Category Archives: Chemistry

The chemistry of lead in drinking water

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Our county, like many in the US and Canada, is served by thousands of miles of lead pipes. Some of these are the property of the government, others sit beneath our homes and are the property of the home-owner. These pipes are usually safe, but sometimes poison us. There is also problem of lead-tin solder. It was used universally to connect iron and copper pipes until it was outlawed in 1986. After years of sitting quietly, this lead caused a poisoning crisis in DC in 2004, and in Flint in 2015-16. Last month my town, Oak Park, registered dangerous lead levels in the drinking water. In an attempt to help, please find the following summary of the relevant lead chemistry. Maybe people in my town, or in other towns, will find some clue here to what’s going on, and what they can do to fix it.

lead pipes showing the three oxides: brown, yellow, and red, PbO2, PbO, and Pb2O3.

Left to itself, lead and solder pipe could be safe; lead is not soluble in clean water. But, if the water becomes corrosive, as happens every now and again, the lead becomes oxidized to one of several compounds that are soluble. These oxides are the main route of poisoning; they can present serious health issues including slow development, joint and muscle pain, memory issues, vomiting, and death. The legal limit for lead content in US drinking water is 15 ppb, a level that is far below that associated with any of the above. The solubility of PbO, lead II oxide, is more than 1000 times this limit 0.017 g/L, or 17,000 ppb. At this concentration serious health issues will show up.

PbO is the yellow lead oxide shown in the center of the figure above, right; the other pipes show other oxides, that are less-soluble, and thus less dangerous. Yellow lead oxide and red lead oxides on the right were used as paint colors until well into the 20th century. Red lead oxide is fairly neutral, but yellow PbO is a base; its solubility is strongly dependent on the PH of the water. In neutral water, its solution can be described by the following reaction.

PbO + H2O(l) –> Pb2+(aq) + 2 OH–(aq).

In high pH water (basic water), there are many OH–(aq) ions, and the solubility is lower. In low pH, acidic water the solubility is even higher. For every 1 point of lower pH the lowubility increases by a factor of 10, for every 1 point of higher pH, it decreases by a factor of ten. In most of our county, the water is slightly basic, about pH 8. It also helps that our water contains carbonate. Yellow lead forms basic lead carbonate, 2PbCO3·Pb(OH)2, the white lead that was used in paint and cosmetics. Its solubliity is lower than that of PbO, 110 ppb, in pure water, or within legal limit in water of pH 8. If you eat white lead, though, it reacts with stomach acid, pH 2, and becomes quite soluble and deadly. Remember, each number here is a factor of ten.

A main reason lead levels a very low today are essentially zero, even in homes with lead solder or pipe, involves involves the interaction with hypochlorite. Most water systems add hypochlorite to kill bacteria (germs) in the water. A side benefit is significant removal of lead ion, Pb2+(aq).

Pb2+(aq) + 2 ClO–(aq) –> Pb(ClO)2(s). 

Any dissolved lead reacts with some hypochlorite ion reacts to form insoluble lead hypochlorite. Lead hypochlorite can slowly convert to Lead IV oxide — the brown pyrophilic form of lead shown on the left pipe in the figure above. This oxide is insoluble. Alkaline waters favor this reaction, decreasing solubility, but unlike with PbO, highly alkaline waters provide no significant advantage.

PbClO+(aq) + H2O(l) –> PbO2(s) + 2 H+(aq) + Cl–(aq)

Lead IV oxide, PbO2 was used in old-fashioned matches; it reacts violently with phosphorus or sulfur. People were sometimes poisoned by sucking on these matches. In the stomach, or the presence of acidic drinking water, PbO2 is decomposed forming soluble PbO:

PbO2(s) +2 H+(aq) + 2 e– –> PbO(s) + H2O(l).

You may wonder at the presence of the two electrons in the reaction above. A common source in water systems is the oxidation of sulphite:

SO3-2(aq)–> SO4-2(aq) + 2 e–.

The presence of sulphite in the water means that hypochlorite is removed.

ClO–(aq) + 2 H+(aq) + 2 e– —> Cl–(aq) + H2O(l).

Removal of hypochlorite can present a serious danger, in part because the PbO2(s) slowly reverts to PbO and becomes soluble, but mostly because bacteria start multiplying. In the Flint crisis of 2016, and in a previous crisis in Washington DC, the main problem, in my opinion was a lack of hypochlorite addition. The lead crisis was preceded by an uptick in legionnaires disease; It killed 12 people in Flint in 2014 and 2015, and 87 were sickened, all before the lead crisis. Eventually, it was the rise of legionaries disease that alerted water officials in Virginia that there was something seriously wrong in Flint. Most folks were unaware because Flint water inspectors seem to have been fudging the lead numbers to make things look better.

Most US systems add phosphate to remove lead from the water. Flint water folks could have stopped the lead crisis, but not the legionnaires, by adding more phosphate. Lead phosphate solubility is 14 ppb at 20°C, and my suspicion is that this is the reason that the legal limit in the US is 15 ppb. Regulators chose 15 ppb, I suspect, not for health reasons, but because the target could be met easily through the addition of phosphate. Some water systems in the US and Canada disinfect with chloramine, not hypochlorite, and these systems rely entirely on phosphate to keep lead levels down. Excess phosphate is used in Canada to lower lead levels below 10 ppb. It works better on systems with hypochlorite.

Chloramine is formed by reacting hypochlorite with ammonia. It may be safer than hypochlorite in terms of chlorite reaction products, a real problem when the water source is polluted. But chloramine is not safe. It sickened 72 soldiers, 36 male and 36 female in 1998. They’d used ammonia and bleach for a “cleaning party” on successive days. Here’s a report and first aid instructions for the poisoning. That switching to chloramine can expose people to lead is called “the chloramine catch”.

Unlike PbO, PbO2 is a weak acid. PbO2 and PbO can react to form red lead, PbO•PbO2(s), the red stuff on the pipe at right in the picture above. Red lead can react with rust to form iron plumbable, an insoluble corrosion resister. A simple version is:

PbO•PbO2(s) + Fe2O3(s) —> 2FePbO3(s).

This reaction is the basis of red-lead, anti-rust compounds. Iron plumbable is considered to be completely insoluble in water, but like PbO it is soluble in acid. Bottom line, slightly basic water is good, as are hypochlorite in moderation, and phosphate.

Robert Buxbaum, November 18, 2019. I ran for water commissioner, and might run again. Even without being water commissioner, I’ll be happy to lend my expertise, for free, to any Michigan town or county that is not too far from my home.

Why concrete cracks and why sealing is worthwhile

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The oil tanker Palo Alto is one of several major ships made with concrete hulls.

The oil tanker Palo Alto is one of several major ships made with concrete hulls.

Modern concrete is a wonderful construction material. Major buildings are constructed of it, and major dams, and even some ships. But under the wrong circumstances, concrete has a surprising tendency to crack and fail. I thought I’d explain why that happens and what you can do about it. Concrete does not have to crack easily; ancient concrete didn’t and military or ship concrete doesn’t today. A lot of the fault lies in the use of cheap concrete — concrete with lots of filler — and with the cheap way that concrete is laid. First off, the major components of modern concrete are pretty uniform: sand and rock, Portland cement powder (made from cooked limestone, mostly), water, air, and sometimes ash. The cement component is what holds it all together — cements it together as it were — but it is not the majority of even the strongest concretes. The formula of cement has changed too, but the cement is not generally the problem. It doesn’t necessarily stick well to the rock or sand component of concrete (It sticks far better to itself) but it sticks well enough that spoliation, isn’t usually a problem by itself.

What causes problem is that the strength of concrete is strongly affected (decreased) by having lots of sand, aggregate and water. The concrete used in sidewalks is as cheap as possible, with lots of sand and aggregate. Highway and wall concrete has less sand and aggregate, and is stronger. Military and ship concrete has little sand, and is quite a lot stronger. The lowest grade, used in sidewalks, is M5, a term that refers to its compressive strength: 5 Mega Pascals. Pascals are European (Standard International) units of pressure and of strength. One Pascal is one Newton per square meter (Here’ a joke about Pascal units). In US (English) units, 5 MPa is 50 atm or 750 psi.

Ratios for concrete mixes of different strength.

Ratios for concrete mixes of different strength; the numbers I use are double these because these numbers don’t include water; that’s my “1”.

The ratio of dry ingredients in various concretes is shown at right. For M5, and including water, the ratio is 1 2 10 20. That is to say there is one part water, two parts cement, 10 parts sand, and 20 parts stone-aggregate (all these by weight). Added to this is 2-3% air, by volume, or nearly as much air as water. At least these are the target ratios; it sometimes happens that extra air and water are added to a concrete mix by greedy or rushed contractors. It’s sometimes done to save money, but more often because the job ran late. The more the mixer turns the more air gets added. If it turns too long there is extra air. It the job runs late, workers will have to add extra water too because the concrete starts hardening. I you see workers hosing down wet concrete as it comes from the truck, this is why. As you might expect, extra air and water decrease the strength of the product. M-10 and M-20 concrete have less sand, stone, and water as a proportion to cement. The result is 10 MPa or 20 MPa strength respectively.

A good on-site inspector is needed to keep the crew from adding too much water. Some water is needed for the polymerization (setting) of the concrete. The rest is excess, and when it evaporates, it leaves voids that are similar to the voids created by having air mix in. It is not uncommon to find 6% voids, in commercial concrete. This is to say that, after the water evaporates, the concrete contains about as much void as cement by volume. To get a sense of how much void space is in the normal concrete outside your house, go outside to a piece of old concrete (10 years old at least) on a hot, dry day, and pour out a cup of water. You will hear a hiss as the water absorbs, and you will see bubbles come out as the water goes in. It used to be common for cities to send inspectors to measuring the void content of the wet (and dry) concrete by a technique called “pycnometry” (that’s Greek for density measurement). I’ve not seen a local city do this in years, but don’t know why. An industrial pycnometer is shown below.

Pyncnometer used for concrete. I don't see these in use much any more.

Pycnometer used for concrete. I don’t see these in use much any more.

One of the main reason that concrete fails has to do with differential expansion, thermal stress, a concept I dealt with some years ago when figuring out how cold it had to be to freeze the balls off of a brass monkey. As an example of the temperature change to destroy M5, consider that the thermal expansion of cement is roughly 1 x 10-5/ °F or 1.8 x10-5/°C. This is to say that a 1 meter slab of cement that is heated or cooled by 100°F will expand or shrink by 10-3 m respectively; 100 x 1×10-5 = 10-3. This is a fairly large thermal expansion coefficient, as these things go. It would not cause stress-failure except that sand and rock have a smaller thermal expansion coefficients, about 0.6×10-5 — barely more than half the value for cement. Consider now what happens to concrete that s poured in the summer when it is 80°F out, and where the concrete heats up 100°F on setting (cement setting releases heat). Now lets come back in winter when it’s 0°F. This is a total of 100°F of temperature change. The differential expansion is 0.4 x 10-5/°F x 100°F =  4 x10-4 meter/meter = 4 x10-4 inch/inch.

The force created by this differential expansion is the elastic modulus of the cement times the relative change in expansion. The elastic modulus for typical cement is 20 GPa or, in English units, 3 million psi. This is to say that, if you had a column of cement (not concrete), one psi of force would compress it by 1/3,000,000. The differential expansion we calculated, cement vs sand and stone is 4×10-4 ; this much expansion times the elastic modulus, 3,000,000 = 1200 psi. Now look at the strength of the M-5 cement; it’s only 750 psi. When M-5 concrete is exposed to these conditions it will not survive. M-10 will fail on its own, from the temperature change, without any help needed from heavy traffic. You’d really like to see cities check the concrete, but I’ve seen little evidence that they do.

Water makes things worse, and not only because it creates voids when it evaporates. Water also messes up the polymerization reaction of the cement. Basic, fast setting cement is mostly Ca3SiO5

2Ca3SiO5 + 6 H2O –> 3Ca0SiO2•H2O +3Ca(OH)2•H2O.

The former of these, 3Ca0SiO2•H2O, forms something of a polymer. Monomer units of SiO4 are linked directly or by partially hydrated CaO linkages. Add too much water and the polymeric linkages are weakened or do not form at all. Over time the Ca(OH)2 can drain away or react with  CO2 in the air to form chalk.

concrete strength versus-curing time. Slow curing of damp concrete helps; fast dry hurts. Carbonate formation adds little or no strength. Jehan Elsamni 2011.

Portland limestone cement strength versus curing time. Slow curing and damp helps; fast dry hurts. Carbonate formation adds little or no strength. Jehan Elsamni 2011.

Ca(OH)2 + CO2 → CaCO3 + H2O

Sorry to say, the chalk adds little or no strength, as the graph at right shows. Concrete made with too much water isn’t very strong at all, and it gets no stronger when dried in air. Hardening goes on for some weeks after pouring, and this is the reason you don’t drive on 1 too 2 day old concrete. Driving on weak concrete can cause cracks that would not form if you waited.

You might think to make better concrete by pouring concrete in the cold, but pouring in the cold makes things worse. Cold poured cement will expand the summer and the cement will detach from the sand and stone. Ideally, pouring should be in spring or fall, when the temperature is moderate, 40-60°F. Any crack that develops grows by a mechanism called Rayleigh crack growth, described here. Basically, once a crack starts, it concentrates the fracture forces, and any wiggling of the concrete makes the crack grow faster.

Based on the above, I’ve come to suspect that putting on a surface coat can (could) help strengthen old concrete, even long after it’s hardened. Mostly this would happen by filling in voids and cracks, but also by extending the polymer chains. I imagine it would be especially helpful to apply the surface coat somewhat watery on a dry day in the summer. In that case, I imagine that Ca3SiO5 and Ca(OH)2 from the surface coat will penetrate and fill the pores of the concrete below — the sales pores that hiss when you pour water on them. I imagine this would fill cracks and voids, and extend existing CaOSiO2•H2O chains. The coat should add strength, and should be attractive as well. At least that was my thought.

I should note that, while Portland cement is mostly Ca3SiO5, there is also a fair amount (25%) of Ca2SiO4. This component reacts with water to form the same calcium-silicate polymer as above, but does so at a slower rate using less water per gram. My hope was that this component would be the main one to diffuse into deep pores of the concrete, reacting there to strengthen the concrete long after surface drying had occurred.

Trump tower: 664', concrete and glass. What grade of concrete would you use?

Trump tower: 664′, concrete and glass. What grade of concrete would you use?

As it happened, I had a chance to test my ideas this summer and also about 3 years ago. The city inspector came by to say the concrete flags outside my house were rough, and thus needed replacing, and that I was to pay or do it myself. Not that I understand the need for smooth concrete, quite, but that’s our fair city. I applied for a building permit to apply a surface coat, and applied it watery. I used “Quickrete” brand concrete patch, and so far it’s sticking OK. Pock-holes in the old concrete have been filled in, and so far surface is smooth. We’ll have to see if my patch lasts 10-20 years like fresh cement. Otherwise, no matter how strong the concrete becomes underneath, the city will be upset, and I’ll have to fix it. I’ve noticed that there is already some crumbling at the sides of flags, something I attribute to the extra water. It’s not a problem yet, but hope this is not the beginning of something worse. If I’m wrong here, and the whole seal-coat flakes off, I’ll be stuck replacing the flags, or continuing to re-coat just to preserve my reputation. But that’s the cost of experimentation. I tried something new, and am blogging about it in the hope that you and I benefit. “Education is what you get when you don’t get what you want.” (It’s one of my wise sayings). At the worst, I’ll have spent 90 lb of patching cement to get an education. And, I’m happy to say that some of the relatively new concrete flags that the city put in are already cracked. I attribute this to: too much sand, air, water or air (they don’t look like they have much rock): Poor oversight.

Dr. Robert E. Buxbaum. March 5, 2019. As an aside, the 664 foot Trump Tower, NY is virtually the only skyscraper in the city to be built of concrete and glass. The others are mostly steel and glass. Concrete and glass is supposed to be stiffer and quieter. The engineer overseeing the project was Barbara Res, the first woman to oversee a major, NY building project. Thought question: if you built the Trump Tower, which quality of concrete would you use, and why.

Isotopic effects in hydrogen diffusion in metals

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For most people, there is a fundamental difference between solids and fluids. Solids have long-term permanence with no apparent diffusion; liquids diffuse and lack permanence. Put a penny on top of a dime, and 20 years later the two coins are as distinct as ever. Put a layer of colored water on top of plain water, and within a few minutes you’ll see that the coloring diffuse into the plain water, or (if you think the other way) you’ll see the plain water diffuse into the colored.

Now consider the transport of hydrogen in metals, the technology behind REB Research’s metallic  membranes and getters. The metals are clearly solid, keeping their shapes and properties for centuries. Still, hydrogen flows into and through the metals at a rate of a light breeze, about 40 cm/minute. Another way of saying this is we transfer 30 to 50 cc/min of hydrogen through each cm2 of membrane at 200 psi and 400°C; divide the volume by the area, and you’ll see that the hydrogen really moves through the metal at a nice clip. It’s like a normal filter, but it’s 100% selective to hydrogen. No other gas goes through.

To explain why hydrogen passes through the solid metal membrane this way, we have to start talking about quantum behavior. It was the quantum behavior of hydrogen that first interested me in hydrogen, some 42 years ago. I used it to explain why water was wet. Below, you will find something a bit more mathematical, a quantum explanation of hydrogen motion in metals. At REB we recently put these ideas towards building a membrane system for concentration of heavy hydrogen isotopes. If you like what follows, you might want to look up my thesis. This is from my 3rd appendix.

Although no-one quite understands why nature should work this way, it seems that nature works by quantum mechanics (and entropy). The basic idea of quantum mechanics you will know that confined atoms can only occupy specific, quantized energy levels as shown below. The energy difference between the lowest energy state and the next level is typically high. Thus, most of the hydrogen atoms in an atom will occupy only the lower state, the so-called zero-point-energy state.

A hydrogen atom, shown occupying an interstitial position between metal atoms (above), is also occupying quantum states (below). The lowest state, ZPE is above the bottom of the well. Higher energy states are degenerate: they appear in pairs. The rate of diffusive motion is related to ∆E* and this degeneracy.

A hydrogen atom, shown occupying an interstitial position between metal atoms (above), is also occupying quantum states (below). The lowest state, ZPE is above the bottom of the well. Higher energy states are degenerate: they appear in pairs. The rate of diffusive motion is related to ∆E* and this degeneracy.

The fraction occupying a higher energy state is calculated as c*/c = exp (-∆E*/RT). where ∆E* is the molar energy difference between the higher energy state and the ground state, R is the gas constant and T is temperature. When thinking about diffusion it is worthwhile to note that this energy is likely temperature dependent. Thus ∆E* = ∆G* = ∆H* – T∆S* where asterisk indicates the key energy level where diffusion takes place — the activated state. If ∆E* is mostly elastic strain energy, we can assume that ∆S* is related to the temperature dependence of the elastic strain.

Thus,

∆S* = -∆E*/Y dY/dT

where Y is the Young’s modulus of elasticity of the metal. For hydrogen diffusion in metals, I find that ∆S* is typically small, while it is often typically significant for the diffusion of other atoms: carbon, nitrogen, oxygen, sulfur…

The rate of diffusion is now calculated assuming a three-dimensional drunkards walk where the step lengths are constant = a. Rayleigh showed that, for a simple cubic lattice, this becomes:

D = a2/6τ

a is the distance between interstitial sites and t is the average time for crossing. For hydrogen in a BCC metal like niobium or iron, D=

a2/9τ; for a FCC metal, like palladium or copper, it’s

a2/3τ. A nice way to think about τ, is to note that it is only at high-energy can a hydrogen atom cross from one interstitial site to another, and as we noted most hydrogen atoms will be at lower energies. Thus,

τ = ω c*/c = ω exp (-∆E*/RT)

where ω is the approach frequency, or the amount of time it takes to go from the left interstitial position to the right one. When I was doing my PhD (and still likely today) the standard approach of physics writers was to use a classical formulation for this time-scale based on the average speed of the interstitial. Thus, ω = 1/2a√(kT/m), and

τ = 1/2a√(kT/m) exp (-∆E*/RT).

In the above, m is the mass of the hydrogen atom, 1.66 x 10-24 g for protium, and twice that for deuterium, etc., a is the distance between interstitial sites, measured in cm, T is temperature, Kelvin, and k is the Boltzmann constant, 1.38 x 10-16 erg/°K. This formulation correctly predicts that heavier isotopes will diffuse slower than light isotopes, but it predicts incorrectly that, at all temperatures, the diffusivity of deuterium is 1/√2 that for protium, and that the diffusivity of tritium is 1/√3 that of protium. It also suggests that the activation energy of diffusion will not depend on isotope mass. I noticed that neither of these predictions is borne out by experiment, and came to wonder if it would not be more correct to assume ω represent the motion of the lattice, breathing, and not the motion of a highly activated hydrogen atom breaking through an immobile lattice. This thought is borne out by experimental diffusion data where you describe hydrogen diffusion as D = D° exp (-∆E*/RT).

Screen Shot 2018-06-21 at 12.08.20 AM

You’ll notice from the above that D° hardly changes with isotope mass, in complete contradiction to the above classical model. Also note that ∆E* is very isotope dependent. This too is in contradiction to the classical formulation above. Further, to the extent that D° does change with isotope mass, D° gets larger for heavier mass hydrogen isotopes. I assume that small difference is the entropy effect of ∆E* mentioned above. There is no simple square-root of mass behavior in contrast to most of the books we had in grad school.

As for why ∆E* varies with isotope mass, I found that I could get a decent explanation of my observations if I assumed that the isotope dependence arose from the zero point energy. Heavier isotopes of hydrogen will have lower zero-point energies, and thus ∆E* will be higher for heavier isotopes of hydrogen. This seems like a far better approach than the semi-classical one, where ∆E* is isotope independent.

I will now go a bit further than I did in my PhD thesis. I’ll make the general assumption that the energy well is sinusoidal, or rather that it consists of two parabolas one opposite the other. The ZPE is easily calculated for parabolic energy surfaces (harmonic oscillators). I find that ZPE = h/aπ √(∆E/m) where m is the mass of the particular hydrogen atom, h is Plank’s constant, 6.63 x 10-27 erg-sec,  and ∆E is ∆E* + ZPE, the zero point energy. For my PhD thesis, I didn’t think to calculate ZPE and thus the isotope effect on the activation energy. I now see how I could have done it relatively easily e.g. by trial and error, and a quick estimate shows it would have worked nicely. Instead, for my PhD, Appendix 3, I only looked at D°, and found that the values of D° were consistent with the idea that ω is about 0.55 times the Debye frequency, ω ≈ .55 ωD. The slight tendency for D° to be larger for heavier isotopes was explained by the temperature dependence of the metal’s elasticity.

Two more comments based on the diagram I presented above. First, notice that there is middle split level of energies. This was an explanation I’d put forward for quantum tunneling atomic migration that some people had seen at energies below the activation energy. I don’t know if this observation was a reality or an optical illusion, but present I the energy picture so that you’ll have the beginnings of a description. The other thing I’d like to address is the question you may have had — why is there no zero-energy effect at the activated energy state. Such a zero energy difference would cancel the one at the ground state and leave you with no isotope effect on activation energy. The simple answer is that all the data showing the isotope effect on activation energy, table A3-2, was for BCC metals. BCC metals have an activation energy barrier, but it is not caused by physical squeezing between atoms, as for a FCC metal, but by a lack of electrons. In a BCC metal there is no physical squeezing, at the activated state so you’d expect to have no ZPE there. This is not be the case for FCC metals, like palladium, copper, or most stainless steels. For these metals there is a much smaller, on non-existent isotope effect on ∆E*.

Robert Buxbaum, June 21, 2018. I should probably try to answer the original question about solids and fluids, too: why solids appear solid, and fluids not. My answer has to do with quantum mechanics: Energies are quantized, and always have a ∆E* for motion. Solid materials are those where ω exp (-∆E*/RT) has unit of centuries. Thus, our ability to understand the world is based on the least understandable bit of physics.

Alkaline batteries have second lives

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Most people assume that alkaline batteries are one-time only, throwaway items. Some have used rechargeable cells, but these are Ni-metal hydride, or Ni-Cads, expensive variants that have lower power densities than normal alkaline batteries, and almost impossible to find in stores. It would be nice to be able to recharge ordinary alkaline batteries, e.g. when a smoke alarm goes off in the middle of the night and you find you’re out, but people assume this is impossible. People assume incorrectly.

Modern alkaline batteries are highly efficient: more efficient than even a few years ago, and that always suggests reversibility. Unlike the acid batteries you learned about in highschool chemistry class (basic chemistry due to Volta) the chemistry of modern alkaline batteries is based on Edison’s alkaline car batteries. They have been tweaked to an extent that even the non-rechargeable versions can be recharged. I’ve found I can reliably recharge an ordinary alkaline cell, 9V, at least once using the crude means of a standard 12 V car battery charger by watching the amperage closely. It only took 10 minutes. I suspect I can get nine lives out of these batteries, but have not tried.

To do this experiment, I took a 9 V alkaline that had recently died, and finding I had no replacement, I attached it to a 6 Amp, 12 V, car battery charger that I had on hand. I would have preferred to use a 2 A charger and ideally a charger designed to output 9-10 V, but a 12 V charger is what I had available, and it worked. I only let it charge for 10 minutes because, at that amperage, I calculated that I’d recharged to the full 1 Amp-hr capacity. Since the new alkaline batteries only claimed 1 amp hr, I figured that more charge would likely do bad things, even perhaps cause the thing to blow up.  After 5 minutes, I found that the voltage had returned to normal and the battery worked fine with no bad effects, but went for the full 10 minutes. Perhaps stopping at 5 would have been safer.

I changed for 10 minutes (1/6 hour) because the battery claimed a capacity of 1 Amp-hour when new. My thought was 1 amp-hour = 1 Amp for 1 hour, = 6 Amps for 1/6 hour = ten minutes. That’s engineering math for you, the reason engineers earn so much. I figured that watching the recharge for ten minutes was less work and quicker than running to the store (20 minutes). I used this battery in my firm alarm, and have tested it twice since then to see that it works. After a few days in my fire alarm, I took it out and checked that the voltage was still 9 V, just like when the battery was new. Confirming experiments like this are a good idea. Another confirmation occurred when I overcooked some eggs and the alarm went off from the smoke.

If you want to experiment, you can try a 9V as I did, or try putting a 1.5 volt AA or AAA battery in a charger designed for rechargeables. Another thought is to see what happens when you overcharge. Keep safe: do this in a wood box outside at a distance, but I’d like to know how close I got to having an exploding energizer. Also, it would be worthwhile to try several charge/ discharge cycles to see how the energy content degrades. I expect you can get ~9 recharges with a “non-rechargeable” alkaline battery because the label says: “9 lives,” but even getting a second life from each battery is a significant savings. Try using a charger that’s made for rechargeables. One last experiment: If you’ve got a cell phone charger that works on a car battery, and you get the polarity right, you’ll find you can use a 9V alkaline to recharge your iPhone or Android. How do I know? I judged a science fair not long ago, and a 4th grader did this for her science fair project.

Robert Buxbaum, April 19, 2018. For more, semi-dangerous electrochemistry and biology experiments.

Keeping your car batteries alive.

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Lithium-battery cost and performance has improved so much that no one uses Ni-Cad or metal hydride batteries any more. These are the choice for tools, phones, and computers, while lead acid batteries are used for car starting and emergency lights. I thought I’d write about the care and trade-offs of these two remaining options.

As things currently stand, you can buy a 12 V, lead-acid car battery with 40 Amp-h capacity for about $95. This suggests a cost of about $200/ kWh. The price rises to $400/kWh if you only discharge half way (good practice). This is cheaper than the per-power cost of lithium batteries, about $500/ kWh or $1000/ kWh if you only discharge half-way (good practice), but people pick lithium because (1) it’s lighter, and (2) it’s generally longer lasting. Lithium generally lasts about 2000 half-discharge cycles vs 500 for lead-acid.

On the basis of cost per cycle, lead acid batteries would have been replaced completely except that they are more tolerant of cold and heat, and they easily output the 400-800 Amps needed to start a car. Lithium batteries have problems at these currents, especially when it’s hot or cold. Lithium batteries deteriorate fast in the heat too (over 40°C, 105°F), and you can not charge a lithium car battery at more than 3-4 Amps at temperatures below about 0°C, 32°F. At higher currents, a coat of lithium metal forms on the anode. This lithium can react with water: 2Li + H2O –> Li2O + H2, or it can form dendrites that puncture the cell separators leading to fire and explosion. If you charge a lead acid battery too fast some hydrogen can form, but that’s much less of a problem. If you are worried about hydrogen, we sell hydrogen getters and catalysts that remove it. Here’s a description of the mechanisms.

The best thing you can do to keep a lead-acid battery alive is to keep it near-fully charged. This can be done by taking long drives, by idling the car (warming it up), or by use of an external trickle charger. I recommend a trickle charger in the winter because it’s non-polluting. A lead-acid battery that’s kept at near full charge will give you enough charge for 3000 to 5000 starts. If you let the battery completely discharge, you get only 50 or so deep cycles or 1000 starts. But beware: full discharge can creep up on you. A new car battery will hold 40 Ampere-hours of current, or 65,000 Ampere-seconds if you half discharge. Starting the car will take 5 seconds of 600 Amps, using 3000 Amp-s or about 5% of the battery’s juice. The battery will recharge as you drive, but not that fast. You’ll have to drive for at least 500 seconds (8 minutes) to recharge from the energy used in starting. But in the winter it is common that your drive will be shorter, and that a lot of your alternator power will be sent to the defrosters, lights, and seat heaters. As a result, your lead-acid battery will not totally charge, even on a 10 minute drive. With every week of short trips, the battery will drain a little, and sooner or later, you’ll find your battery is dead. Beware and recharge, ideally before 50% discharge

A little chemistry will help explain why full discharging is bad for battery life (for a different version see Wikipedia). For the first half discharge of a lead-acid battery, the reaction Is:

Pb + 2PbO2 + 2H2SO4  –> PbSO4 + Pb2O2SO4 + 2H2O.

This reaction involves 2 electrons and has a -∆G° of >394 kJ, suggesting a reversible voltage more than 2.04 V per cell with voltage decreasing as H2SO4 is used up. Any discharge forms PbSO4 on the positive plate (the lead anode) and converts lead oxide on the cathode (the negative plate) to Pb2O2SO4. Discharging to more than 50% involves this reaction converting the Pb2O2SO4 on the cathode to PbSO4.

Pb + Pb2O2SO4 + 2H2SO4  –> 2PbSO4 + 2H2O.

This also involves two electrons, but -∆G < 394 kJ, and voltage is less than 2.04 V. Not only is the voltage less, the maximum current is less. As it happens Pb2O2SO4 is amorphous, adherent, and conductive, while PbSO4 is crystalline, not that adherent, and not-so conductive. Operating at more than 50% results in less voltage, increased internal resistance, decreased H2SO4 concentrations, and lead sulfate flaking off the electrode. Even letting a battery sit at low voltage contributes to PbSO4 flaking off. If the weather is cold enough, the low concentration H2SO4 freezes and the battery case cracks. My advice: Get out your battery charger and top up your battery. Don’t worry about overcharging; your battery charger will sense when the charge is complete. A lead-acid battery operated at near full charge, between 67 and 100% will provide 1500 cycles, about as many as lithium. 

Trickle charging my wife's car. Good for battery life. At 6 Amps, expect this to take 3-6 hours.

Trickle charging my wife’s car: good for battery life. At 6 Amps, expect a full charge to take 6 hours or more. You might want to recharge the battery in your emergency lights too. 

Lithium batteries are the choice for tools and electric vehicles, but the chemistry is different. For longest life with lithium batteries, they should not be charged fully. If you change fully they deteriorate and self-discharge, especially when warm (100°F, 40°C). If you operate at 20°C between 75% and 25% charge, a lithium-ion battery will last 2000 cycles; at 100% to 0%, expect only 200 cycles or so.

Tesla cars use lithium batteries of a special type, lithium cobalt. Such batteries have been known to explode, but and Tesla adds sophisticated electronics and cooling systems to prevent this. The Chevy Volt and Bolt use lithium batteries too, but they are less energy-dense. In either case, assuming $1000/kWh and a 2000 cycle life, the battery cost of an EV is about 50¢/kWh-cycle. Add to this the cost of electricity, 15¢/kWh including the over-potential needed to charge, and I find a total cost of operation of 65¢/kWh. EVs get about 3 miles per kWh, suggesting an energy cost of about 22¢/mile. By comparison, a 23 mpg car that uses gasoline at $2.80 / gal, the energy cost is 12¢/mile, about half that of the EVs. For now, I stick to gasoline for normal driving, and for long trips, suggest buses, trains, and flying.

Robert Buxbaum, January 4, 2018.

magnetic separation of air

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As some of you will know, oxygen is paramagnetic, attracted slightly by a magnet. Oxygen’s paramagnetism is due to the two unpaired electrons in every O2 molecule. Oxygen has a triple-bond structure as discussed here (much of the chemistry you were taught is wrong). Virtually every other common gas is diamagnetic, repelled by a magnet. These include nitrogen, water, CO2, and argon — all diamagnetic. As a result, you can do a reasonable job of extracting oxygen from air by the use of a magnet. This is awfully cool, and could make for a good science fair project, if anyone is of a mind.

But first some math, or physics, if you like. To a good approximation the magnetization of a material, M = CH/T where M is magnetization, H is magnetic field strength, C is the Curie constant for the material, and T is absolute temperature.

Ignoring for now, the difference between entropy and internal energy, but thinking only in terms of work derived by lowering a magnet towards a volume of gas, we can say that the work extracted, and thus the decrease in energy of the magnetic gas is ∫∫HdM  = MH/2. At constant temperature and pressure, we can say ∆G = -CH2/2T.

The maximum magnetization you’re likely to get with any permanent magnet (not achieved to date) is about 50 Tesla, or 40,000 ampere meters. At 20°C, the per-mol, magnetic susceptibility of oxygen is 1.34×10−6  This suggests that the Curie constant is 1.34 ×10−6 x 293 = 3.93 ×10−4. Applying this value to oxygen in a 50 Tesla magnet at 20°C, we find the energy difference, ∆G is 1072 J/mole = RT ln ß where ß is a concentration ratio factor between the O2 content of the magnetized and un-magnetized gas, C1/C2 =ß

At room temperature, 298K ß = 1.6, and thus we find that the maximum oxygen concentration you’re likely to get is about 1.6 x 21% = 33%. It’s slightly more than this due to nitrogen’s diamagnetism, but this effect is too small the matter. What does matter is that 33% O2 is a good amount for a variety of medical uses.

I show below my simple design for a magnetic O2 concentrator. The dotted line is a permeable membrane of no selectivity – with a little O2 permeability the design will work better. All you need is a blower or pump. A coffee filter could serve as a membrane.bux magneitc air separator

This design is as simple as the standard membrane-based O2 concentrator – those based on semi-permeable membranes, but this design should require less pressure differential — just enough to overcome the magnet. Less pressure means the blower should be smaller, and less noisy, with less energy use.  I figure this could be really convenient for people who need portable oxygen. With current magnets it would take 4-5 stages or low temperatures to reach this concentration, still this design could have commercial use, I’d think.

On the theoretical end, an interesting thing I find concerns the effect on the entropy of the magnetic oxygen. (Please ignore this paragraph if you have not learned statistical thermodynamics.) While you might imagine that magnetization decreases entropy, other-things being equal because the molecules are somewhat aligned with the field, temperature and pressure being fixed, I’ve come to realize that entropy is likely higher. A sea of semi-aligned molecules will have a slightly higher heat capacity than nonaligned molecules because the vibrational Cp is higher, other things being equal. Thus, unless I’m wrong, the temperature of the gas will be slightly lower in the magnetic area than in the non-magnetic field area. Temperature and pressure are not the same within the separator as out, by the way; the blower is something of a compressor, though a much less-energy intense one than used for most air separators. Because of the blower, both the magnetic and the non magnetic air will be slightly warmer than in the surround (blower Work = ∆T/Cp). This heat will be mostly lost when the gas leaves the system, that is when it flows to lower pressure, both gas streams will be, essentially at room temperature. Again, this is not the case with the classic membrane-based oxygen concentrators — there the nitrogen-rich stream is notably warm.

Robert E. Buxbaum, October 11, 2017. I find thermodynamics wonderful, both as science and as an analog for society.

A clever, sorption-based, hydrogen compressor

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Hydrogen-powered fuel cells provide weight and cost advantages over batteries, important e.g. for drones and extended range vehicles, but they require highly compressed hydrogen and it’s often a challenge compressing the hydrogen. A large-scale solution I like is pneumatic compression, e.g. this compressor. One would combine it with a membrane reactor hydrogen generator, to fill tanks for fuel cells. The problem is that this pump is somewhat complex, and would likely add air impurities to the hydrogen. I’d now like to describe a different, very clever hydrogen pump, one that suited to smaller outputs, but adds no impurities and and provides very high pressure. It operates by metallic hydride sorption at low temperature, followed by desorption at high temperature.

Hydride sorption -desorption pressures vs temperature.

Hydride sorption -desorption pressures vs temperature, from Dhinesh et al.

The metal hydriding reaction is M + nH2 <–> MH2n. Where M is a metal or metallic alloy and MH2n is the hydride. While most metals will undergo this reaction at some appropriate temperature and pressure, the materials of practical interest are exothermic hydrides that is hydrides that give off heat on hydriding. They also must undergo a nearly stoichiometric absorption or desorption reaction at reasonable temperatures and pressures. The plot at right presents the plateau pressure for hydrogen absorption/ desorption in several, exothermic metal hydrides. The most attractive of these are shown in the red box near the center. These sorb or desorb between 1 and 10 atmospheres and 25 and 100 °C.

In this plot, the slope of the sorption line is proportional to the heat of sorption. The most attractive materials for this pump are the ones in the box (or near) with a high slope to the line implying a high heat of sorption. A high heat of sorption means you can get very high compression without too much of a temperature swing.

To me, NaAlH4 appears to be the best of the materials. Though I have not built a pump yet with this material, I’d like to. It certainly serves as a good example for how the pump might work. The basic reaction is:

NaAl + 2H2 <–> NaAlH4

suggesting that each mol of NaAl material (50g) will absorb 2 mols of hydrogen (44.8 std liters). The sorption line for this reaction crosses the 1 atm horizontal line at about 30°C. This suggests that sorption will occur at 1 am and normal room temperature: 20-25°C. Assume the pump contains 100 g of NaAl (2.0 mols). Under ideal conditions, these 100g will 4 mols of hydrogen gas, about 90 liters. If this material in now heated to 226°C, it will desorb the hydrogen (more like 80%, 72 liters) at a pressure in excess of 100 atm, or 1500 psi. The pressure line extends beyond the graph, but the sense is that one could pressure in the neighborhood of 5000 psi or more: enough to use filling the high pressure tank of a hydrogen-based, fuel cell car.

The problem with this pump for larger volume H2 users is time. It will take 2-3 hours to cycle the sober, that is, to absorb hydrogen at low pressure, to heat the material to 226°C, to desorb the H2 and cycle back to low temperature. At a pump rate of 72 liters in 2-3 hours, this will not be an effective pump for a fuel-cell car. The output, 72 liters is only enough to generate 0.12kWh, perhaps enough for the tank of a fuel cell drone, or for augmenting the mpg of gasoline automobiles. If one is interested in these materials, my company, REB Research will be happy to manufacture some in research quantities (the prices blow are for materials cost, only I will charge significantly more for the manufactured product, and more yet if you want a heater/cooler system).

Properties of Metal Hydride materials; Dhanesh Chandra,* Wen-Ming Chien and Anjali Talekar, Material Matters, Volume 6 Article 2

Properties of Metal Hydride materials; Dhanesh Chandra,* Wen-Ming Chien and Anjali Talekar, Material Matters, Volume 6 Article 2

One could increase the output of a pump by using more sorbent, perhaps 10 kg distributed over 100 cells. With this much sorbent, you’ll pump 100 times faster, enough to take the output of a fairly large hydrogen generator, like this one from REB. I’m not sure you get economies of scale, though. With a mechanical pump, or the pneumatical pump,  you get an economy of scale: typically it costs 3 times as much for each 10 times increase in output. For the hydride pump, a ten times increase might cost 7-8 times as much. For this reason, the sorption pump lends itself to low volume applications. At high volume, you’re going to want a mechanical pump, perhaps with a getter to remove small amounts of air impurities.

Materials with sorption lines near the middle of the graph above are suited for long-term hydrogen storage. Uranium hydride is popular in the nuclear industry, though I have also provided Pd-coated niobium for this purpose. Materials whose graph appear at the far, lower left, titanium TiH2, can be used for permanent hydrogen removal (gettering). I have sold Pd-niobium screws for this application, and will be happy to provide other shapes and other materials, e.g. for reversible vacuum pumping from a fusion reactor.

Robert Buxbaum, May 26, 2017 (updated Apr. 4, 2022). 

The chemistry of sewage treatment

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The first thing to know about sewage is that it’s mostly water and only about 250 ppm solids. That is, if you boiled down a pot of sewage, only about 1/40 of 1% of it would remain as solids at the bottom of the pot. There would be some dried poop, some bits of lint and soap, the remains of potato peelings… Mostly, the sewage is water, and mostly it would have boiled away. The second thing to know, is that the solids, the bio-solids, are a lot like soil but better: more valuable, brown gold if used right. While our county mostly burns and landfills the solids remnant of our treated sewage, the wiser choice would be to convert it to fertilizer. Here is a comparison between the composition of soil and bio-solids.

The composition of soil and the composition of bio-solid waste. biosolids are like soil, just better.

The composition of soil and the composition of bio-solid waste. biosolids are like soil, just better.

Most of Oakland’s sewage goes to Detroit where they mostly dry and burn it, and land fill the rest. These processes are expensive and engineering- problematic. It takes a lot of energy to dry these solids to the point where they burn (they’re like really wet wood), and even then they don’t burn nicely. As shown above, the biosolids contain lots of sulfur and that makes combustion smelly. They also contain nitrate, and that makes combustion dangerous. It’s sort of like burning natural gun powder.

The preferred solution is partial combustion (oxidation) at room temperature by bacteria followed by conversion to fertilizer. In Detroit we do this first stage of treatment, the slow partial combustion by bacteria. Consider glucose, a typical carbohydrate,

-HCOH- + O2 –> CO2 + H2O.    ∆G°= -114.6 kcal/mol.

The value of ∆G°, is relevant as a determinate of whether the reaction will proceed. A negative value of ∆G°, as above, indicates that the reaction can progress substantially to completion at standard conditions of 25°C and 1 atm pressure. In a sewage plant, many different carbohydrates are treated by many different bacteria (amoebae, paramnesia, and lactobacilli), and the temperature is slightly cooler than room, about 10-15°C, but this value of ∆G° suggests that near total biological oxidation is possible.

The Detroit plant, like most others, do this biological oxidation treatment using either large stirred tanks, of million gallon volume or so, or in flow reactors with a large fraction of cellular-material returning as recycle. Recycle is needed also in the stirred tank process because of the low solid content. The reaction is approximately first order in oxygen, carbohydrate, and bacteria. Thus a 50% cell recycle more or less doubles the speed of the reaction. Air is typically bubbled through the reactor to provide the oxygen, but in Detroit, pure oxygen is used. About half the organic carbon is oxidized and the remainder is sent to a settling pond. The decant (top) water is sent for “polishing” and dumped in the river, while the goop (the bottom) is currently dried for burning or carted off for landfill. The Holly, MI sewage plant uses a heterogeneous reactors for the oxidation: a trickle bed followed by a rotating disk contractor. These have higher bio-content and thus lower area demands and separation costs, but there is a somewhat higher capital cost.

A major component of bio-solids is nitrogen. Much of this in enters the form of urea, NH2-CO-NH2. In an oxidizing environment, bacteria turns the urea and other nitrogen compounds into nitrate. Consider the reaction the presence of washing soda, Na2CO3. The urea is turned into nitrate, a product suitable for gun powder manufacture. The value of ∆G° is negative, and the reaction is highly favorable.

NH2-CO-NH2 + Na2CO3 + 4 O2 –> 2 Na(NO3) + 2 CO2 + 2 H2O.     ∆G° = -177.5 kcal/mol

The mixture of nitrates and dry bio-solids is highly flammable, and there was recently a fire in the Detroit biosolids dryer. If we wished to make fertilizer, we’d probably want to replace the drier with a further stage of bio-treatment. In Wisconsin, and on a smaller scale in Oakland MI, biosolids are treated by higher temperature (thermophilic) bacteria in the absence of air, that is anaerobically. Anaerobic digestion produces hydrogen and methane, and produces highly useful forms of organic carbon.

2 (-HCOH-) –> CO2 + CH4        ∆G° = -33.7 Kcal/mol

3 (-HCOH-) + H2O –> -CH2COOH + CO2 +  2 1/2 H2        ∆G° = -21.9 kcal/mol

In a well-designed plant, the methane is recovered to provide heat to the plant, and sometimes to generate power. In Wisconsin, enough methane is produced to cook the fertilizer to sterilization. The product is called “Milorganite” as much of it comes from Milwaukee and much of the nitrate is bound to organics.

Egg-shaped, anaerobic biosolid digestors.

Egg-shaped, anaerobic biosolid digestors, Singapore.

The hydrogen could be recovered too, but typically reacts further within the anaerobic digester. Typically it will reduce the iron oxide in the biosolids from the brown, ferric form, Fe2O3, to black FeO.  In a reducing atmosphere,

Fe2O3 + H2 –> 2 FeO + H2O.

Fe2O3 is the reason leaves turn brown in the fall and is the reason that most poop is brown. FeO is the reason that composted soil is typically black. You’ll notice that swamps are filled with black goo, that’s because of a lack of oxygen at the bottom. Sulphate and phosphorous can be bound to ferrous iron and this is good for fertilizer. Generally you want the reduction reactions to go no further.

Weir dam on the river dour. Used to manage floods, increase residence time, and oxygenate the flow.

Weir dam on the river Dour in Scotland. Dams of this type increase residence time, and oxygenate the flow. They’re good for fish, pollution, and flooding.

When allowed to continue, the hydrogen produced by anaerobic digestion begins to reduce sulfate to H2S.

NaSO4 + 4.5 H2 –>  NaOH + 3H2O + H2S.

I’m running for Oakland county, MI water commissioner, and one of my aims is to stop wasting our biosolids. Oakland produces nearly 1000,000 pounds of dry biosolids per day. This is either a blessing or a curse depending on how we use it.

Another issue, Oakland county dumps unpasteurized, smelly black goo into Lake St. Clair every other week, whenever it rains more than one inch. I’d like to stop this by separating the storm and “sanitary” sewage. There is a capital cost, but it can save money because we’d no longer have to pay to treat our rainwater at the Detroit sewage plant. To clean the storm runoff, I’d use mini wetlands and weir dams to increase residence time and provide oxygen. Done right, it would look beautiful and would avoid the flash floods. It should also bring natural fish back to the Clinton River.

Robert Buxbaum, May 24 – Sept. 15, 2016 Thermodynamics plays a big role in my posts. You can show that, when the global ∆G is negative, there is an increase in the entropy of the universe.

How to help Flint and avoid lead here.

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As most folks know, Flint has a lead-poisoning problem that seems to have begun in April, 2014 when the city switched its water supply from Detroit-supplied, Lake Huron water to their own source, water from the Flint River. Here are some thoughts on how to help the affected population, and how to avoid a repeat in Oakland county, where I’m running for water commissioner. First observation, it is not enough to make sure that the source water does not contain lead. The people who decided on the switch had found that the Flint river water had no significant content of lead or other obvious toxins. A key problem, it seems: the river water did not contain anticorrosion phosphates, and none, it seems, were added by the Flint water folks. It also seems that insufficient levels of chlorine (hypochlorite) were added. After the switch, citizens started seeing disgusting, brown water come from their taps, and citizens with lead pipes or solder were poisoned with ppb-levels of lead. There was also an outbreak of legionaries disease that killed 12 people. It was the legionaries that alerted the CDC to the possibility of lead, since it seems the water folks were fudging the numbers there, and hiding that part of the problem.

Flint water, Sept 2015, before switching back to Lake Huron.

Flint water after 5 hours of flushing, Sept 2015, before switching back to Lake Huron.

The city began solving its problem by switching back to Detroit-supplied, Lake Huron water in October, 2015. Beginning in December, 2015, they started adding triple doses of phosphate to the wate. As a result, Flint tap-water is now back within EPA standards, but it’s still fairly unsafe, see here for more details.

There has been a fair amount of finger-pointing. At Detroit for raising the price of water so Flint had to switch, at water officials ignoring the early signs of lead and fudging their reports, at other employees for not adding phosphate or enough chlorine, and at “the system” for not providing Flint’s government with better oversight. My take is that a lot of the problem came from the ignorance of the water commission, and it’s commissioner. We elect our water commissioners to be competent overseers of complex infrastructure, but in may counties folks seem to pick them the same way they pick aldermen: for a nice smile, a great handshake, and an ability to remember names. That, anyway, seems to be the way that Oakland got its current water commissioner. When you pick your commissioner that way, it’s no surprise that he (or she) isn’t particularly up on corrosion chemistry, something that few people understand, and fewer care about until it bites them.

Flint river water contains corrosive chloride that probably helped dissolve the lead from pipes and solder. Contributing to the corrosion problem, I’m going to guess that Flint River water also contains, relatively little carbonate, but significant amounts of chelating chemicals, like EDTA, in 10s of ppb concentration. EDTA isn’t poisonous at these concentrations, but it’s common in industry and is the most commonly used antidote for lead poisoning. EDTA extracts lead and other metals from people and would tend to contribute to the process of extracting lead and iron oxide from the pipes surface into the drinking water. With EDTA in the water, a lot of phosphate or hypochlorite would be needed to avoid the lead poisoning problem and the deadly multiplication of disease.

Detroit ex-mayor Kwame Kilpatrick has claimed that both Flint water and Detroit water were known to be poisoned even a decade before the switch. I find these claims believable given the high levels of lead in kids blood even before the switch. Also, I note that there are areas of Detroit where the blood-lead levels are higher than Flint. Flint tested at the taps in a way that fudged the data during the first days of the poisoning, and I suspect many of our MI cities do this today — just to make the numbers look better. My first suggestion therefore is to test correctly, both at the pipes and at the taps; lead pipes are most-often found in the last few feet before the tap. In particular, we should test at all schools and other places where the state has direct authorization to fix the problem. A MI senate bill has been proposed to this effect, but I’m not sure where it stands in the MI house. It seems there are movements to add lots of ‘riders’ and that’s usually a bad sign.

Another thought is that citizens should be encouraged to test their private taps and helped to fix them. The state can’t come in and test or rip out your private pipes, even if they suspect lead, but the private owner has that authorization. The state could condemn a private property where they believe the water is bad, but I doubt they could evict the residents. It’s a democratic republic, as I understand; you have the right to be deadly stupid. But I’ll take my own suggestion to encourage you: If you think your water has lead, take a sample and call (517) 335-8184. Do it.

Another suggestion, perhaps the easiest and most important, is drink bottled water for now, and if you feel you’ve been poisoned, take an antidote.  As I understand things, the state is already providing bottles of imported water. The most common antidote is, as I’d mentioned, EDTA. Assuming that Flint River water had enough EDTA to significantly worsen the problem, the cheapest antidote might be Flint River water, assuming you drew it in lead-free pipes and chlorinated sufficiently to rid it of bugs. If there is EDTA it will help the poisoned. Another antidote is Succinic acid, something sold by REB Research, my company. As with EDTA it is non-toxic, even in fairly large doses, but its use would have to be doctor- approved.

Robert E. Buxbaum, January 19-31, 2016. I hope this helps. We’d have to check Flint River water for levels of EDTA, but I suspect we’d find biologically significant concentrations. If you think Oakland should have an engineer in charge of the water, elect Buxbaum for water commissioner.

Why are glaciers blue

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i recently returned from a cruse trip to Alaska and, as is typical for such, a highlight of the trip was a visit to Alaska’s glaciers, in our case Hubbard Glacier, Glacier bay, and Mendenhall Glacier. All were blue — bright blue, as were the small icebergs that broke off. Glacier blocks only 2 feet across were bright blue like the glaciers themselves.

Hubbard Glacier, Alaska. Note how blue the ice is

Hubbard Glacier, Alaska. My photo. Note how blue the ice is

What made this interesting/ surprising is that I’ve seen ice sculptures that are 5 foot thick or more, and they are not significantly blue. They have a very slight tinge, but are generally more colorless than glass to my ability to tell. I asked the park rangers why the glaciers were blue, but was given no satisfactory answer. The claim was that glacier ice contained small air bubbles that scattered light the same way that air did. Another park ranger claimed that water is blue by nature, so of course the glaciers were too. The “proof” to this was that the sea was blue. Neither of these seem quite true to me, though there seamed some grains of truth. Sea water, I notice, is sort of blue, but isn’t this shade of blue, certainly not in areas that I’ve lived. Instead, sea water is a rather grayish similar to mud and sea-weeds that I’d expect to find on the sea floor. What’s more, if you look through the relatively clear water of a swimming-pool water to the white-tile bottom, you see only a slight shade of blue-green, even at the 9 foot depth where the light you see has passed through 18 feet of water. This is far more water than an iceberg thickness, and the color is nowhere near as pure blue and the intensity nowhere near as strong.

Plymouth, MI Ice sculpture -- the ice is fairly clear, like swimming pool water

Plymouth, MI Ice sculpture — the ice is fairly clear, like swimming pool water

As for the bubble explanation, it doesn’t seem quite right, either. The bubble size would be non-uniform, with many quite large resulting in a mix of scattered colors — an off white– something seen with the sky of mars. Our earth sky is a purer blue, but this is not because of scattering off of ice-crystals, dust or any other small particles, but rather scattering off the air molecules themselves. The clear blue of glaciers, and of overturned icebergs, suggests (to me) a single-size scattering entity, larger than air molecules, but much smaller than the wavelength of visible light. My preferred entity would be a new compound, a clathrate structure compound, that would be formed from air and ice at high pressures.

An overturned ice-burg is remarkably blue: far bluer than an Ice sculpture. I claim clathrates are the reason.

An overturned ice-burg is remarkably blue: far bluer than an Ice sculpture. I claim clathrates are the reason.

Sea-water forms clathrate compounds with natural gas at high pressures found at great depth. My thought is that similar compounds form between ice and one or more components of air (nitrogen, oxygen, or perhaps argon). Though no compounds of this sort have been quite identified, all these gases are reasonably soluble in water so that suggestion isn’t entirely implausible. The clathrates would be spheres, bigger than air molecules and thus should have more scattering power than the original molecules. An uneven distribution would explain the observation that the blue of glaciers is not uniform, but instead has deeper and lighter blue edges and stripes. Perhaps some parts of the glacier were formed at higher pressures one could expect that these would form more clathrate compounds, and thus more blue. One sees the most intense blue in overturned icebergs — the parts that were under the most pressure.

Robert Buxbaum, October 12, 2015. By the way, some of Alaska’s glaciers are growing and others shrinking. The rangers claimed this was the bad effect of global warming: that the shrinking glaciers should be growing and the growing ones shrinking. They also worried that despite Alaska temperatures reaching 40° below reasonably regularly, it was too warm (for whom?). The lowest recorded temperature in Fairbanks was -66°F in 1961.