why halogenoalkanes are largely insoluble in water, but they mix well with organic (non-polar) solvents.

You have successfully navigated the forces that hold the molecular world together! Understanding these forces is the key to unlocking the Energetics and Organic Chemistry sections. Keep practicing comparison questions, and you'll master this in no time!