Would it be possible to safely cool and drink a glass of water with anything else than the Ih form of ice?

Here and here you can see that some alternative forms of ice have a higher density than water, hence they would sink.

Would it be possible to have the coolest party ever, where the attendees would drink from glasses where the ice is sinking instead of floating?

  • $\begingroup$ Not sure about the alternate forms of the ice, but if you can arrange it so nobody drinks from that glass, you can use heavy water to make ice cubes that sink. Pretty cool trick, and you can win some small bets that way. ;) $\endgroup$ – CoilKid Oct 1 '15 at 0:29
  • $\begingroup$ On a similar note, you may find this interesting (pop-sci article) or this (from his website). In my opinion, that book is totally worth the cost. $\endgroup$ – CoilKid Oct 1 '15 at 3:06

As you can see from the phase diagram plot in the first link you provided, the only other ice phase which is stable at atmospheric pressure is ice XI, and its density is about the same as that of the most familiar ice phase (ice Ih). The other denser ice phases that you see on the phase diagram are only stable at pressures significantly above 1 atmosphere. As far as I'm aware, none of those high-pressure phases of ice are metastable, so you would have no chance of synthesizing any of those high-pressure ice phases in a high-pressure device (e.g., diamond anvil cell, Paris-Edinburgh cell, etc.) and then trying to retain the phase as you download the pressure back to atmospheric pressure.

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    $\begingroup$ Wikipedia thinks that IceIV is both metastable and denser than water. Amorphous ice forms at atmospheric pressures in a low-density variety, which is apparently the common in outer space as the surface of comets and icy planets. It might not sink in water, but cooling ones drink with space-ice sounds like a pretty cool trick and talking point. $\endgroup$ – tpg2114 Oct 1 '15 at 3:48
  • $\begingroup$ I thought maybe this particular kind of ice could be formed under high-pressure/low-teperature, then moved to atmospheric conditions. Kind of what happens when diamonds are formed under high-pressure/high-temperature and then they remain stable. $\endgroup$ – Leo Nov 12 '15 at 15:29
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    $\begingroup$ @Leo - You mean a metastable phase? I don't believe that this ice phase is metastable. I work in high-pressure research, and can state that metastable high-pressure phases that can be brought all the way down to atmospheric pressure are not very common at all. Diamond is a notable exception, but note that the energy difference per atom between the diamond and graphite phases of carbon at room temperature and pressure is only about 0.03 eV. That's very small and there is a large energy barrier between the two phases. $\endgroup$ – Samuel Weir Nov 12 '15 at 17:53

From the diagrams on the webpages you linked, it appears that other ice phases begin to form at around 200 MPa of pressure, and about $-20\text{C}^{\circ}$. Keep in mind normal water freezes at $0\text{C}^{\circ}$, and air pressure at sea level is around 0.101MPa. That means an ice cube made of one of these phases would sublimate or explode very quickly. If it didn't explode, the water in the cup would probably freeze in pretty short order.


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