All matter is governed by attractive forces operating at the atomic and molecular level. These forces are categorized into two types: intramolecular forces, which exist within a molecule, and intermolecular forces (IMFs), which exist between separate molecules. Understanding the distinction and relative strength between these forces is fundamental to predicting a substance’s chemical stability and physical state. The intramolecular force is typically a covalent bond, while the attraction drawing separate units close is the intermolecular force. This difference in strength dictates nearly every observable property of a compound.
Covalent Bonds: The Internal Structure
Covalent bonds represent a powerful type of intramolecular force, serving as the internal structure that defines a chemical compound. These bonds are formed when two atoms achieve stability by sharing one or more pairs of valence electrons. The energy required to break the bond, known as bond dissociation energy, is a direct measure of the bond’s strength.
For example, the energy needed to break a single oxygen-hydrogen bond in a water molecule is approximately 463 kilojoules per mole (kJ/mol). The sharing of electrons establishes a stable balance between the attractive forces of the nuclei and the repulsive forces between the positively charged nuclei. Breaking a covalent bond fundamentally changes the chemical identity of the substance, converting a molecule into individual atoms or different molecular fragments.
The shared electrons are often not distributed equally between the two atoms, particularly when the atoms have different electronegativities, leading to a polar covalent bond. This unequal sharing creates partial positive and negative charges on the bonded atoms, which are still locked within the molecule. The robust nature of these electron-sharing interactions is why molecules maintain their integrity through most chemical processes.
Intermolecular Forces: Connections Between Molecules
Intermolecular forces (IMFs) are the attractive forces that occur between distinct molecules, acting as a kind of molecular “stickiness” that draws them together. These forces are non-covalent and do not involve the sharing or transfer of electrons, which makes them significantly weaker than the bonds within a molecule. IMFs are based on electrostatic attractions, arising from temporary or permanent partial charges on the molecules.
The weakest type of IMF, present in all substances, is the London Dispersion Force, which results from the instantaneous, fleeting polarization of electron clouds in a molecule. Slightly stronger are Dipole-Dipole interactions, which occur between permanent positive and negative ends of neighboring polar molecules. These attractions are a direct consequence of the unequal electron sharing that creates polar covalent bonds within the molecules themselves.
The strongest common type of intermolecular force is the Hydrogen Bond, which is a special, powerful form of dipole-dipole interaction. This attraction occurs when a hydrogen atom covalently bonded to a highly electronegative atom (nitrogen, oxygen, or fluorine) is attracted to a lone pair of electrons on a neighboring molecule. These forces are responsible for holding molecules together in the liquid and solid states, but they are attractions, not full chemical bonds.
Relative Strength of Chemical Interactions
Covalent bonds are substantially stronger than any type of intermolecular force, often by a factor ranging from 10 to 100 times. The fundamental difference lies in the nature of the interaction: covalent bonds involve overcoming the nuclear attraction to a shared electron pair, while IMFs involve only overcoming a lesser electrostatic attraction between partial charges.
To illustrate this magnitude difference, the energy required to break a typical covalent bond falls in the range of 100 to 400 kJ/mol. In stark contrast, the strongest IMFs, such as hydrogen bonds, typically have dissociation energies around 15 to 25 kJ/mol, and weaker London Dispersion Forces can be as low as 0.05 kJ/mol. This immense energy gap means that a molecule is far more likely to separate from its neighbors than to break apart internally.
Considering water as a specific example, it takes 927 kJ to break the two internal O-H covalent bonds in one mole of water molecules. However, it only requires about 41 kJ to overcome the IMFs holding one mole of liquid water together, converting it into steam. This vast difference in energy demonstrates why covalent bonds are the dominant force in determining a molecule’s chemical structure.
How Interaction Strength Determines Physical Properties
The disparity in strength between covalent bonds and intermolecular forces is directly responsible for how substances behave in the physical world. When a substance undergoes a phase change, such as melting a solid or boiling a liquid, the energy being supplied is used to overcome the weaker intermolecular forces. The covalent bonds within the individual molecules remain completely intact during these physical transitions.
For instance, when liquid water boils, the energy input is sufficient to separate the individual H₂O molecules from one another, allowing them to escape into the gaseous state. The atoms of hydrogen and oxygen remain bonded to each other, meaning the substance is still water, just in a different state. This is why substances with stronger IMFs, like water with its hydrogen bonds, have higher boiling and melting points than those with only weak London Dispersion Forces.
Only in a true chemical reaction, or through the application of massive amounts of energy, do the stronger covalent bonds break, resulting in a change in the substance’s chemical composition. The influence of IMFs on physical properties extends beyond phase changes to include viscosity and surface tension. Stronger attractions make the molecules resist movement and separation more effectively, determining the bulk properties of nearly all molecular materials.