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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.  Distinctions between points of views of experts (represented by citations) and the opinions of the author are made as clearly as possible.  If you discover any factual errors, please write to Dorothy E. Pugh at Contact Us.

This essay was written mainly as an exercise to apply what I'd learned in General Chemistry I and II to a real-world problem.  I chose the soft water problem because it was an obvious problem in my daily life.  Many hypotheses are set forth in this essay, are undergoing regular revision, and probably sound wild and crazy in places.  Anthropomorphic analogies are used with frequent abandon and speculation runs rampant.

Although this essay has the superficial appearance of a scientific paper, it's really meant to be much more informal.  "I" and "you" are used often to stress the difference between the author's opinions and information drawn from other sources.  I personally don't think these departures from the conventions of the science world hurt this essay's scientific value in themselves, although they may make it harder to read in some ways.  You don't get any tidy conclusions, i.e., in the "Conclusions" section where personal speculation is usually confined in conventional papers;  they're scattered throughout.  What you do see are thought processes leading to hypotheses.

This text has not been reviewed by any professionals in the field, although it may have been read by some.   I have had some difficulty about establishing the facts in this respect in the past and apologize for any misunderstanding this may have caused the reader.  It is difficult to find such a reviewer because so few people are interested in these issues. 

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


Very soft water can be difficult to bathe with because when it's mixed with soap it's hard to remove (but feels "soft!")  However, rinsing in cool water can make this process easier.  This simple personal problem and proposed remedy are described using explanations of ionic bonds, intermolecular bonds, solution formation, stable and unstable colloids, factors affecting solubility, equilibrium in reversible reactions, precipitate formation in irreversible reactions, mechanisms of and environmental influences on water acidity.   The water softener backwash issue is discussed, too, with references to the same concepts. 


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, but we'll discuss that later.)   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.

There are ways for an individual to cope with excessively soft water.   The main points: 1) ideally, 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:  Soap used with very soft water doesn't just feel as though it's not coming off.  The scientific proof: that soapy smell on towels used afterwards, which can become a stench after several days!

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

Aqueous solutions: water, hydration and polar/nonpolar intermolecular bonds

When certain small molecules and H2O are mixed together, they form a chemical solution, commonly described as a “homogeneous mixture” (1).   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 these two substances form intermolecular bonds with each other.  These bonds are much weaker than those that hold atoms together in molecules, but, since they require much less input energy to form, they do so easily.  They also come apart easily, but at any given moment, a lot of these bonds exist.

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 (2).  Polar molecules have, in the simplest cases, a partial positive charge on one end and a partial negative charge on the other because of the unequal sharing of electrons within a molecule that arise when the magnitude of difference in electronegativity between atoms (of two different elements) which share a covalent bond (one in which bonding electrons are shared by both atoms) is above a certain minimum value: the more electronegative atom exerts more force on its valence electrons, and therefore on those shared in that covalent bond.   As a result, the more electronegative atom has a partial negative charge, while the less electronegative one has a partial positive charge.  Generally speaking, radially asymmetrical molecules are polar, but there are some important exceptions, i.e., carbon (C) and hydrogen (H) have very similar electronegativity values and therefore hydrocarbons, i.e., all molecules composed of combinations of just C and H atoms, have such small polarity that they are effectively nonpolar.  Molecules composed completely of C atoms, i.e., the graphite making up pencil "lead," are completely nonpolar.

Substances composed of polar molecules mix by forming 1) relatively strong intermolecular bonds called dipole-dipole bonds in which a partially negatively charged atom of one polar molecule is attracted to a partially positively charged atom of another and 2) by weaker London dispersion forces (2a), which are actually temporary dipole-dipole bonds that occur when electrons briefly cluster together relatively near the nucleus of an atom as a result of the constantly shifting electron positions within atoms and movement of those atoms with respect to one another: these clusters create short-term or induced dipole-dipole bonds.  On the other hand, substances composed of nonpolar molecules mix by forming only those London dispersion forces.

For example, H2O  (3) is a polar molecule: its slightly negatively charged oxygen atom forms an intermolecular bond with two slightly positively charged hydrogen atoms of neighboring H2O molecules; this occurs because the more massive oxygen atoms pull the hydrogen electrons toward them.   Because other polar atoms (such as sodium or potassium) or molecules tend to form this type of intermolecular bond with it, they are called “hydrophilic.”  The tendency of water molecules to form a relatively strong type of dipole-dipole bond called a “hydrogen bond” with one another creates surface tension.  On the other hand, when ionic compounds  are added to H2O, they break up into ions which are hydrated, meaning that H2O molecules glom onto the surface of these ions, covering them; the average number of these H2O molecules covering a sodium atom, i.e., the hydration number for sodium, is 4.6, with 5 the most typical (4).  Hydrated sodium is therefore (usually) represented as Na•5H2O.

For example, salt or sodium chloride (NaCl) is an ionic compound (5) which separates into its component ions (Na+ and Cl-)  in H2O as hydration of these ions takes place.  Sodium and chlorine are especially susceptible to this, i.e., are especially soluble, because they each have a lone electron in their outer shells that is easily lost to other atoms, which is reflected in these atoms' low first ionization potential (6), the energy (force) required to pull this lone electron away from these atoms.  Because both sodium and potassium (ionized as K +) have these characteristics, salt water is an excellent conductor of electricity.   For this reason, sodium and potassium are known as electrolytes (7) and perform important functions in maintaining H2O balance across cell walls in the human body.  NOTE: This solubility isn't absolute; for instance, we learned early that if you left a string in a glass of salt water that crystals would eventually form on it.  Of course, the glass had to be left undisturbed for a long time and the lower the temperature, the better.

Nonpolar molecules mix too, but in a different way.  Free fatty acids (8), which are mainly chains of carbon atoms with a couple of hydrogen atoms attached to all or most of them, can lie very compactly together, which makes the overall London dispersion forces among them stronger.   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. 

Lipids (9), the fatty/waxy substances that you use soap to remove from your skin, are nonpolar.   Lipids are a rather heterogeneous group of substances defined by their nonpolarity, which makes them hydrophobic.  When most people use the term, I think, they are thinking of substances naturally found in the human body that play a key role in its function.  Fats (10) are lipids that have component fatty acids; in fact what we commonly call "fat" is really the triglyceride molecule, composed of glycerol and three fatty acids. 

Bigger, Messier, Murkier but Eventually Cleaner: The Colloid and Soap

When certain large lipid molecules and H2O are brought together, they form a chemical colloid (11), a kind of “heterogeneous mixture,” a relatively opaque, rather lumpy substance that is very unstable, i.e., the clumps of lipid are always breaking up and moving around.  The dispersion forces that allow lipids to mix together just aren't that strong.  But certain ions covering the outside of these clumps of lipid can stabilize them by keeping the lipid inside and the  H2O outside.  And you ask, why would I care whether a colloid is stabilized?  After all, I enjoy eating eggs and drinking (maybe) homogenized whole milk (well, buttermilk), which are definitely unstabilized colloids! 

Everything is different if you want to get that greasy stuff that inevitably collects on your skin off. 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 to your problem: 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? This is accomplished with the miracle of soap!  Soap (12) is a polar-nonpolar hybrid.   The polar part mixes with H2O (forming one type of solution), and the nonpolar part mixes with those lipids that you're trying to remove (forming another type of solution).  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:

The organic molecule sodium stearate (IUPAC name sodium octadecanoate, formula C18H35NaO2) (13), a typical soap, has a slightly positively charged sodium atom “head” (the “polar” end) which is chemically bonded to the end of an fatty acid “tail” where it ends with a slightly negatively charged oxygen atom. This fatty acid in its standalone state is called "stearic acid" (formula C18H36O2) and consists mainly of a chain of 18 carbon atoms (C) with two hydrogen atoms (H) branching off each carbon atom except those at the ends, i.e., three from the carbon at one end and one at the cargon at the other end, bonded to two oxygen atoms (O), one of which is bonded to a hydrogen atom (are you still with me?)   Sodium stearate is formed when a sodium atom (Na) comes along, bumps the hydrogen atom on the end out of the way and takes its place.   The slightly positively charged sodium atom head is attracted in turn via dipole-dipole intermolecular forces 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 bond weaker than an atomic 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 bonds causes a reduction in "surface tension;" that's why we call a "pure" soap such as sodium stearate a "surfactant."  Commercial soaps, of course, contain many more ingredients intended to make the effect on the skin gentler.

Meanwhile, the stearate fatty acid tail mixes with the lipids on your skin.  Actually, to be entirely fair, stearate is not 100% polar because of the oxygens on the end.  But because of that long, symmetrical hydrocarbon chain, it behaves much more like a purely nonpolar molecule than a purely polar one.  As it happens, 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 hydrophilic sodium heads, which in turn face outward from the middle of that glob and are joined to H2O molecules.  These globs, called micelles (9), are extremely tiny and impossible to see individually; this is how soap creates a stabilized colloid, a process known as emulsification.  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 (14). 

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 (and soap, i.e., a polar-nonpolar hybrid) 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.  

Naturally 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), CO32- (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" (15).  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 CO32- 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- , CO32-  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 HCO32- and part H+ (which joins H2O to become H3O+ ), and all are hydrated:  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 (in the aqueous diffusion layer of the cell membrane and inside the cell), 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 H2O were left to stand in an unpolluted room overnight, the CO2 in the air would 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 (16).   This result would be the chemical equivalent of "normal" rainwater, which is actually infinitely soft water.

One more point, though: it isn’t the acidity of this new solution of mostly water in itself that makes it “soft,” although there is an indirect relationship.  The H3O+ ions aren’t really part of the action.  But some of those other ions mentioned above which contribute to the formation of 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-) (17).   Some of these Ca2+ ions react irreversibly with the carbonate ions already existing in soft water to form calcium carbonate:

Ca2+ (aq) + CO32- (aq) à CaCO3 (s)

Calcium carbonate (18) is a precipitate, i.e., a solid (s), insoluble in H2O (at least at "normal" temperatures) and notorious for forming a kind of cement called "limescale" (19) 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.  Meanwhile, the calcium ions (Ca2+) are bumping the sodium atoms (Na) in the sodium stearate out of the way and taking their place to form calcium stearate, which forms "soap scum" if it collects on the bottom of the bathtub or shower stall.  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.  The flip side of this is that the hard water can remove the soap before it finishes its cleansing function.

This raises the pH of the water by reducing the number of CO32- ions in the process of creating more CaCO3 molecules, the indirect measurement of which is used to calculate water hardness (20)  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 (21).   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 CaCO3 can form protective thin coatings on lead pipe plumbing which keep the lead from leaching out into the (acidic) water going through the pipes.  Large amounts can block or narrow the pipes too much, however.   Also, "soap scum" can form, providing a challenge for those cleaning bathtubs and sinks; it's formed by the Ca2+ ions that haven't formed CaCO3 and the stearate from the soap: the Ca2+ bumps the Na+  out of the way, producing calcium stearate (22), which makes a softer deposit which can be removed by mechanical means (and some sweat!)  With less sodium stearate, the active ingredient in your soap product, you get less cleansing action.  As a result, some Na+ is released into the water, with which it forms a solution via hydration, and it's carried out into the sewer system eventually.  This amount is very little per shower, at least when compared to that released by water softeners.  The next section will explain this, as well as why Ca2+ has more "pull" than Na+ does.

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.  According to one Canadian government agency, the ideal tradeoff is 80 to 100 mg/L of CaCO3 (23).  Many municipalities apparently do in fact accomplish this with their water treatment plans.

Water Hardness, Water Acidity and Water Softeners

No discussion of public water softness is complete without a discussion of water softeners.  The public water quality issue is actually much more complex because 1) some water is naturally hard, i.e., already containing metal ions, especially calcium (Ca2+) and magnesium (Mg2+) ions and 2) water softeners may arrive with your water, removing these ions and sometimes causing undesirable side effects.   This is not to say that municipal water treatment plants use water softeners (they don't); however, some of their input wastewater has been treated with water softeners by residential and commercial customers and water treatment at the plant necessarily involves removal of enough of these contaminants to meet any relevant state and federal regulations at the minimum.   For example, the city of Paso Robles, CA was motivated by the size and expense of this task to devote a page on its website to educating its residents about how to reduce the "discharge of salt brines into the wastewater collection system" by moderating and altering their use of water softeners (24).  The implication here is that, in the absence of sufficient regulations or oversight, a less conscientious city might decide to cut its water treatment costs by skimping on the processing of these discharges, thereby softening the water it provides its citizens.

Where does all this salt come from?  One method of water softening is ion exchange (25), where introduced cations (elements that tend to lose electrons rather than gain them in chemical reactions) with low ionization potential displace others with greater ionization potential.  Introduced Group I sodium (Na+) and/or potassium (K+) cations replace Group II Ca2+ and Mg2+  cations. This is an equation describing a reaction produced by a water softener designed to remove calcium sulfate (CaSO4), a common well water mineral, using the salt sodium carbonate (Na2CO3) (20):

Na2CO3 + CaSO4 → Na2SO4 + CaCO3

 This is made clearer by specifying the phases of the reactants and products:

Na2CO3 (aq)+ CaSO4 (aq) → Na2SO4 (aq) + CaCO3 (s)

In reality, this is just a bunch of ions except for the CaCO3.  These ions may have been joined as "solid" molecules when they entered this system, but they really aren't distinguishable as "sodium carbonate" (Na2CO3) or "calcium sulfate" (CaSO4) in the water: they are really just sodium (Na+), calcium (Ca2+), sulfate (SO42-) and carbonate (CO32-) ions, each covered with H2O molecules held to them by dipole-dipole forces.  As you can probably guess, this process takes place before the water comes out of the tap, its solid precipitated CaCO3 disposed of and lots of Na+ and SO42- kicking around.  Oh yes, there's a "resin" involved, full of beads with (in this particular example) negatively charged carbonate ions (CO32-) that bond with the Na+ ions when that resin is "washed" in a salt-rich solution created by dropping sodium carbonate pellets into it, where they sit until each Ca2+ in the water finds its way to the resin, bumps the Na+ out of the way and forms a genuine chemical bond via the above reaction.  The harder the water, the more sodium is released into it, and the bigger the water disposal problem.

How does this work on an atomic level?   Group I metals, e.g. Na+,  have the lowest first ionization potentials because they have just one electron distant enough from the nucleus to be lost in aqueous solution (creating an ion represented with a single "+") and therefore form relatively weak bonds with anions, i.e., are relatively "soluble" compounds with anions (those elements that tend to gain electrons in chemical reactions).  On the other hand, Group II cations, e.g. Ca2+, which tend to lose two electrons in chemical reactions (creating an ion represented with "2+"), hold onto them and onto the anions that take their electrons away more firmly.   The less the ionization potential, the smaller the chance that an "insoluble" (relatively speaking) compound (such as CaCO3) will form.   That explains why sodium and potassium ions are preferred to calcium and magnesium, and introduced to replace them: they are less likely to form troublesome precipitates.  

There is concern about the safety of the ion exchange water-softening process (26,27) because extra sodium and potassium intake in drinking water can upset one's electrolyte balance and (in the case of sodium) aggravate high blood pressure.  In California, one city is taking the first steps toward restricting the residential and commercial use of sodium-based ion exchange water-softeners to reduce groundwater salinity (28).  And one North Carolinian claims that some such sodium-softened water kills houseplants (29). 

There's a funny relationship between the hardness of household water and the pH of the rainwater (by definition, mineral-free and therefore extremely soft) that it was originally.  We know that the main acidity factor, CO2, produces carbonate (CO32-) ions, which may bind irreversibly with calcium ions (Ca2+)  to produce CaCO3.  And we know that a popular indirect measurement of water hardness is the CaCO3 concentration (mg of precipitated CaCO3 per L of water) in that water (20).  So if there's an excess of CaCO3 in the water after the CO32- has all been "used up," that calcium won't be figured into the hardness measurement.   So what this measure is really giving us is nuisance hardness.  So the higher local CO2 emissions are, the more problematic a certain degree of existing water hardness is and, by extension, the greater the motivation of affected people to use water softeners.

There is considerable regional variation within the US for acidity of rainwater and water hardness, incidentally.  One extreme case is part of Ohio, where water hardness is high, i.e., over 181 CaCO3 mg/L, while the pH of rainwater is low, in the 4.1-4.3 range.   So if people in this part of the country use enough water softeners to achieve complete softness, that would mean a lot of salinity in their wastewater.  On the other hand, North Carolina water is very soft, i.e., less than 60 CaCO3 mg/L (below the Canadian recommended range of 80 and 100 mg/L discussed above) and the pH of North Carolina Piedmont area rainwater is moderately acidic, in the 4.5-4.7 range (30).  Yet a trip to my local hardware store invariably means getting to see piles of large bags of table salt (NaCl) for sale for use in water softeners!

 Another method of water softening is adding EDTA (31) , an amino acid that performs chelation on all metal ions, i.e., chemically binds with them so it can carry them out of the body.  EDTA is not absorbed by the body, but is indirectly capable of causing environmental pollution: disposing of it safely requires special measures that are prescribed by government organizations, since it needs to be treated with sodium to release the minerals it is bonded to.   Hazards (32) are known in detail.  The Environmental Protection Agency (EPA) has apparently not set any EDTA water quality standards (33).

It's worth noting that the only use of EDTA and other chelating agents in standard allopathic medical practice is for intravenous treatment of acute heavy metal poisoning in adults (34).  However, its use is promoted vigorously online as "chelation therapy" for a variety of ailments (35).   It seems to be common sense to me that the benefits of its FDA-approved uses would not carry over to those other uses because 1) chelating agents each remove a wide range of minerals from the body, 2) in the case of acute heavy metal poisoning, most of the metal removed would be the target toxic metal and 3) this therapy would be short-term and the patient would suffer only temporary (and presumably moderate) essential mineral deficiencies as a result.

How Does Water Softener Saline Brine Discharge Stress Public Water Treatment Systems?

There really isn't a standard wastewater processing step designed with salt brine discharges in mind.  So these discharges do present problems during an early stage of wastewater processing, when dense solids (sludge) are allowed to settle on the bottom of large containers known as "sedimentation tanks," while the less dense lipid-containing compounds (scum) rise to the water surface.  Scrapers on the bottom of the tanks deliver the sludge to a hopper through which it's selectively removed.   The remaining water is sent to drainfields that are one to three feet above the groundwater table to be treated, where microbes in the soil break down the remaining chemicals as the water soaks down through the soil, eventually entering the groundwater.  When the sludge completes treatment, it is either compressed into pellets and sold as fertilizer or delivered to landfills.  Any sodium entering the system, then, is returned to the environment (36). 

 The problem with salt brine (hydrated sodium, i.e., most typically Na•5H2O) is that it can interfere with the microorganism-mediated breakdown process because it is toxic to these microbes, upsetting their osmotic balance: it creates a hypertonic solution outside them, causing H2O to flow out of their cells via diffusion (37).  If sludge sodium is not disposed of properly, that sodium and its companion anions can also cause calcium- and magnesium-displacement problems in the soil.   

One Michigan business providing septic tank systems to those without access to municipal wastewater treatment classifies "backwash from Water Softener regeneration" in the "Chemicals & Toxins" category that "kill the microbes necessary for the biological treatment (of wastewater) to occur" (38).

The Impact of Water Softener Saline Discharge on Individual Households

Alas, you won't see any footnotes here, because what is all theoretical (really, hypothetical) on my part.  What do a bunch of sodium ions (Na+) do when they enter your water supply and you don't have a water softener system, complete with a resin to collect Group II metal ions such as calcium and magnesium?  This is what I suspect happens:  the sodium ions, on entering the water, become hydrated (remember, that means each gets "covered" by about 5 H2O molecules).  These units then get in the way when calcium and carbonate ions try to get together to produce CaCO3 (just as CaCO3 gets in the way of soap joining water as described above, thereby giving hard water its bad reputation for interfering with cleaning, an issue, of course, apart from that of producing limescale).  With less CaCO3, the water would be softer.  It's possible that the water is the softest during the resin-regeneration process when the water is flooded with Na+


Proposed (Personal) Remedy #1:  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 (1).  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.

It is generally accepted that hard water shortens the life of washing machines and clothes alike because of the CaCO3 deposits (and those of some other minerals) it leaves behind (25).

Proposed (Personal) Remedy #2:  Keep Hard and Soft Water in Separate Systems

One set of guidelines proposes having separate water systems: hard water for drinking and soft water for washing clothes (39).   There are some significant drawbacks to implementing this plan, in addition to the effort and expense of installing and maintaining personal and commercial water softeners.   As stated above, using standard water softening methods involves adding sodium to the water.

Conclusion:  Community Remedies

If you have water that's too soft for your tastes, it probably is just right for your washing machine and you can cope to some degree by rinsing yourself in cool water.  But the soft water/hard water issue is actually quite complicated because of this discrepancy, and is taking on major environmental dimensions.  We have yet to find a solution that provides optimal comfort, minimizes wear on washing machines and is safe for the environment.

One thing we can do now as a group is to make sure that people have a genuine choice about whether water-softening chemicals get into their water.  Public wastewater treatment plant personnel can be open about how big the water softener output treatment problem is and to make an effort to educate customers about how water softener use can contaminate groundwater and what they can do to minimize the harm their households inflict on it.  Part of that education should include informing them about the actual hardness of the water in their region and, by implication, how much softening their water really needs to produce a significant reduction in wear on their washing machines and clothes.   Water acidity is a factor, too:  if customers know that their local rainwater is already too acid and that softening the water makes it even more so, they may be motivated to cut back on water softeners.  There are, of course, kits available to test one's own water.

If a lot of sodium leaves the water softener system to enter the sewer system, that suggests it's wasted.  Modifying available technology to recycle the sodium would help.  One obvious method would be to reduce the water temperature to near freezing, but that would be expensive.  Or we could go back to what we learned in elementary school: have the water stand in a tank with strings hanging down in it so that salt crystals form and recycle the salt crystals afterwards.  However, this system would work only if softened water were used only in washing machines and dishwashers, which use much less water than people do taking showers.

Part of the municipal solution might involve charging more for wastewater processing, with the individual household cost being set according to the sodium concentration in that household's (or, more realistically, neighborhood's) wastewater.  But this would be more effective if customers can be convinced that their need for water softeners is less than they had previously assumed.

In the long term, we as a nation may need to come up with new water-softening technologies that take minerals out of the water instead of putting more in it.  But in some communities, mine among them, this really doesn't seem to be a pressing problem. 

Finally, reducing CO2 emissions will reduce the need for water softening since it will reduce the amount of CaCO3 and calcium stearate produced.  But that's obviously a very long-term goal!


1. Solution (Wikipedia)

2. "Chemical Polarity" (Wikipedia)

2a. Purdue University (General Chemistry Help),"London Dispersion Forces"

3. "Water (Molecule)" (Wikipedia)

4. Rempe, S.A. and L.R. Pratt (2007) The hydration number of Na+ in liquid water, 14th Symposium on Thermophysical Properties, V. 183-184, pp. 121-132, Boulder, CO 

5. "Ionic_compound" (Wikipedia)

6. "Ionization_potential" (Wikipedia)

7. "Electrolyte" (Wikipedia)

8. "Fatty Acid" (Wikipedia)

9. "Lipid" (Wikipedia)

10. "Fat" (Wikipedia)

11. "Colloid" (Wikipedia)

12. "Soap" (Wikipedia)

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

14. "Soap Bubble" (Wikipedia)

"Carbonic Acid" (Wikipedia)

"Measuring Acid Rain" (Environmental Protection Agency)

17. "Calcium Chloride" (Wikipedia)

18. "Calcium Carbonate" (Wikipedia)

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

20. "Hard Water" (Wikipedia)

21. "Limescale" (Wikipedia)

22. Technology to enable the removal of calcium carbonate deposits from surfaces at elevated pH’s (flintbox.com)

23. Health Canada: Environmental and Workplace Health: Guidelines For Canadian Drinking Water Quality - Supporting Documents - Hardness

24. City of Paso Robles, "Water Softener Issues"

25. "Ion Exchange" (Wikipedia)

26.  "Water Softening" (Wikipedia)

27. Farmington's Municipal Water

28. Murphy, M. "Dixon may back water softener bill," 7/20/2008. Vacaville, CA: The Reporter

29. http://www.city-data.com/forum/north-carolina/95070-hard-water-2.html

30. "It's raining, it's pouring: A chemical analysis of rainwater" (ScienceBuddies.com)

31. Mallinckrodt and Baker (2007) "Material Safety Data Sheet: EDTA, Iron (III) Derivative, Sodium Salt (13% Iron)"

32. Dow Chemical Company (2008) "Product Safety Assessment (PSA): Salts of Ethylenediaminetetraacetic Acid (EDTA)"

33. http://www.pesticideinfo.org/Detail_Chemical.jsp?Rec_Id=PC36146

34. Barile, F.A. (online) Clinical Toxicology: Principles and Mechanisms, p. 310. CRC Press.

35. Green, S. (2007) Chelation Therapy: Unproven Claims and Unsound Theories

36. http://en.wikipedia.org/wiki/Sewage_treatment

37. http://www.tvdsb.on.ca/WESTMIN/science/sbi3a1/Cells/Osmosis.htm

Ryan's Excavating (MI), "Private Onsite Wastewater Treatment Systems"

39. Guidelines for Canadian Drinking Water Quality



Other Sources of Information








State of Minnesota, "Sulfate in Well Water"


Science, Vol 214, Issue 4526, 1241-1244 Copyright © 1981 by American Association for the Advancement of Science Intestinal diffusion barrier: unstirred water layer or membrane surface mucous coat? KW Smithson, DB Millar, LR Jacobs, and GM Gray

SEGURA-BONO M. J. ; GARRIGUES T. M. ; MERINO V. ; BERMEJO M. V. ; 1994, vol. 107, no2, pp. 159-166 Compared effects of synthetic and natural bile acid surfactants on xenobiotic absorption. III: studies with mixed micelles 






Chang, R. (2007) Chemistry. 9th ed. /em>

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



"Hard Water" (Wikipedia)


"Calcium Carbonate" (Wikipedia)

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

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


Copyright © 2008 by Dorothy E. Pugh