Intermolecular Forces (IMFs) are the weaker attractions that keep molecules close, often called “molecular glue.” The strength of these IMFs determines a substance’s physical state (solid, liquid, or gas) and properties like boiling point or solubility. To understand the diverse behavior of matter, it is necessary to compare the relative strengths of different IMFs. This analysis focuses on two common forms of attraction—dipole-dipole forces and hydrogen bonding—to determine which interaction is stronger.
The Foundation: Understanding Dipole-Dipole Forces
Dipole-dipole forces are a form of electrostatic attraction that occurs between polar molecules. A molecule becomes polar when electrons are shared unequally between atoms due to a difference in electronegativity. This uneven sharing results in a permanent separation of charge, creating a partial negative end (\(\delta-\)) and a partial positive end (\(\delta+\)) on the molecule, which is called a dipole.
These forces arise from the attraction between the partial negative pole of one molecule and the partial positive pole of a neighboring molecule. This attraction is directional, meaning the molecules orient themselves to maximize the favorable electrostatic interaction. A simple example is hydrogen chloride (\(\text{HCl}\)), where the highly electronegative chlorine atom pulls electron density away from the hydrogen atom, creating a permanent dipole moment. Dipole-dipole forces are present in all molecules that possess this permanent charge separation.
The energy required to overcome these attractions is relatively modest, typically falling in the range of \(5\) to \(25\text{ kJ/mol}\). For instance, the dipole-dipole interaction in \(\text{HCl}\) is weak, around \(3.3\text{ kJ/mol}\), which is why hydrogen chloride is a gas at room temperature. These forces are strong enough to influence a substance’s physical state.
The Special Case: What Makes Hydrogen Bonding Unique
Hydrogen bonding is not a separate force entirely but is categorized as a particularly strong type of dipole-dipole interaction. This special strength is due to a very specific set of structural requirements involving only three highly electronegative elements: nitrogen (\(\text{N}\)), oxygen (\(\text{O}\)), or fluorine (\(\text{F}\)). For a hydrogen bond to form, a hydrogen atom must be covalently bonded to one of these three atoms in one molecule.
Because these atoms (\(\text{N}\), \(\text{O}\), or \(\text{F}\)) are highly electronegative, they pull the bonding electrons strongly towards themselves. This leaves the small hydrogen atom with a large partial positive charge (\(\delta+\)). The tiny size of the hydrogen atom means its positively charged nucleus is nearly exposed. This concentrated positive charge creates an exceptionally powerful electrostatic attraction to a lone pair of electrons on an adjacent \(\text{N}\), \(\text{O}\), or \(\text{F}\) atom in a neighboring molecule. This intense, directional interaction is far stronger than the broader charge distribution found in a typical dipole-dipole interaction.
Water (\(\text{H}_2\text{O}\)) serves as the classic example, where each molecule can act as both a hydrogen-bond donor and an acceptor. The resulting large partial charges allow water molecules to form an extensive network of strong hydrogen bonds, fundamentally shaping the properties of liquid water and ice.
Direct Comparison: Answering the Strength Question
The direct answer to the strength comparison is that hydrogen bonding is significantly stronger than a typical dipole-dipole interaction. While both are driven by the attraction between partial positive and partial negative charges, the unique structure of hydrogen bonding results in a far greater interaction energy. Hydrogen bonds typically exhibit strengths in the range of \(10\) to \(40\text{ kJ/mol}\), with some examples reaching \(50\text{ kJ/mol}\).
Comparing the strongest hydrogen bonds to the weakest dipole-dipole forces reveals a substantial difference in magnitude. For instance, the energy to break the intermolecular forces in water is approximately \(19\text{ kJ/mol}\), while the dipole-dipole interaction in \(\text{HCl}\) is only \(3.3\text{ kJ/mol}\). This demonstrates that hydrogen bonding can be several times stronger than standard dipole-dipole forces.
The difference in energy results from the efficiency of the charge interaction. The extreme polarity and small size of the hydrogen atom in \(\text{H-N}\), \(\text{H-O}\), or \(\text{H-F}\) bonds allow for a much closer approach and a more focused electrostatic pull between molecules. This intense partial charge concentration requires a much greater input of energy to separate the molecules compared to a general dipole-dipole interaction.
Practical Impact: Why Molecular Forces Matter
The difference in strength between hydrogen bonding and dipole-dipole forces has profound, observable effects on the physical properties of substances. The most commonly noted consequence is the elevation of boiling points for hydrogen-bonded compounds. Since boiling requires supplying enough energy to overcome the intermolecular forces and separate the molecules into the gas phase, stronger attractions demand higher temperatures.
This is clearly demonstrated when comparing water (\(\text{H}_2\text{O}\)) to hydrogen sulfide (\(\text{H}_2\text{S}\)), a molecule with a similar structure but no hydrogen bonding. Despite \(\text{H}_2\text{S}\) being a heavier molecule, its lack of hydrogen bonds means it boils at a very low \(-60^{\circ}\text{C}\). In stark contrast, the powerful hydrogen bonds in water force its boiling point to be \(100^{\circ}\text{C}\).
The ability of a substance to form hydrogen bonds also dictates its solubility in water. Water is often called the universal solvent because it can form hydrogen bonds with other molecules, such as alcohols. This principle, summarized as “like dissolves like,” explains why substances like ethanol, which can form hydrogen bonds with water, mix readily, while non-hydrogen-bonding molecules like oils do not. The unique strength of the hydrogen bond is responsible for many of the anomalous properties of water.