Jim Elsey is a mechanical engineer who has focused on rotating equipment design and applications for the military and several large original equipment manufacturers for 47 years in most industrial markets around the world. Elsey is an active member of the American Society of Mechanical Engineers, the National Association of Corrosion Engineers and the American Society for Metals. He is the general manager for Summit Pump Inc. and the principal of MaDDog Pump Consultants LLC. Elsey may be reached at firstname.lastname@example.org.
- Vapor pressure is the pressure required for a liquid to boil at a given temperature.
- Each liquid has its own unique vapor pressure values and characteristics.
- The vapor pressure of a liquid will increase with temperature and vice versa.
- The vapor pressure of a liquid at a given temperature is the pressure that it will flash (change state) to vapor if heat is added.
- The opposite is also true—the pressure where the vapor at a given temperature will condense to liquid form if heat is removed.
- The surface area of the liquid has no effect on the vapor pressure.
- Equilibrium is when the evaporation rate of the fluid is equal to the condensation rate measured at the surface where the fluid (liquid) and the atmosphere (gas) meet.
- When the vapor pressure is equal to the atmospheric (or ambient) pressure that will be the boiling point.
- Saturation temperature means boiling point.
- The saturation temperature is the temperature for a corresponding saturation pressure at which a liquid boils into its vapor phase.
Any liquid in an open container will evaporate to a vapor unless some other force is present to prevent the change from occurring. In most examples, that force is simply atmospheric pressure.
The fluid exerts a certain pressure to the atmosphere and in return the atmosphere exerts a counter pressure. For water below 212 F, the atmospheric pressure is greater than the vapor pressure.
Those two pressures (liquid and atmospheric) would be in equilibrium if the water was at 212 F at sea level where the water would begin to boil. If you were at higher elevations (lower ambient pressure) the water would boil at a lower temperature.
For instance, at 10,000 feet above sea level, water would boil at about 194 F.
A slightly different way to think of vapor pressure is to equate it with the boiling point for that liquid at some pressure. Just keep in mind that, in reality, it is more complicated. As previously mentioned, water at room temperature and pressure will not vaporize until you raise the temperature to 212 F.
However, you can also boil water at room temperature by reducing the pressure. For 68 F water the boiling point could be obtained by lowering the pressure to 0.34 psia. Note this value of absolute pressure is approaching an extremely high level of vacuum, so this does not often occur except in special processes.
Why are we discussing this subject at all? Because cavitation occurs in the pump when the fluid pressure drops below its vapor pressure.
This phenomena can occur by either increasing the temperature or reducing the pressure, but in almost all cavitation cases it is because the lowest pressure the fluid will experience in the pump system is the area immediately in front of the impeller eye.
If the pressure is below the vapor pressure, it will flash to vapor—not because of a temperature change, but because of a pressure change.