Field and Swamp: Animals and Their Habitats

Bubble, Bubble, Toil and Trouble:

The Hard Problem of (Very) Soft Water in Daily Life

Disclaimer:  This essay represent the views of a science student.  Nothing in this essay should be interpreted as a recommendation for a scientific experiment.  Some ideas presented in this essay are simply speculation.  This essay is meant to be about hard science, although it has its lighter moments.  If you discover any factual errors, please write to Dorothy E. Pugh at Contact Us.

Mark Matthews, Ph.D., from whom the author took General Chemistry I and II, generously reviewed this essay at my request for the soundness of the general scientific principles discussed here.  This essay is a tribute to the clarity of his presentation of that material as well as to the enthusiasm that he inspired in this student.  However, this does not mean that he subscribes to the rather eccentric point of view expressed in this article (and I think he may be a little embarrassed to have his name associated with it.)  At any rate, I, the author, take full responsibility for the correctness of the text.

For more details about the background of the author, Dorothy E. Pugh, see About Us. 

Abstract

Very soft water can be difficult to bathe with because when it's mixed with soap it's hard to remove.  However, rinsing in cool water can make this process easier.  This simple problem and proposed remedy are described using explanations of intermolecular bonds, solution formation, factors affecting solubility, equilibrium in reversible reactions, precipitate formation in irreversible reactions, mechanisms of and environmental influences on water acidity.

Introduction

Maybe Shakespeare got his inspiration for Macbeth's wife's "Out, damn'd spot!" hallucinatory scene from trying to bathe in very soft water (actually, she would have gotten the spot out, but not the soap off.)   Soap mixes nicely with soft water, but then it doesn’t rinse off, at least if you depend on more water to accomplish that.  If you’re used to hard water and it suddenly goes soft, it can be doubly frustrating.

Suppose a city changed its water treatment plan to make the water more acidic, and therefore softer, reportedly to reduce costs?  Suppose the city's water testing discovered this was leaching lead out of the pipes, and the city promised to shift back to the old way of doing things?  Suppose the soft water stuck around anyway?  What would you, a resident of that city, do?

There are ways to cope with excessively soft water.   The main points: 1) soap and pure water  form a “partial” solution while the soap forms a "partial" solution with body lipids; 2) minerals in the water affect that solution (determining whether it’s “hard” or “soft” water), and 3) one other condition, i.e., lowered temperature, can offset the problems of rinsing off soapy soft water.

NOTE:  From now on, we will observe this naming convention: "H₂O" will mean "pure water," while we will use the word "water" will mean a mixture of H₂O and the various substances that naturally find their way into it via reactions with CO₂ in the air, i.e., water as we normally experience it.

Water, lipids and solid-liquid solutions: the basics

When you take a shower, soap and H₂O form a chemical solution, commonly described as a “homogeneous mixture.”   One type of solution is formed by the joining of a solid and a liquid in a special way that makes them seem like a single entity.  If you view a solution in a glass, it has a uniform appearance: it is at least somewhat transparent although it may be colored.  The reason this magic happens is that the former liquid and solid substances form intermolecular bonds with each other.  These bonds are much weaker than those that hold atoms together in molecules, but require much less input energy to form.

Not every pair of substances can mix in this special way, however.  The general rule is “Like mixes with like,” where “like” means  the molecules of each are either both “polar” or both “nonpolar.”  Polar molecules have a positive charge on one end and a negative charge on the other because of the unequal sharing of electrons within a molecule that is radially asymmetrical in any plane.  They mix by forming intermolecular bonds in which a negatively charged atom of one polar molecule is attracted to a positively charged atom of another. 

H₂O is a polar molecule: its slightly negatively charged oxygen atom forms an intermolecular bond with two slightly positively charged hydrogen atoms of neighboring water molecules.   Because other polar molecules tend to form this type of intermolecular bond with it, they are called “hydrophilic.”  The tendency of water molecules to form these relatively strong intermolecular bonds (called “hydrogen bonds”) with one another creates surface tension.

Nonpolar molecules mix too, but in a different way:  they neither attract nor repel one another.  Free fatty acids, which are mainly chains of carbon atoms with a couple of hydrogen atoms attached to most of them, can lie very compactly together.   When these chains have the same number of molecules, and therefore the same mass, it all pretty much looks the same, although it isn’t as clear as a solution of polar substances because it’s not as orderly.  They don’t form intermolecular bonds with water as a result, so they’re called “hydrophobic.” This is why oil (less dense than liquid water) dropped on water floats on top of it. 

Lipids, the fatty/waxy substances that you use soap to remove from your skin, are nonpolar.   You can use the purely physical force of water to push lipids out of the way to some degree, but you would still feel pretty grimy because some of the lipids would stay behind in the minute folds of your skin.  So you need a chemical solution: you have to create a solution that joins the water to the lipids so that the water coming off pulls the lipids with it.  But if water (polar) can't form intermolecular bonds with lipids (nonpolar), how is this possible?

How Soap Works

This is accomplished with the miracle of soap!  Soap is a polar-nonpolar hybrid.   The polar part mixes with H₂O, and the nonpolar part mixes with those lipids that you're trying to remove.  So when you remove the H₂O, it pulls the soap after it, which in turn pulls the lipids after it.   This is how it works:

Sodium stearate, a typical soap, has a slightly positively charged sodium atom “head” (the “polar” end) which is attached to an uncharged fatty acid “tail”  (the nonpolar end, called "stearic acid"), which consists mainly of a chain of 18 carbon atoms with two hydrogen atoms branching off each carbon atoms except those at the ends.  The sodium atom head is attracted to the slightly negatively charged oxygen atoms in the polar H₂O, causing those H₂O molecules to crowd onto the exposed surface of the sodium atom, forming a kind of intermolecular bond in a process called "hydration."  This soap-H₂O bond isn't as strong as another kind of intermolecular bond called the "hydrogen bond" that joins H₂O molecules together (by linking a hydrogen atom of one H₂O molecule to the oxygen atom of another).  As a result, the formation of these soap-H₂O causes a reduction in "surface tension;" that's why we call soap a "surfactant." 

Meanwhile, the fatty acid tail mixes with the lipids on your skin.   In fact, the many long, skinny fatty acid tails of the soap molecules fit closely together with the skin lipids in a greasy, hydrophobic glob bounded by the sodium heads, which in turn are joined to H₂O molecules.  These globs, called micelles, are extremely tiny and impossible to see individually.  In fact, these body lipids form a very small part of this water-soap-lipid complex.  This water-and-soap combination can form a film around any air that's introduced into the soap-H₂O combination, producing bubbles that can escape the water surface and even float into the air. 

On the other hand, the strong surface tension of pure H₂O keeps this from happening: you may see fizzing in the water, but the bubbles will never push their way past the water surface at room temperature and one atmosphere of air pressure.   On the other hand, if you heat liquid water to boiling in a pot, put a lid on it, the reduced volume will cause both the air pressure and temperature to increase.  That temperature increase represents an increase in the energy of the water molecules, making them push away from one another and  exert an increasing force on their surroundings (yes, this is what happens during what we call a "phase change," in this case from liquid to gas).   This force overwhelms the water's surface tension by breaking those hydrogen bonds which held the H₂O together in its liquid state.  If you put a transparent lid on the pot, you'll see lots of large bubbles, all the the way up to the lid, as the water turns from liquid to gas form and rises.  This will continue if you turn down the stove to "simmering."  But if you then reduce the air pressure to one atmosphere by removing the lid, the bubbles will quickly disappear.  (Of course, these are steam bubbles, not air bubbles.)

Back to the water-and-soap scenario:  So if you then introduce air, e.g., by lathering, it produces bubbles: air pockets bounded by that film.  So, when you rubbed soap all over yourself, you’ve got a sometimes bubbly, sometimes slippery, H₂O-soap-lipid complex.  All you have to do, then, is get the whole business off.  

That’s when the hard water/soft water issues start kicking in: you need relatively soft water to get bubbles, and sometimes that can be too much of a good thing.  It’s nice to know that all that grease and grime on you is trapped in those bubbles.  But when you try to wash it off and all you get are more bubbles or just slippery stuff, than you’ve got too much of a good thing: your water is too soft – or maybe just too warm.  

Soft water and acidity

Although H₂O is naturally nearly acid-base neutral, it’s usually not found that way in nature, although carbon dioxide-free freshly distilled water (where the distillation takes place in a vacuum) comes very close.   In typical life situations, carbon dioxide (CO₂) finds its way into the water and reacts with it in a way that eventually makes the result acidic, i.e., with some hydrogen ions (H⁺, actually protons) in the mix.  Lots of carbon dioxide in the environment will make it really acidic, hence acid rain in some polluted places.  But there is a natural limit to the acidity of the resulting combination of H₂O, HCO₃^- (hydrogen carbonate ions), CO₃^2- (carbonate ions), H⁺ ions (which react with H₂O molecules to produce covalently-bonded H₃O⁺ ions) and carbonic acid molecules (H₂CO₃) that eventually results (this result being what we still think of as "water.")   This is how it happens:

Let’s say you have pure water, by definition extremely "soft" because there are no minerals in it.  As carbon dioxide enters the picture, it begins to react with the water to form carbonic acid.  But at the same time, some of the carbonic acid is turning back into water and carbon dioxide.  Eventually this system reaches "equilibrium", when the net proportions of all these entities stabilize, with those on the left side predominating.  (This happens because carbonic acid is a "weak" acid, i.e., it's very resistant to breaking down into its component ions, H⁺ being one of them.)  This is the most obvious part of the story:

CO₂ (g) + H₂O (l) ß à H₂CO₃ (aq)

So where does the acid-producing H⁺ come from?  A very small proportion of the H₂CO₃ breaks down to form HCO₃^-  ions and some H⁺ ions via another such equilibrium, and another even smaller number of those HCO₃^-  ions breaks down to form CO₃^2- ions and more H⁺ ions.  (NOTE: “g” means the molecule is in gas or vapor form; “l” means it’s in liquid form; “aq” means it’s dissociated in water, i.e., in a solution with water.  HCO₃^- , CO₃^2-  and H⁺ ions are all "aqueous" in this case.  And what does "aqueous" mean?  It describes that ongoing dynamic behavior of certain molecules in solution in H₂O: for example, "aqueous" H₂CO₃ is part molecular H₂CO₃ (mostly), part HCO₃^- , part HCO₃^2- and part H⁺ (which it becomes part of H₃O⁺ ), and all are solvated:  that means the closest H₂O molecules glom onto them, as many as it takes to cover the surface of these molecules and ions, forming intermolecular bonds with them.

Incidentally, all of those reactions you learned about in biology class, e.g., cellular respiration and DNA transcription, take place in aqueous solutions, the reason you keep hearing about water being absolutely necessary for life to exist!

How does this affect water acidity?  To make a very long story short, if a glass of pure water is left to stand in an unpolluted room overnight, the CO₂ in the air will bring its pH (a measure of its H₃O⁺ ion concentration) down from a close-to-neutral 7.0 to a mildly acidic one of about 5.6 at equilibrium.   The addition of sodium stearate soap, with its fatty acids, lowers the pH even more.

One more point, though: it isn’t the acidity of this new solution of mostly water in itself that makes it “soft.”  The H₃O⁺ ions aren’t really part of the action.  But those other ions (much more numerous than the H₃O⁺ ions) do present certain problems if you try “hardening” the water, which is sometimes done to reverse the acidity.  Here is why:

Very hard water vs. very soft water

Back to the problem of getting off that H₂O-soap-lipid complex on your skin. You’re not greasy any more, just slippery, and that won’t come off unless you turn off the water, get a towel and painstakingly rub it off.  The shower water will burst most of the bubbles and push them and their contents off your body through sheer mechanical force.   But some will remain behind in the tiny folds in your skin that trap this complex.  This happens because some water you add will establish hydrogen bonds with the water molecules on the surface of the soapy film, and you are stuck with a slippery coating.   NOTE:  "Slippery" isn't the same as "soap-free:" the popular expression "squeaky clean," which alludes to friction produced by completely washed and rinsed surfaces, implies the opposite!

You can harden soft water by adding, say, calcium compounds (most likely calcium chloride, CaCl₂, commonly used in water treatment) that readily dissociate in the water, producing calcium ions (Ca²⁺) and chloride ions (Cl^-).   Some of these Ca²⁺ ions react irreversibly with the carbonate ions already existing in soft water to form calcium carbonate:

Ca²⁺ (aq) + CO₃^2- (aq) à CaCO₃ (s)

Calcium carbonate is a precipitate, i.e., a solid (s), insoluble in H₂O and notorious for forming a kind of cement on surfaces where it’s not wanted.  It also interferes with the formation of the water-soap solution by getting in the way: water and soap molecules need to contact each other to form bonds with each other.  But the fewer Ca²⁺ ions and CaCO₃ molecules around, the more the soap and water will bond, the more lipids that the soap is able is to link to the water and the more bubbles will be created if air is introduced with turbulence.  On the other hand, when water pushes the nonreactive, CaCO₃ molecules forward, that CaCO₃ and Ca²⁺ and Cl^- will be able to push other molecules and ions out of the way, making the rinsing process more effective by purely mechanical means.   On the other hand, rinsing this way can remove the soap before it finishes its cleansing function.

This raises the pH of the water by reducing the number of CO₃^2- ions (in the process of creating more CaCO₃ molecules, the number of which is used to calculate water hardness) as shown in the above equation, which in turn lowers the number of H₃O⁺ ions.  This is often the chosen purpose for hardening water: to achieve a neutral pH, thereby putting less stress on our bodies' buffer-producing functions.   Another concern is that harder water provides essential minerals that may be deficient in our diets.   So hard water is likely to be with us regardless of enthusiasm from certain quarters about soft water.

There are other considerations.  Moderate levels of calcium carbonate can form protective thin coatings on lead pipe plumbing which keep the lead from leaching out into the water going through the pipes.  Large amounts can block or narrow the pipes too much. 

So it comes down to this: very soft water can make it hard to rinse yourself; very hard water can make it hard for you to clean yourself.  So somewhere between these extremes you can find an optimum middle ground.  Many municipalities do in fact accomplish this with their water treatment plans.

A proposed remedy: cool things down

You may be wondering at this point: how do washing machines get clothes clean and rinsed with soft water when you can't get anywhere with your own skin, even with a washcloth?  You'll notice that some settings will allow you to wash in warm or hot water, and rinse in cold water.  Unless you're washing cotton clothes that haven't been pre-shrunk (an increasingly common phenomenon these days), picking the "warm" wash and the "cold" rinse is the ticket for soft water, in my humble opinion.  Granted, you probably couldn't stand the cold water the washing machine uses, but tolerably cool water still can make a big difference.

This is why:  In most cases, the solubility of a solid in water, i.e., the proportion of that solid that actually gets into solution with H₂O, decreases as that temperature goes down.  I still remember vividly a story, told  by my chemistry instructor, about how sugar dissolves much more completely in hot tea than in iced tea:  he was annoyed by having to stir the sugar in the latter every time he took a sip of the iced tea to get the sugar evenly (and just temporarily) distributed throughout it.  In the same vein, using cool enough shower water gets the soap to come out of the soap-water solution.

The reason for this is that, for solutions to form, bonds (in this case, intermolecular ones) need to be made, and this requires additional energy.  Heat adds energy, providing the energy for these bonds.  Similarly, cooling the solution removes energy, resulting in the loss of the bonds.  NOTE:  This isn't an iron-clad rule for solid-water solutions: in a small minority of solution types, temperature has the reverse effect, but even then seldom dramatically so.

Conclusion

The ideal solution to the problem of bathing in water that’s too soft is this: 1) put the soap on with warm water and 2) rinse it off with cool water.  The water-soap  bonds (really, between the water and the sodium) formed in the warm environment will break in the cool one.

 References

Chang, R. (2007) Chemistry. 9th ed.

Ball, P. "Water -- an enduring mystery."  Nature 452, 291-292.

"Soap Bubble" (Wikipedia)

"Fatty Acid" (Wikipedia)

http://www.cem.msu.edu/~reusch/VirtualText/lipids.htm

Rzepa, H.S., "Stearic Acid and Sodium Stearate" (Chemistry Dept., Imperial College London)

"Hard Water" (Wikipedia)

"Chemical Polarity" (Wikipedia)

"Carbonic Acid" (Wikipedia)

"Measuring Acid Rain" (Environmental Protection Agency)

"Calcium Carbonate" (Wikipedia)

Kožišek, F., Health significance of drinking water calcium and magnesium

Shakhashiri, B.Z., "Soft Water and Suds" (Chemistry Dept., University of Wisconsin at Madison)

"Calcium Chloride" (Wikipedia)

Wardchem, "Advantages of Using Calcium Chloride for Effluent Treatment" (Express Press Release Maryland)