Solution - Class 8 PDF Download

Yes. you can, every thing is possible in this world.
In fact the simple combustion reaction of O2(g)+2H2(g)→2H2O(g)O2(g)+2H2(g)→2H2O(g)will give you the desired result, once you’ve condensed out the water.

However: Be forewarned. Combining 1 mole of Oxygen with 2 moles of H2, corresponds to about 70L of gas at STP. This will make 2 moles of water (or about 32ml) + release a lot of heat. So for a 300 ml glass, you’ll need ~227L of O2 gas and 454L of H2 gas (at STP)…and a you’ll probably wake up your neighbors if you do this all at once.

Alternatively, you can be smart about it and make a hydrogen fuel cell and catalyze the reaction on a Pt surface. Note that catalyzing a reaction can work both ways. You can accelerate slow reactions or slow down fast reactions. In this case you want to slow down the reaction by first dissociating O2 and H2 on a Pt surface (following a Langmuir-Hinschelwood mechanism).

O2(g)⇌2O(s)O2(g)⇌2O(s)

H2(g)⇌2H(s)H2(g)⇌2H(s)

Then, have

O(s)+H(s)⇌HO(s)O(s)+H(s)⇌HO(s)

and finally

HO(s)+H(s)→H2O(g)→H2O(liq)HO(s)+H(s)→H2O(g)→H2O(liq)

where X(s), X(g), and X(liq) denote whether the species is adsorbed on the surface (s), in the gas phase (g), or in liquid phase (liq).

One can then use the excess heat (carried off by the H2O(g)H2O(g) and condensed to liquid) as an energy source for like a steam turbine. But, that’s now an engineering problem.

I’ll leave it as an exercise for the reader (and my p-chem class this spring!) to work out the rate law.