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Seeing some cool properties of water through the lens of its molecular structure

We all know the importance of water—our bodies are mostly water, we need it to survive, it’s the second most important ingredient in coffee… Geologically, it facilitates almost everything we know, from erosion to magma formation to rock fracture. I’m often struck by how so many of water’s unusual properties are determined by its chemistry and molecular structure –and in a very understandable way.

Waterfalls and cliff, New Zealand.

waterfalls in Fjordland, South Island, New Zealand.

Water molecules are polar
Many of water’s properties stem directly from its polar nature –and its polar nature comes right from its molecular structure. Here’s how. The water molecule’s two hydrogen atoms bond to the much larger oxygen atom in a way that forms a near right triangle. As a result, the hydrogens crowd one side of the oxygen while leaving the other side open. This asymmetry creates water’s well-known polarity: even though the two hydrogens and one oxygen balance each other, a slight positive charge exists on the side of the hydrogen atoms and a water molecule2slight negative charge exists on the far side of the oxygen.

 

And it’s this polarity that, among other things, creates water’s capillary action, its surface tension, its amazing ability to dissolve things, and its powerful ability to buffer extreme temperature changes.

 

Water ‘s polarity causes its individual molecules to stick to each other
120218-1The diagram below shows multiple water molecules –and even though they’re separate, they’re connected. The slight positive charge on the side of the hydrogen atoms of one molecule bonds with the slight negative charge on the hydrogen-poor side of another molecule. These bonds, called hydrogen bonds, are weak, but they keep the molecules together. This “stickiness” gives water its surface tension, which lets water extend over the brim of a cup, form large drops that beautifully reflect sunlight, or support creatures like water-striders that move over its surface. It also gives water capillary strength, so that trees, for example, can carry water up into their canopies.

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Water molecules joined by hydrogen bonds –USGS

The salt crystals that grow upwards from cracks in Death Valley’s salt pan give a dramatic example of capillary action. Water that ends up on the salt pan dissolves the salt and sinks into the ground. As the air temperature warms, however, water still on the surface begins to evaporate, which pulls the water just below the surface upwards by capillary action. In turn, that upward-moving water pulls the water below it upwards and so on. What’s more, as the water evaporates, it leaves behind salt, crystallizing into needles and ridges that grow upwards.

saltpanspreadmr

Salt crystals in Death Valley, California. Left: new salt crystals growing upwards from damp valley floor –note concentration along cracks. Right: After months to years of crystallization, large polygons of salt are separated by raised, salt-filled fractures.

It’s water’s polar nature that allows the salt to dissolve so readily in the first place. Salt, NaCl, is itself polar, with a slight positive charge on the side of the sodium atom and a slight negative charge on the side of the chlorine atom. With water then, the sodium ion readily bonds with an oxygen ion while the chlorine ion readily bonds with the hydrogen. The result is that the salt molecules break up and dissolve into the water. Water‘s polarity allows it to dissolve a wider range of substances than anything else if given enough time. Hence its description as the “universal solvent”.

Heat Capacity and mellow coastal climates
Heat Capacity is a measure of how much energy it takes to raise a substance’s temperature – and water requires a lot of energy, also because of its hydrogen bonding. Because the hydrogen bonds keep the water molecules together, energy is required to first break those bonds before individual water molecules can vibrate and create heat. This quality of high heat capacity gives water the ability to buffer large changes in temperature. As a result, coastal areas experience smaller temperature variations than continental interiors. Look at the graph below, showing the much narrower range of high temperatures for Newport, Oregon than for Ontario, Oregon –both of which are at the same latitude. Newport, though, is on the coast, whereas Ontario is inland.

Avg high temps-N-O-Oregon

Newport, being on the coast, experiences a narrower range of temperatures than Ontario, which is inland.

Freezing water
Then there’s water’s tendency to expand when it freezes. We see this phenomenon all the time. In fact, we’re so used to it that we probably don’t give it a second thought.  But nearly all other materials shrink and become denser when they freeze. If water did too, lakes would freeze solid in winter and lake-dwelling creatures would die as they became encased in ice! How about rock weathering and erosion? Expansion of freezing water in cracks and soils drives both frost cracking and solifluction, two processes critical to landscapes. The list goes on.
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This expansion-while-freezing also has to do with hydrogen bonds, as with decreasing temperatures, they become increasingly stable and incorporate freezing water into a hexagonal crystal lattice. These hexagons, while beautiful and symmetrical, are also more open-structured than uncrystalline (liquid) water and so significantly less dense.

Of course, there are dozens of other properties specific to water. If you’re interested, you can start with the US Geological Survey’s water properties site–the USGS, as usual, gives out a lot of free and delicious information!

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3 thoughts on “Seeing some cool properties of water through the lens of its molecular structure

  1. Thanks as always, Marli. Three comments:

    1. You have the structure of hexagonal ice wrong. You will find a good description at http://www1.lsbu.ac.uk/water/hexagonal_ice.html The individual molecules do *not* distort, but the structure is open, accounting for the lower density. Note that each oxygen links to 4 other oxygens through hydrgen bridges. This will matter in comment 3.

    2. “Salt, NaCl, is itself polar, with a slight positive charge on the side of the sodium atom and a slight negative charge on the side of the chlorine atom… salt molecules”

    To a good approximation, in solid salt the sodium has a charge of +1 and the chloride a charge of -1; full transfer, not slight. In the solid, there are no “salt molecules” but each Na+ is surrouned by 6 Cl- and vice versa in the familiar cubic NaCl crystal structure.

    3. Siica, SiO2, is another material that is less dense at the melting point as a solid rather than as a liquid. The reason is a bit like that for ice; each Si bridges to four others, via equally shared oxygens, just as in ice each oxygen bridges to four others via (partly) shared hydrogens.

    A consequence for both ice and silica is that the melting point *decreases* at higher pressure. This is relevant to why granite tends to form plutons, whil basalt flows. The silica in granite crystallises out as the melt rises, making the mixture stickier. Basalt, on the other hand, behaves normally; as it rises, the reduced pressure makes it more fluid.

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    • Thanks for the comments, Paul –very helpful! I think I probably oversimplified things too much and will modify the post –especially that part about crystallizing into ice.
      However, I’m not sure I understand what you’re saying about silica. Silica-rich magmas (rhyolitic/granitic) are far more viscous than silica-poor (basaltic) ones, which as I understand it, is the reason why basalt tends to reach the surface more often than rhyolite. It’s not that the silica is crystallizing out of the magma as it rises –in fact, it’s the silica-poor crystals that drop out first, so magmas tend to become more silica-rich as they rise.
      Thanks again.

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      • Seemslike I was wrong about the granite. I had thought that the silica was crystallising out as a result of reduced pressure, making the magma stickier as it rose, but cliearly I’m wrong on this. So what I said in (3) is wrong; the fact that silica (and also silicon and germanium, where the solid structure is also based on local tetrahedra) share the same behaviour as ice is irrelevant

        “Hexagonal” is a misleadng term; it refers to the symmetry of the unit cell, rather than to the presence of actual hexagons.

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