Category Archives: Chemistry

Alkaline batteries have second lives

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.

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

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.

With a neodymium magnet, you should be able to get 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 x293 = 3.93 ×10−4  At 20°C, this energy difference is 1072 J/mole. = RT ln ß where ß is the concentration ratio between the O2 content of the magnetized and un-magnetized gas.

From the above, we find that, at room temperature, 298K ß = 1.6, and thus 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 several stages and low temperature operation, this design could have commercial use.

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 pump

Hydrogen-power ed fuel cells provide a lot of advantages over batteries, e.g. for drones and extended range vehicles, but part of the challenge is compressing the hydrogen. On solution I’d proposed is a larger version of this steam-powered compressor, another is a membrane reactor hydrogen generator, and a few weeks ago, I wrote about an other clever innovative solutions: an electrochemical hydrogen pump. It was a fuel cell operating backwards, pumping was very efficient and compact, but the pressure was borne by the fuel cell membranes, so the pump is only suitable at low pressure differentials. I’d now like to describe a different, very clever hydrogen pump, one that operates by metallic hydride sorption and provides very high pressure.

Hydride sorption -desorption pressures vs temperature.

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

The basic metal hydride reaction is M + nH2 <–> MH2n. Where M is a metal or metallic alloy. While most metals will undergo this reaction at some appropriate temperature and pressure, the materials of interest are exothermic hydrides that undergo a nearly stoichiometric absorption or desorption reaction at temperatures near 1 atm, temperatures near room temperature. The plot at right presents the plateau pressure for hydrogen absorption/ desorption in several, common metal hydrides. The slope is proportionals to the heat of sorption. There is a red box shown for the candidates that sorb or desorb between 1 and 10 atmospheres and 25 and 100 °C. Sorbants whose lines pass through that box are good candidates for pump use. The ones with a high slope (high heat of sorption) in particular, if you want a convenient source of very high pressure.

To me, NaAlH4 is among the best of the materials, and certainly serves as a good example for how the pump works. The basic reaction, in this case is:

NaAl + 2H2 <–> NaAlH4

The line for this reaction crosses the 1 atm red line at about 30°C suggesting that each mol of NaAl material will absorb 2 mols of hydrogen at 1 am and normal room temperatures: 20-30°C. Assume the pump contains 100 g of NaAl (2.0 mols). We can expect it will 4 mols of hydrogen gas, about 90 liters at this temperature. If this material in now heated to 250°C, it will desorb most of the hydrogen (80% perhaps, 72 liters) at 100 atm, or 1500 psi. This is a remarkably high pressure boost; 1500 psi hydrogen is suitable for use filling the high pressure tank of a hydrogen-based, fuel cell car.

But there is a problem: it will take 2-3 hours to cycle the sober; the absorb hydrogen at low pressure, heat, desorb and cycle back to low temperature. If you only can pump 72 liters in 2-3 hours, this will not be an effective pump for automobiles. Even with several cells operating in parallel, it will be hard to fill the fuel tank of a fuel-cell car. The output is enough for electric generators, or for the small gas 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 supply them in research quantities.

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

At this point, I can imagine you saying that there is a simple way to make up for the low output of a pump with 100g of sorbent: use more, perhaps 10 kg distributed over 100 cells. The alloys don’t cost much in bulk, see chart above (they’re a lot more expensive in small quantities). With 100 times more sorbent, you’ll pump 100 times faster, enough for a fairly large hydrogen generator, like this one from REB. This will work, but you don’t get economies of scale. With standard, mechanical pumps give you a decent economy of scale — it costs 3-4 times as much for each 10 times increase in output. For this reason, the hydride sorption pump, though clever appears to be destined for low volume applications. Though low volume might involve hundreds of kg of sorbent, at some larger value, you’re going to want to use a mechanical pump.

Other uses of these materials include hydrogen storageremoval of hydrogen from a volume, e.g. so it does not mess up electronics, or for vacuum pumping from a futon reactor. I have sold niobium screws for hydrogen sorption in electronic packages, and my company provides chemical sorbers for hydrogen removal from air. For more of our products, visit www.rebresearch.com/catalog.html

Robert Buxbaum, May 26, 2017. 

The chemistry of sewage treatment

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- + O–> CO+ 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-) –> COCH4        ∆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.

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 plant pulling 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 source water does not contain lead. The people who decided on the switch had found that the Flint river water had no significant 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. 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.

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 switched back to Detroit-supplied, Lake Huron water in October, 2015, and they started adding triple doses of phosphates to the water in December. As a result, Flint tap-water is now back within EPA standards, but it’s still likely 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 workers for ignoring the early signs of lead, at other employees for not adding the additive, and at “the system” for not caring, or providing Flint with decent infrastructure. I suspect that a lot of the problem is ignorance in the water commission. We elect our water commissioners, and folks seem to pick them the same way we pick presidents: 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 sensitive to corrosion, an invisible chemical phenomenon that few people understand.

Flint river water contains corrosive chloride. Contributing to the corrosion problem, I’m going to guess that Flint River water also contains an industrial chelating chemical used in plating, EDTA in 10s of ppb concentrations. EDTA isn’t poisonous at these concentrations, but it’s the most commonly used antidote for lead poisoning and commonly used in industry. At these concentrations, EDTA extracts lead and other metals from people and I’m going to guess that this same chemical, or something very similar, contributed to the process that extracted lead and iron oxide from the pipes. With EDTA in the water, no amount of phosphate would avoid or solve the lead poisoning problem.

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 did not test at the taps in a scientifically acceptable way during the first days of the poisoning, and neither, I suspect, do many of our MI cities today. 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 to provide an antidote. The main antidotes for lead are chelating compounds, and we’re already providing bottles of imported water. Why not provide some of the water with compounds that help extract lead from people. And here I have an interesting thought. Assuming I am right that Flint River water had enough EDTA to cause/ worsen the problem, the cheapest/ best antidote might be Flint River water. You’d want to draw the water with plastic pipes and chlorinate it to rid it of bugs, but if there is EDTA it will help the poisoned. EDTA is a known lead-poisoning antidote. Another antidote is Succinct acid, something sold by REB Research, my company. There are other antidotes too, but wouldn’t it be cool if Flint river water worked?

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 it at 50 ppb, or so, a biologically significant concentration. If you think Oakland should have an engineer in charge of the water, elect Buxbaum for water commissioner.

Why are glaciers blue

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.

Much of the chemistry you learned is wrong

When you were in school, you probably learned that understanding chemistry involved understanding the bonds between atoms. That all the things of the world were made of molecules, and that these molecules were fixed proportion combinations of the chemical elements held together by one of the 2 or 3 types of electron-sharing bonds. You were taught that water was H2O, that table salt was NaCl, that glass was SIO2, and rust was Fe2O3, and perhaps that the bonds involved an electron transferring between an electron-giver: H, Na, Si, or Fe… to an electron receiver: O or Cl above.

Sorry to say, none of that is true. These are fictions perpetrated by well-meaning, and sometime ignorant teachers. All of the materials mentioned above are grand polymers. Any of them can have extra or fewer atoms of any species, and as a result the stoichiometry isn’t quite fixed. They are not molecules at all in the sense you knew them. Also, ionic bonds hardly exist. Not in any chemical you’re familiar with. There are no common electron compounds. The world works, almost entirely on covalent, shared bonds. If bonds were ionic you could separate most materials by direct electrolysis of the pure compound, but you can not. You can not, for example, make iron by electrolysis of rust, nor can you make silicon by electrolysis of pure SiO2, or titanium by electrolysis of pure TiO. If you could, you’d make a lot of money and titanium would be very cheap. On the other hand, the fact that stoichiometry is rarely fixed allows you to make many useful devices, e.g. solid oxide fuel cells — things that should not work based on the chemistry you were taught.

Iron -zinc forms compounds, but they don't have fixed stoichiometry. As an example the compound at 60 atom % Zn is, I guess Zn3Fe2, but the composition varies quite a bit from there.

Iron -zinc forms compounds, but they don’t have fixed stoichiometry. As an example the compound at 68-80 atom% Zn is, I guess Zn7Fe3 with many substituted atoms, especially at temperatures near 665°C.

Because most bonds are covalent many compounds form that you would not expect. Most metal pairs form compounds with unusual stoicheometric composition. Here, for example, is the phase diagram for zinc and Iron –the materials behind galvanized sheet metal: iron that does not rust readily. The delta phase has a composition between 85 and 92 atom% Zn (8 and 15 a% iron): Perhaps the main compound is Zn5Fe2, not the sort of compound you’d expect, and it has a very variable compositions.

You may now ask why your teachers didn’t tell you this sort of stuff, but instead told you a pack of lies and half-truths. In part it’s because we don’t quite understand this ourselves. We don’t like to admit that. And besides, the lies serve a useful purpose: it gives us something to test you on. That is, a way to tell if you are a good student. The good students are those who memorize well and spit our lies back without asking too many questions of the wrong sort. We give students who do this good grades. I’m going to guess you were a good student (congratulations, so was I). The dullards got confused by our explanations. They asked too many questions, and asked, “can you explain that again? Or why? We get mad at these dullards and give them low grades. Eventually, the dullards feel bad enough about themselves to allow themselves to be ruled by us. We graduates who are confident in our ignorance rule the world, but inventions come from the dullards who don’t feel bad about their ignorance. They survive despite our best efforts. A few more of these folks survive in the west, and especially in America, than survive elsewhere. If you’re one, be happy you live here. In most countries you’d be beheaded.

Back to chemistry. It’s very difficult to know where to start to un-teach someone. Lets start with EMF and ionic bonds. While it is generally easier to remove an electron from a free metal atom than from a free non-metal atom, e.g. from a sodium atom instead of oxygen, removing an electron is always energetically unfavored, for all atoms. Similarly, while oxygen takes an extra electron easier than iron would, adding an electron is energetically unfavored. The figure below shows the classic ion bond, left, and two electron sharing options (center right) One is a bonding option the other anti-bonding. Nature prefers this to electron sharing to ionic bonds, even with blatantly ionic elements like sodium and chlorine.

Bond options in NaCl. Note that covalent is the stronger bond option though it requires less ionization.

Bond options in NaCl. Note that covalent is the stronger bond option though it requires less ionization.

There is a very small degree of ionic bonding in NaCl (left picture), but in virtually every case, covalent bonds (center) are easier to form and stronger when formed. And then there is the key anti-bonding state (right picture). The anti bond is hardly ever mentioned in high school or college chemistry, but it is critical — it’s this bond that keeps all mater from shrinking into nothingness.

I’ve discussed hydrogen bonds before. I find them fascinating since they make water wet and make life possible. I’d mentioned that they are just like regular bonds except that the quantum hydrogen atom (proton) plays the role that the electron plays. I now have to add that this is not a transfer, but a covalent spot. The H atom (proton) divides up like the electron did in the NaCl above. Thus, two water molecules are attracted by having partial bits of a proton half-way between the two oxygen atoms. The proton does not stay put at the center, there, but bobs between them as a quantum cloud. I should also mention that the hydrogen bond has an anti-bond state just like the electron above. We were never “taught” the hydrogen bond in high school or college — fortunately — that’s how I came to understand them. My professors, at Princeton saw hydrogen atoms as solid. It was their ignorance that allowed me to discover new things and get a PhD. One must be thankful for the folly of others: without it, no talented person could succeed.

And now I get to really weird bonds: entropy bonds. Have you ever noticed that meat gets softer when its aged in the freezer? That’s because most of the chemicals of life are held together by a sort of anti-bond called entropy, or randomness. The molecules in meat are unstable energetically, but actually increase the entropy of the water around them by their formation. When you lower the temperature you case the inherent instability of the bonds to cause them to let go. Unfortunately, this happens only slowly at low temperatures so you’ve got to age meat to tenderize it.

A nice thing about the entropy bond is that it is not particularly specific. A consequence of this is that all protein bonds are more-or-less the same strength. This allows proteins to form in a wide variety of compositions, but also means that deuterium oxide (heavy water) is toxic — it has a different entropic profile than regular water.

Robert Buxbaum, March 19, 2015. Unlearning false facts one lie at a time.

Change your underwear; of mites and men

The underware bomber mites make it right.

Umar, the underwear bomber.

For those who don’t know it, the underwear bomber, Umar Farook Abdulmutallab, wore his pair of explosive underwear for 3 weeks straight before trying to detonate them while flying over Detroit in 2009. They didn’t go off, leaving him scarred for life. It’s quite possible that the nasty little mites that live in underwear stopped the underwear bomber. They are a main source of US allergens too.

Dust mite, skin, and pollen seen with a light  microscope. Gimmie some skin.

Dust mite, skin, and pollen seen with a light microscope. Gimmie some skin.

If you’ve ever used an electron microscope to look at household objects, you’ll find them covered with brick-like flakes of dried out skin-cells: yours and your friends’. Each person sheds his or her skin every month, on average. The outer layer dries out and flakes off as new skin grows in behind it. Skin flakes are the single largest source of household dust, and if not for the fact that these flakes are the main food for mites, your house would be chock full of your left over skin. When sunlight shines in your window, you see the shimmer of skin-flakes hanging in the air. Under the electron microscope, the fresh skin flakes look like bricks, but mite-eaten skin flakes look irregular. Less common, but more busy are the mites.

The facial mite movie. They live on in us, about 1 per hair follicle, particularly favoring eyelashes. Whenever you shower, your shower with a friend.

The facial mite movie. They live on in us, about 1 per hair follicle, particularly favoring eyelashes. Whenever you shower, your shower with a friend.

Dry skin is mostly protein (keratin), plus cholesterol and squalene. This provides great nutrition for dust mites and their associated bacteria. In warm, damp environments, as in your underwear or mattress, these beasties multiply and eat the old skin. The average density of dust mites on a mattress is greater than 2500/gram of dust.[1]  The mites leave behind excrement and broken off mite-limbs: nasty bits that are the most common allergens in the US today.

An allergy to dust shows up as sneezing, coughing, clogged lungs, and eczema. The most effective cure is a high level of in-home hygiene; mites don’t like soap or dry air. You’ve go to mop and vacuum regularly. Clean and change your clothing, particularly your undergarments; rotate your mattresses, and shake the dust out of your bedding. Vacuuming is less-effective as a significant fraction of the nasties go through the filter and get spread around by the vacuum blower.

As it turns out, dust mites and their bacteria eat more than skin. They also eat dried body fluids, poop residue, and the particular explosive used by Umar Farook, pentaerythritol tetra nitrate, PETN (humans can eat and metabolize this stuff too — it’s an angina treatment). The mites turn PETN into less-explosive versions, plus more mites.

Mighty mites as seen with electron microscopy. They eat more than skin.

Mighty mites as seen with electron microscopy. They eat more than skin.

There are many varieties of mite living on and among us. Belly button mites, for example, and face mites as shown above (click on the image to see it move). On average, people have one facial mite per hair follicle. It’s also possible that the bomber was stopped by poor quality control engineering and not mites at all. Religion tends to be at odds with a science like quality control, and followers tend to put their faith in miracles.

Chigger turning on a dime

Chigger turning on a dime

larger than the dust mite is the chigger, shown at left. Chiggers leave visible bites, particularly along the underwear waste-band. There are larger-yet critters in the family: lice, bed bugs, crabs. Bathing regularly, and cleaning your stuff will rid yourself of all these beasties, at least temporarily. Keeping your hair short and your windows open helps too. Mites multiply in humid, warm environments. Opening the windows dries and cools the air, and blows out mite-bits that could cause wheezing. Benjamin Franklin and took air-baths too: walking around naked with the windows open, even in winter. It helped that he lived on the second floor. Other ways to minimize mite growth include sunlight, DOT (a modern version of DDT), and eucalyptus oil. At the very minimum, change your underwear regularly. It goes a long way to reduce dust embarrassing moments at the jihadist convention.

Dr. Robert E. Buxbaum, Sept 21, 2014. Not all science or life is this weird and wonderful, but a lot is, and I prefer to write about the weird and wonderful bits. See e.g. the hazards of health food, the value of sunshine, or the cancer hazard of living near a river. Or the grammar of pirates.

New mixed drink, the R°

Earlier this week, R__ turned 21, the drinking age in most of the USA. As a gift to her, I thought I might invent a new mixed drink that would suit her taste, and make her birthday more special. My requirements: that it should be kosher, that it’s made with widely available ingredients; that it should be relatively sophisticated, that it should be lower in alcohol (a fatherly concern), and that it should taste good to her and the general public.

The R___: gin tonic and grenadine

The R°: gin, tonic , ice, and grenadine

What I came up with, is something I call,The R°. It’s a modification of one of the great drinks of the western world, the gin and tonic. My modification is to use less gin, and to use grenadine instead of the traditional squeeze of lime. As she gets older, she may want to increase the gin content. The recipe: put 2/3 shot gin in a 10 oz straight-sided glass. Fill the glass 2/3 full of ice, near-fill with tonic water, and add a dash of grenadine, 1/4 shot or so (I used Rose’s). Stir slightly so the pink color stays mostly on the bottom. The result is slightly sweeter than the traditional gin and tonic, kosher in almost all places (you’ve got to check, but generally true), fairly sophisticated, good-tasting, and a reminder of Israel, a country where pomegranates grow all over. If you order one at a place with black lights and doesn’t stir much,you’ll discover that the tonic water glows electric-blue.

The verdict: R__ liked it. My hope is that you will enjoy it too. As a literary note, grenade is French for pomegranate; hand grenades got their name because of the shape. This drink is also suitable for talk like a pirate day (September 19).

Sept 14, 2014. My only previous gastronomic post was a recipe to make great lemonade. For a song by my daughter, go here, or here. For a joke about a neutron walking into a bar, go here.