Surface tension in liquids is caused by cohesive forces, the attractive forces between molecules of the same type. Every molecule inside a liquid is pulled equally in all directions by its neighbors, so those forces cancel out. But molecules at the surface have no neighbors above them, only air. They experience a net inward pull that draws them back toward the bulk of the liquid, creating a tight, elastic-like “skin” at the surface.
How Cohesive Forces Create Surface Tension
Picture a molecule sitting deep inside a glass of water. It’s surrounded on all sides by other water molecules, each tugging on it with equal strength. Those pulls balance out perfectly, so the molecule sits in equilibrium. Now picture a molecule right at the surface. It has neighbors beside it and below it, but nothing pulling from above. The result is a net downward and inward force that yanks surface molecules toward the interior of the liquid.
This imbalance has a visible consequence: the liquid minimizes the number of molecules exposed at the surface. That’s why small drops of water form spheres. A sphere has the smallest possible surface area for a given volume, which satisfies the inward pull most efficiently. The force per unit length along the surface is what physicists measure as surface tension, expressed in newtons per meter (N/m).
Why Water’s Surface Tension Is Unusually High
Not all cohesive forces are created equal. In most organic liquids like benzene, the molecules stick together through relatively weak attractions called van der Waals forces. Benzene’s surface tension is about 28 millinewtons per meter. Water, by contrast, measures 72.75 mN/m at 20°C, nearly three times higher.
The difference comes down to hydrogen bonding. Each water molecule can form up to four hydrogen bonds with its neighbors, creating a dense, interconnected network. Hydrogen bonds are substantially stronger than van der Waals forces, so breaking even a few of them at the surface costs a disproportionate amount of energy. That high energy cost is what gives water its remarkably strong surface tension compared to most other common liquids.
Temperature Changes the Strength
Heating a liquid weakens its surface tension. As temperature rises, molecules move faster and spend less time in close contact with their neighbors. The cohesive forces between them become less effective, and the surface “skin” loosens. This relationship is roughly linear for most liquids: the hotter the liquid, the lower the surface tension. It’s one reason hot water is better at wetting surfaces and spreading into tight spaces than cold water.
Cohesion vs. Adhesion
Cohesive forces (attraction between identical molecules) are only half the picture when a liquid touches a solid surface. Adhesive forces, the attraction between different types of molecules, also come into play. The balance between the two determines how a liquid behaves in contact with a container or a narrow tube.
When adhesive forces between water and glass are stronger than water’s internal cohesion, the liquid climbs the walls of a glass tube and forms a concave meniscus, the curved surface you see at the top of water in a graduated cylinder. This climbing effect is called capillary action, and it’s how water travels upward through plant roots and thin soil pores.
When cohesion dominates, the opposite happens. Mercury, for example, has very strong cohesive forces and weak adhesion to glass. It forms a convex meniscus and is actually pushed down in a narrow glass tube rather than pulled up. The contact angle between the liquid and the surface tells you which force is winning: angles less than 90 degrees mean adhesion dominates, angles greater than 90 degrees mean cohesion wins.
How Soap Disrupts the Force
Surfactants, the active ingredients in soap and detergent, are molecules with a split personality: one end attracts water and the other end repels it. When you add a surfactant to water, these molecules migrate to the surface and wedge themselves between water molecules. They physically interrupt the hydrogen-bonding network, reducing the net inward pull on surface molecules. The result is a measurable drop in surface tension, which is why soapy water spreads more easily across surfaces and penetrates fabrics better than plain water.
Surface Tension in the Real World
Water striders are the classic demonstration of surface tension at work. These insects are light enough that their weight never exceeds the upward force generated by the water’s surface tension. Their legs dimple the surface without breaking through, and the cohesive forces in the water push back against the depression like a trampoline. Research on several species shows that their jumping thrust comes primarily from surface tension rather than from pushing water downward. Their bodies are tuned so precisely to this force that the water surface remains unbroken even as they launch themselves vertically to escape underwater predators.
The same physics explains why a carefully placed steel needle can float on water despite being far denser than it, why raindrops are spherical, and why bubbles form. In every case, the cohesive forces between liquid molecules pull the surface into the smallest possible area, creating a measurable tension that resists stretching or breaking.