According to this video, when a valve is quickly shut, the water hammer effect can lead to a high negative pressure in a pipe downstream. This leads to dissolved gases coming out of solution with the water. At what pressures can this occur? The experimental setup in the video reaches 100kPa below atmospheric pressure.

Does this also lead to water splitting?

  • $\begingroup$ Not a specific pressure ; the lower the pressure , the more gasses will escape. $\endgroup$ Jan 17, 2021 at 15:30
  • $\begingroup$ re: electrolysis of water? no. $\endgroup$
    – Pete W
    Jan 17, 2021 at 17:16

2 Answers 2


The vaporization and dissolution of gases out of and into, respectively, of water is a dynamic equilibrium, meaning that gases are continually coming out of solution, and re-entering. We may assume that water that has been sitting for some time has reached its dynamic equilibrium point, of a certain quantity of dissolved gases for a fixed mass of water, under steady temperature and pressure.

There are various ways this equilibrium may be upset, the main factors being temperature and pressure (or more specifically, partial pressures of the dissolved species' in question). Also, nucleation points, ultrasound, etc.

Basically, any reduction in pressure under the equilibrium point will result in some outgassing. If you lower the pressure too much, then not only do the dissolved gases bubble out, water itself "boils" at a lower temperature and forms bubbles of water vapor.

So to summarize, there isn't a single point for gases coming out of solution, but a continuous variation. However, at every temperature there is a pressure at which water undergoes a liquid -> vapor phase change, and this relationship is described by water's phase diagram.


"negative pressure downstream, dissolved gases coming out" - this is correct, and that's a great video showing it. It's called cavitation in some contexts.

The question of at what pressure this happens has three parts, as far as I understand it.

(1) The local temperature and pressure (i.e. resulting from valve action) need to be such that you are at a supersaturated condition. Vapor pressure of would-be-gas-species > pressure. First and foremost, this depends on the amount of the species in question in solution to start with. There are also considerations of physical chemistry that affect the vapor pressure curve. However this is only a high level description of the equilibrium, you won't actually see bubbles when the temperature and pressure of this condition is reached, at least not instantly. This condition will be roughly good enough to slowly grow bubbles in corners or scratches in the side walls, but the fast phenomenon we are talking about is homogeneous nucleation.

(2) the classical description of bubble nucleation. This brings in the effect of curved liquid-gas interfaces. The effect of surface tension is called Laplace pressure, which increases pressure with curvature and tends to drive the gas back into solution. Vapor pressure is also altered, per the Kelvin equation.

(3) Lastly, classical predictions based on Laplace pressure have contradicted experiments, especially at small but well-above-molecular scales, which drags us deeper into thermodynamics. There may be some unresolved points of theory here such as how surface tension varies, and I am not that solid even on the basics at this level, but you need to be past something called the spinodal point, where the second derivative of Gibbs free energy is zero.

For engineering purposes, our goal is simply to ensure that there won't be any unintentional gas phase, so it is good enough to make calculations based on the equilibrium of vapor pressure, #1 in the list above. The solution is to control some or all of: pressure, temperature, concentration of would-be-gas-species in the fluid, or valve opening speed.


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