Why doesn't water start to boil in soils? Consider the following two scenarios:
I take a vessel filled with some water and evacuate the air, the water will (or rather can) start to boil; if the pressure in the vessel is << atmospheric pressure. (Say -1e3 hPa)
Now, consider a nearly dry porous medium such as an arbitrary soil clod. In this
water can move according to cappilary forces from one place to an (energetically) lower place. 
Which force is commonly measured in "water potential" which can be as low as ~ -1e7 hPa.
Why does water boil in the evacuated vessel (-1e3 hPa) and not in the dry soils with water potentials < -1e3 hPa?
Where is the error in my thinking?
 A: While the capillary pressure in soil is many orders of magnitude lower than the atmospheric pressure, you also need to remember that in soil, the water is still in contact with the atmosphere, and thus is at atmospheric pressure plus capillary pressure.
Since atmospheric pressure is orders of magnitude larger than capillary pressure, the pressure on the water in the soil is (I think?)
$$P=P_\text{atm}+P_\text{cap}\approx P_\text{atm}.$$
Since water doesn't boil at atmospheric pressure, the water in soil doesn't boil, even though the capillary pressure component alone is low.
A: If a monolayer of water molecules is adsorbed on the surface of pore, the van der Waales potential acting on a molecule is increased due to negative curvature of the water surface (eg. concave meniscus) and capillary condensation sets in already below the saturation pressure of water. It shifts the sorption isotherm of the porous material towards higher equilibrium moisture. As a consequence a lower pressure is needed in this case to evaporate the water as compared to ordinary evaporation.
A: As mentioned in DumpsterDoofus's answer, those potentials are likely relative to atmospheric pressures but at magnitudes of 1 GPa that's pretty much irrelevant.
It's the pressure of the gas vapor that determines equilibrium. The pressure of the liquid is nearly irrelevant, so in this case even if the capillary forces lowered the pressure below absolute zero, the evaporation rate would remain unaffected.
A good example of this in nature are trees that are taller than 15m. The pressure near the roots is near atmospheric and reduces past absolute zero, up to the leaves where the capillary pressures allow negative liquid pressures to be in equilibrium with the positive atmospheric pressure.
A molecular explanation as to why it is the vapor pressure that determines equilibrium is here:  Thermodynamics of evaporation
