In the realm of chemical interactions, while covalent bonds hold atoms together within a molecule, it's the intermolecular forces that determine how these molecules interact with one another.
Classification of Intermolecular Forces
Intermolecular forces can be broadly classified into three main types based on their strength and the nature of interacting particles:
1. Van der Waals' Forces (or London Dispersion Forces)
- Definition: These are the weakest type of intermolecular forces and arise due to instantaneous dipoles in molecules.
- Occurrence: Present in all molecules, regardless of their polarity.
- Cause: Fluctuations in the electron cloud around a molecule lead to temporary dipoles, which induce dipoles in adjacent molecules.
The formation of temporary dipoles induces dipoles in adjacent molecules.
Image courtesy of OpenStax
2. Dipole-Dipole Interactions
- Definition: These forces arise between molecules that have permanent dipoles.
- Occurrence: Common in polar molecules.
- Cause: Positive end of one polar molecule attracts the negative end of another, leading to attractions.
Image courtesy of Adam Rędzikowski
3. Hydrogen Bonding
- Definition: A special case of dipole-dipole interactions where a hydrogen atom is bonded to a highly electronegative atom and is attracted to another electronegative atom.
- Occurrence: Commonly observed in molecules where H is bonded to N, O, or F.
Deducing Types of Intermolecular Forces from Molecular Structures
To determine which intermolecular forces are present:
1. Analyse the Molecule for Polarity:
- If the molecule is non-polar and has no hydrogen bonded to N, O, or F, then only van der Waals' forces are present.
- If the molecule is polar, it has dipole-dipole interactions.
- If the molecule has a H bonded to N, O, or F, then hydrogen bonding may also be present.
2. Check for Hydrogen Bonding Capabilities:
- Not all molecules with H, N, O, or F will have hydrogen bonding. It requires a specific arrangement where hydrogen is bonded directly to one of the three aforementioned atoms.
Hydrogen Bonding: A Unique Intermolecular Force
Hydrogen bonding stands out from other intermolecular forces for several reasons:
- Strength: It's significantly stronger than other dipole-dipole interactions and van der Waals' forces but is still much weaker than covalent bonds.
- Special Requirement: Occurs when hydrogen is bonded to fluorine, nitrogen, or oxygen (highly electronegative atoms).
- Structural Implications: Hydrogen bonding often leads to unique structural properties in compounds. For example, the anomalous properties of water, such as its high boiling point, are attributed to hydrogen bonding.
- Biological Importance: Hydrogen bonding plays a key role in the structure of DNA and the secondary structures of proteins.
Hydrogen bonding between water molecules. The dotted line represents a hydrogen bond.
Image courtesy of Roland.chem
IUPAC Definition of the Hydrogen Bond
The IUPAC (International Union of Pure and Applied Chemistry) recently updated its definition of a hydrogen bond. Previously, it was considered an interaction mainly displayed by hydrogen attached to nitrogen, oxygen, or fluorine. However, the modern definition acknowledges the existence of hydrogen bonds with other atoms and in various conditions. The new definition emphasises:
- The electrostatic nature of the hydrogen bond.
- It can be stronger or weaker than the typical dipole-dipole interaction.
- The possibility of interaction of hydrogen with atoms other than N, O, or F in certain circumstances.
Note: Students should consult the latest IUPAC publications or trusted academic sources for the exact phrasing and detailed definition.
FAQ
No, not all molecules containing hydrogen atoms can form hydrogen bonds. Hydrogen bonding is a special kind of dipole-dipole interaction and is particularly strong. For a molecule to participate in hydrogen bonding, the hydrogen atom must be covalently bonded to a highly electronegative atom, specifically nitrogen (N), oxygen (O), or fluorine (F). This arrangement ensures that there's a significant positive charge on the hydrogen atom, making it capable of interacting with a lone pair of electrons on another electronegative atom, resulting in a hydrogen bond.
Water's high boiling point, compared to other molecules of similar size, is due to its capacity for hydrogen bonding. Each water molecule can form up to four hydrogen bonds: two through its hydrogen atoms and two through the lone pairs on its oxygen atom. This dense network of strong hydrogen bonds requires a significant amount of energy to break, leading to water's high boiling point. Without these hydrogen bonds, water would boil at a temperature similar to other molecules of its size that only experience van der Waals' forces.
Intermolecular forces play a pivotal role in determining the physical state of a substance at room temperature. Substances with strong intermolecular forces, like hydrogen bonding, tend to be solids or liquids at room temperature due to the strong attractions between molecules. In contrast, substances with weak intermolecular forces, such as London Dispersion Forces, are often gases at room temperature because these weak forces are easily overcome, allowing molecules to move apart from each other. For instance, water, with its hydrogen bonding, is a liquid at room temperature, whereas methane, which only has van der Waals' forces, is a gas.
The updated IUPAC definition is seen as more inclusive because it moves away from the rigid idea that hydrogen bonding can only occur when hydrogen is bonded to nitrogen, oxygen, or fluorine. The new definition recognises that hydrogen bonds can form under a wider range of conditions and with other atoms, provided that the conditions for significant charge disparities exist. By focusing on the electrostatic nature of the interaction, the definition includes a broader range of molecular interactions, highlighting the versatility and varying strengths of hydrogen bonds in different contexts.
Larger atoms or molecules possess more electrons, leading to a larger electron cloud. The size of this electron cloud means that there are more opportunities for fluctuations in electron density. As a result, there's a greater chance of temporary dipoles forming, which induces dipoles in neighbouring molecules. This chain of events leads to stronger van der Waals' forces, or London Dispersion Forces, in larger molecules or atoms. So, generally, as molecular size or molar mass increases, the strength of these forces increases, leading to higher boiling and melting points for such substances.
Practice Questions
Van der Waals' forces, also known as London Dispersion Forces, are the weakest type of intermolecular forces. They arise due to temporary dipoles created by fluctuations in the electron cloud around molecules and are present in all molecules. Dipole-dipole interactions are stronger and occur between molecules that have permanent dipoles, typically in polar molecules. Hydrogen bonding is a special type of dipole-dipole interaction and is stronger than the aforementioned forces. It occurs when hydrogen is bonded to highly electronegative atoms like nitrogen, oxygen, or fluorine. To determine which force is present, one must first analyse the molecule for polarity and the presence of hydrogen bonded to N, O, or F.
The updated IUPAC definition of a hydrogen bond broadens our understanding of this type of interaction. Previously, hydrogen bonding was mainly considered an interaction when hydrogen was attached to nitrogen, oxygen, or fluorine. However, the new definition acknowledges the possibility of hydrogen bonds with other atoms under certain conditions. It emphasises the electrostatic nature of the bond and states that its strength can be stronger or weaker than typical dipole-dipole interactions. The significance of this update lies in its acknowledgement of the diverse scenarios in which hydrogen bonding can occur, making it more encompassing and reflective of the complexities of molecular interactions.
