Water typically boils and turns to steam at \(100^\circ\text{C}\) (\(212^\circ\text{F}\)) at standard sea-level atmospheric pressure. However, the boiling point of any liquid is not fixed; it is a function of the surrounding pressure. The simple answer to whether pure water can exist as a liquid at \(110^\circ\text{C}\) is yes, but achieving this state requires manipulating the physical environment. Liquid water can exceed its normal boiling temperature through two primary physical mechanisms: one that is thermodynamically stable and one that is temporary and unstable.
How Pressure Keeps Water Liquid Above \(100^\circ\text{C}\)
The most stable method for maintaining water in its liquid phase at \(110^\circ\text{C}\) involves increasing the external pressure exerted on the liquid. Boiling is a physical process that occurs when the vapor pressure produced by the water molecules equals the pressure of the surrounding atmosphere pressing down on the liquid’s surface. At standard atmospheric pressure (one atmosphere, or \(760\) torr), this balance is achieved precisely at \(100^\circ\text{C}\).
When heat is applied to water at \(100^\circ\text{C}\), the energy is consumed entirely in the work of changing the liquid to a gas, known as the latent heat of vaporization. This added energy allows water molecules to overcome the attraction forces holding them together in the liquid state. By increasing the external pressure, the boiling point is effectively raised because the water molecules require more kinetic energy to generate a vapor pressure strong enough to push back against the greater downward force.
To keep water liquid at \(110^\circ\text{C}\), the external pressure must be raised to match the water’s saturated vapor pressure at that temperature. Scientific data shows that the vapor pressure of water at \(110^\circ\text{C}\) is approximately \(1075\) torr. Since \(1\) atmosphere is \(760\) torr, this means the pressure must be increased to about \(1.42\) atmospheres. At any pressure slightly above \(1075\) torr, the water will remain a stable liquid even as its temperature reaches \(110^\circ\text{C}\).
This relationship is mapped out on a phase diagram, which shows the conditions of temperature and pressure necessary for water to exist as a solid, liquid, or gas. The liquid-gas boundary line indicates how the boiling point increases as pressure rises, confirming that the liquid state can be extended far beyond \(100^\circ\text{C}\). This high-temperature liquid water is scientifically known as “superheated water” or “subcritical water.” The liquid remains stable because the elevated pressure continuously forces the water molecules to remain compressed, preventing the formation of steam bubbles.
Superheating: Liquid Water Without Boiling
A separate and far less stable condition that allows water to exceed \(100^\circ\text{C}\) is known as superheating. This phenomenon occurs even at normal atmospheric pressure, but it represents a metastable state, meaning it is only temporarily stable. In a superheated state, the liquid’s temperature is above its boiling point, but the phase change to gas has not yet begun.
The main reason this temporary state can exist is the absence of nucleation sites. Nucleation sites are microscopic imperfections like tiny scratches on a container wall, dissolved gas pockets, or dust particles, all of which provide a surface for vapor bubbles to initially form. When pure water is heated in a container with exceptionally smooth surfaces and no impurities, these sites are unavailable.
Without a nucleation site, the water molecules must rely on surface tension to suppress the formation of a vapor bubble. The high surface tension of water acts like an elastic skin, requiring a small amount of extra energy, and therefore a slightly higher temperature, to overcome the pressure and start the boiling process. The superheated water appears deceptively calm and still, having no visible bubbles, even as it reaches temperatures like \(110^\circ\text{C}\).
This condition is highly dangerous because the water holds excess thermal energy that can be instantly released. Any sudden disturbance, such as jostling the container, inserting a spoon, or adding a powder, can provide the necessary nucleation sites. When this happens, the stored energy is released rapidly, causing the water to flash boil into steam explosively. This sudden eruption can spray scalding water and cause severe burns.
Where High-Temperature Liquid Water Exists
High-temperature liquid water is a feature of both advanced engineering and natural geological systems. Engineered devices rely on the principle of increased pressure to safely raise the boiling point, using a robust sealed vessel. A household pressure cooker is a common example, where steam builds up inside the sealed pot, increasing the internal pressure and allowing the water to reach temperatures well above \(100^\circ\text{C}\) for faster cooking.
Larger industrial applications also employ this principle, such as steam boilers in power plants and the cooling systems of pressurized water nuclear reactors. In these systems, water is kept at extremely high temperatures, often hundreds of degrees Celsius, to efficiently transfer or convert thermal energy while remaining in the liquid phase. Even the cooling system in a standard car operates under slight pressure to allow the coolant mixture to safely exceed \(100^\circ\text{C}\) without boiling.
Nature also produces high-temperature liquid water under pressure in deep-sea hydrothermal vents. At the ocean floor, the immense weight of the water column creates pressures that can be hundreds of times greater than at the surface. This pressure allows water erupting from the Earth’s crust to remain a liquid even at temperatures that can reach \(350^\circ\text{C}\) or more. Conversely, the unstable phenomenon of superheating is most often encountered accidentally when heating purified water in a smooth container, like a glass mug in a microwave oven.