# Boil Water At Low Temparatures

I'm describing a little experiment that i conducted. Water or any liquid boils when its vapor pressure is greater than the pressure of the air above it.You can demonstrate this easily by using a disposable syringe. Suck water in the syringe so that it is about a quarter filled.(do it underwater so that there are no bubbles). Now close the open end of the syringe with your thumb and pull the piston backwards. Now there is a temporary vacuum created b/w the nozzle's end and the piston, where the pressure is zero. The water will now start to boil!

PS: Remove the needle while conducting this experiment...

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Danger
Gold Member
PS: Remove the needle while conducting this experiment...
:rofl:

That's a great example of a simple experiment that kids can do at home. If you have others that demonstrate different scientific principles (and are safe), please share.

Defennder
Homework Helper
I never knew we could do that. Did anyone try it out to see if it works?

Gold Member
I love these little experiments. I have done quite a few, most of them came from New Scientist books. I'll make a thred of them.

Thank you guys

Thank you guys for the support. And Mr.Danger I'm a die-hard fan of Mathematica...three cheers for it.

uby
Water or any liquid boils when its vapor pressure is greater than the pressure of the air above it.
be careful how you describe this! from the standpoint of chemical thermodynamics:

any liquid in its standard state (for convenience, this is usually pure liquid at 1atm total pressure and 298.15K although you can certainly pick any reference state you wish!) will be in equilibrium with its vapor at the liquid/gas interface by the reaction:
X(l) <--> X(g)
The equilibrium partial pressure of X(g) over X(l) at the liquid/gas interface is estimated using some assumptions about the ideal behavior of the gas and it's reference state (Raoultian or Henrian depending on what you're trying to calculate). Tabulated Gibbs free energy of formation data can be used to find the equilibrium activity and corresponding partial pressure under these assumptions.

The equilibrium partial pressure is typically tabulated with 1atm total pressure. If this value is lowered, then your driving force to form new gas species increases as your are perturbing the equilibrium in the gas forming direction (Le'Chatlier's principle). The shift in the equilbrium as a function of total pressure can be calculated as well, and you can estimate the temperature at which any liquid will boil for a given total pressure by accounting for the equilibrium shift.

I'm new to these forums, so I don't yet know how to write out the equations. But the idea is simple: going from a ratio of activities to partial pressures involves normalization to the total pressure (by convention: partial pressure = concentration * total pressure !for an ideal gas only!)

I never knew we could do that. Did anyone try it out to see if it works?
I tried with water and with 99% ethanol, with about a quarter volume and with smaller quantities. No boiling observed. What did I miss?

uby
I tried with water and with 99% ethanol, with about a quarter volume and with smaller quantities. No boiling observed. What did I miss?
when you lift the plunger in the syringe, you are doing work on the system. to a good approximation, this is isothermal work. so, the change in volume is inversely proportional to the change in pressure.

if the plunger is not sealed, then the pressure will not be changed from its environment.

if the change in pressure at room temperature is not enough of a shift in the equilibrium to cause boiling, you will not see much of a difference except some condensation on the syringe tube walls.

i would recommend finding a low-boiling azeotrope for your demonstration. look on wikipedia for "Azeotrope_data" for a good tabulation of azeotropic mixture data. use what you're comfortable using in a lab environment. be sure the syringe is compatible with the compositions your work with!

Redbelly98
Staff Emeritus
Homework Helper
when you lift the plunger in the syringe, you are doing work on the system. to a good approximation, this is isothermal work. so, the change in volume is inversely proportional to the change in pressure.
It's not clear that this is the case. Volume and pressure are inversely proportional in a gas, where the amount of gas is fixed. Since we have liquid evaporating, the amount of vapor is not fixed. Also, the instructions are to have just liquid, no air, initially.

if the plunger is not sealed, then the pressure will not be changed from its environment.
That may be the key factor. In normal use, a syringe must only seal against liquid leakage, at low pressure differential. To seal against air leakage, at close to 1 atm differential, requires a higher quality seal. It should only need to seal for several seconds in order to demonstrate the effect.

if the change in pressure at room temperature is not enough of a shift in the equilibrium to cause boiling, you will not see much of a difference except some condensation on the syringe tube walls.
mannattil appears to have made it work, using ordinary water:

mannattil said:
I'm describing a little experiment that i conducted.

Redbelly98
Staff Emeritus
Homework Helper
I tried with water and with 99% ethanol, with about a quarter volume and with smaller quantities. No boiling observed. What did I miss?
I don't know. Hot tap water should boil more readily than room-temperature or cold water. Pulling the plunger quickly vs. slowly might make a difference (not sure which would work better, so try both).

I think we have syringes at work, I might try this out later.

Here are some vapor pressure numbers. Hot water should work better than room-temperature ethanol:

Water:
20 C, 18 Torr
50 C, 92 Torr

Ethanol
20 C, 40-50 Torr

Isopropanol
20 C, 20-25 Torr

Danger
Gold Member
And Mr.Danger I'm a die-hard fan of Mathematica...
That's just 'Danger'; I'm only Mr. Danger to people I don't like.

Redbelly98
Staff Emeritus