Odorless Fertilizer from Household Urine

103_9206Using an analysis of urine and  some basic arithmetic,  I realized yesterday how many useful compounds we unfortunately flush down the toilet.  Based on the fact that an individual’s daily output of urine contains about 3 g of potassium ion, 2.5 grams(g)  of phosphate,  16 g of urea, 3 g of amino acids, a gram of ammonia,  a city of 3 million like Montreal, wastes 75 000 kg of potential fertilizer every day, none of which is recovered by sewage treatment.

Luckily there is a solution. At the elevated pH of aged urine, which will also have more ammonium, the latter ion and phosphate will not dissolve in the presence of magnesium ion. If we add more Mg²+ we can create far more of a precipitate known as struvite. The product is odorless, free of heavy metal contaminants and releases ammonium and phosphate ions slowly into the soil, whose pH is lower than that of aged urine.  Over a temperature range of 25 to  35 oC, 169 to 213 mg of struvite dissolve per liter at pH = 7. The Swiss Federal Institute of Aquatic Science is actively involved in building STUN reactors to recover struvite, but I had to try it myself on a small scale.

The easiest way to do it at home is to use Epsom salts. Every 500 L of urine can potentially produce 1 kg of struvite.  The typical trip to the washroom produces 0.150 L  of urine, which proportionally can yield  0.3 g of MgNH4PO4.6H­2O (struvite) or 0.3/245 = 0.00245 moles. Epsom salts are pure hydrated magnesium sulfate MgSO4.7H2O (246g/mole), and a one to one molar ratio, implies that 0.603 g of  the magnesium salt are needed.

With time, the urine’s pH climbs because urea, urine’s 2nd most common ingredient, breaks down into ammonia, which dissolves in urine’s primary ingredient to produce ammonium hydroxide(NH4OH). At an elevated pH, MgNH4PO4.6H­2O  is poorly soluble, which makes recovery possible. Hoping to eventually form larger crystals, I added the Epsom salts while  the urine was fresh so that the struvite would precipitate gradually. As expected no crystals were apparent without the formation of ammonium.


Struvite crystals, six days after the addition of Epsom salts to fresh urine. The white crystals are at the lower edge of the mason jar.


Filtered struvite crystals, seven days after the start of the experiment.

It took between 15 and 34 hours of aging  to get the initial crystals, a time that was probably elongated by the fact that the mason jar of urine was kept in a cold storage room. More crystals formed during the week. I checked the pH of the urine 7 days after first adding the epsom salts, and it had gone alkaline.

To make sure I was not getting magnesium sulfate instead of struvite crystallizing (unlikely since it’s soluble in ocean water), I added baking soda to a magnesium sulfate solution and obtained no precipitate. For an additional control I added baking soda, magnesium sulfate to filtered aged urine and again obtained no crystals.

The cost of Epsom salt is about $ 3.00/kg .  One kg of Epsom salt will produce 0.3/0.603 = 0.5 kg of struvite; so it would cost $6 to make a kilogram of  struvite, which exceeds the cost of commercial fertilizer , and the latter also has the advantage of including plant-required K+. Magnesium carbonate, is cheaper, and can be more easily converted into soluble MgO, but the process releases carbon dioxide into the environment. So the most ecological and economical source of Mg2+  would be bittern, a magnesium-rich liquid that remains after salt (NaCl) is extracted from seawater.

You may be unaware, as I was, that struvite is found in a type of kidney stones. Unlike oxalate or uric acid stones, the struvite -variety sometimes form after frequent urinary infections. The bacteria responsible for the infection hydrolyzes urea to ammonium, raising the pH and mirroring the precipitating effect used in the recovery of struvite from aged urine.


 Other Sources:
  http://oatao.univ-toulouse.fr/4645/1/Hanhoun_4645.pdf ( a detailed study of the solubility and thermodynamics of struvite )


The Causes and Consequences Of Grade Inflation

Grade inflation is common. It recognizes no borders, occurring in public and private schools, at the elementary level and in Ivy League universities. These graphics from the NY Times demonstrate how grades have mushroomed while SAT critical reading scores have not changed much between 1980 and 2011! (source: College Board) grades1It is a serious problem, and yet frank and open discussions about the matter are as rare as a school with a C-average. Here’s an insider’s look at both the consequences and causes.

A- The Consequences Of Grade Inflation

1. Within elementary and high schools, grade inflation leads to improper placement of students. Kids who have so far displayed only mediocre ability end up in difficult science and mathematics classes. Eventually they get turned off and some develop long lasting hangups towards the courses, whereas if they had been placed in a more appropriate level, they would have stumbled less and learned more, and they still would have been able to eventually enroll in more rigorous courses.

2. When no one fails a subject, problems are possibly being swept under the rug. Many teachers have always set up evaluation schemes so that a good worker who has the prerequisites should not fail. But so much can happen despite the best intentions: students slip in between the cracks and get into a course without the necessary background; some get overconfident and procrastinate; they get depressed over relationships; get over-involved in activities or are often absent for various reasons, or they hit some kind of intellectual wall and don’t have the persistence or motivation to get around it. But if all is made too easy, the student is never obliged to make adjustments to life’s curve balls, and they are not redirected to a more suitable or remedial avenue.

3. Inflation blurs the landscape for college admissions offices. It can be argued that colleges are guilty of accepting too many students. In our province, for instance, 46% of young adults between the ages of 18 and 20 are in college.  But when so many high schools have 40% of their seniors on honor rolls, the few colleges that do not constantly increase their quotas have no choice but to lift the cutoff point for accepting students. In an atmosphere of pumped-up grades, some marks may be more exaggerated than others, which allows some weaker students to be accepted, and more deserving ones to be unfairly rejected.

4. When the majority of specialty college programs or university graduates get A’s, industry has little basis for hiring someone. Yet society is well-served when a grading system is fair and reserves the best grades only for the most with most aptitude, motivation and persistence. The following examples are anecdotal, but I don’t think they were unusual in an atmosphere when there was little grade inflation. Two guys I knew since high school were at the top of their class in electrical engineering at McGill University in the early 1980’s. They were both recruited by a Canadian company on campus, and they proved to be excellent choices because the department had good standards. Without the latter, industries have to resort to subjective interviews, their own testing and possibly patronage.
In my chemistry program, summer jobs were given to students who had the highest grades in a quantitative analytical laboratory course because to get the right result, one had to be dexterous, meticulous and know how to handle the raw numerical results. No one could perform the experiment for you, and you could not copy off someone else because everyone received a different unknown sample.

5. When students see that numbers can be manipulated and pumped up, it sets the stage for future “creative accounting” in their own personal and/or corporate lives. Schools are a microcosm of society, so if we plant the seeds of bad habits, they will some day create more monstrous examples within our economy or even within a research environment.

B- The Causes of Grade Inflation

1. Millions of parents have the monotonous expectation that their children will become university graduates. Marks have become tickets to that sometimes unrealistic goal. This expectation and the accompanying  huge flow of money pressures the school system into “printing” as many tickets as they can to let students enter the next level. Meanwhile, we experience a shortage of tradespeople because we are too busy granting a disproportionate number of degrees in certain fields.  31% of 2008 Canadian degrees were awarded to the humanities, sociology and other social sciences and many of the life science degrees are awarded to students who are only using them as second attempts to get into medical school. Yet excellent 3-year specialty programs are completely overlooked. One of many examples is that due to low demand by students, there is only one program in our entire province that trains them to work in the plastics industry. Yet the industrial demand is there: 100% of the graduates get a job immediately. The percentage is almost as high for a program that trains medical lab technicians.

2. Schools are in competition with each other and use marks to market themselves. Depending on the quality of the school, this leads to varying degrees of indirect or not-so-subtle pressure into churning out high grades.

3. Education departments create marking policies that lead to unintended consequences. Here are some examples:
a) The passing grade in our province was 50% several decades ago, and it was lifted to 60%, hoping to raise standards. But in more subjective marking courses, the same poor quality exam paper that was graded as a 50%, was merely re-branded as a 60%. Overall averages climbed but reflected no real improvement in the quality of education.
b) More recently, to encourage more lab work in science courses, 40% of the overall grade has to have a practical component. But the lab exams, unlike theory, are not standardized, and in most schools, the lab reports and lab tests (if the latter exist) simply serve to inflate grades.
c) No zero policies dissuade teachers from giving zeros for blank papers. Luckily these highly dubious ideas are being challenged and abandoned in some provinces.

4. Too often the so-called “hard-markers” are stereotyped as being callous to the impact that low grades can have on young, impressionable minds and how, when handed out early in the year, they prevent students from having a fighting chance. The reality is that the “hard markers” have always been a minority, even in the absence of overall grade inflation.  Teachers who realize they’re in the business of both coaching and refereeing, both of which need outside input, both of which need separation at the time of grading, are more likely to be objective markers.
5. Insecurity on part of teachers who are either not tenured or teaching in areas outside of their expertise makes them more likely to inflate marks, either deliberately or unconsciously.It takes experience to create and balance challenging tasks with easier ones. What makes matters worse is that many teacher-training faculties currently use a fuzzy, sociological approach towards teaching and marking and evaluation.
6. Bureaucracies occasionally implement poorly tested programs that place students at a disadvantage. In response, some educators, with good intentions, overcompensate for poor organizational design that is beyond their control and consequently mark too leniently.

7. It’s easy to get away with inflating grades. Most people will not complain about something that seemingly favors them, even if it may be unfair to others. Besides, their vanity may blind them from the fact that there’s something wrong in the first place.

8. Education researcher and retired Duke University professor Stuart Rojstaczer offers an alternative principal reason for grade inflation. In his own words, he believes that from the 1980s onward, students were perceived as customers in search of a degree, not as acolytes who tried to gain knowledge. This cultural shift in colleges brought on student-based course reviews, which unconsciously or not, pressures professors to inflate grades in order to receive favorable ratings from students.

Gradeinflation.com does a more detailed analysis of grade inflation, and also points out the exceptions to the pattern.

It should be pointed out that most college science faculties have been far less impacted by grade inflation. While they should be applauded for maintaining their standards, if other faculties don’t do likewise, it causes some students to shy away from studying such an important field. They gravitate towards the path of least resistance, towards other faculties where grades are more likely to be given away like candy on Halloween.

Specific Heat: a Beautiful Characteristic Property

More than once, I intended to prepare  hard boiled eggs  for my kids in the morning, placed water in a pot but forgot the egg. If I remembered within 30 seconds or so, it was still safe to place my hand in the water, but it was  a bad idea to touch the pot itself. Metals warm up faster than water does—that’s what a low specific heat implies. Specific heat is a characteristic property that measures how much energy it takes to raise the temperature of 1 gram of a material by 1 oC.  The higher the c value is, the more difficult it is to warm up that substance. By the same token, substances with high specific heats also lose their heat with difficulty, while metals cool off with ease.

The high specific heat capacity of water helps temper the rate at which air changes temperature, which is why temperature changes between seasons is gradual, especially near large lakes or the ocean. Water is the reason why Toronto is milder than Montreal.

It’s the reason that the Gulf stream can retain the heat of water heated in the Caribbean and carry it towards London, England,  a city with an average temperature of 7 ºC, even though its latitude is six degrees more North than that of Montreal, whose average January- temperature is only -9 ºC.

If water’s specific heat was lower than it actually is, life would not be possible. The rate of evaporation would be too high, and it would be too difficult for the evolutionary precursors of cells to maintain homeostasis. Why is water special? The reason it is difficult to raise water’s temperature is that hydrogen bonds exist between molecules of water. The hydrogen atoms of one molecule are attracted to the oxygen atoms of other molecules. (see the dots = . . . . in the adjacent diagram. )


Five molecules linked by hydrogen bonds, represented by the four-dot sequence.

Each set represents a hydrogen bond. Five molecules are shown in all)) To overcome this attraction, energy is needed. The bonds between the water molecules are like the links between the wagons of a train. Just like it is difficult to get a big train to reach a high speed, it takes a lot of energy to warm up water. Once the train is moving, it is difficult to stop. Similarly it is difficult to cool water. Of course, in the liquid state, the links are not as fixed as those of train links. As molecules rotate, they constantly break and reform H-bonds, but their overall effect is to maintain attractions within the group.

A very useful formula allows us to calculate the amount of heat either absorbed or lost by a substance during a physical change: Q = m c ΔT , where Q = quantity of heat,  m = mass being heated or cooled, c is the specific heat and  ΔT is the change in temperature.

Now imagine two identical masses, of say, water and copper,  each at the same temperature, receiving the same amount of heat. Since Q is equal, then  

 the masses in the equation cancel, and we end up with:


The inverse relationship graphically represented within the cross reveals that for water’s product cwΔTw to be equal to that of copper, a bigger c  entails a smaller ΔT. For copper it’s the reverse. A smaller c entails a bigger ΔT.

While water’s specific heat of 4.19 J/(g º C) is elevated to that figure by its intermolecular bonds, no such attractions exist between copper atoms. If temperature extremes are excluded, the specific heat of most metals can be predicted from c = 3 R/M, where R = gas constant 8.31 J/mole K and M = molar mass in grams per mole. Since c, in our context , is based on a temperature difference, any  ΔT is equivalent in both C(Celsius) or K(Kelvin) scale of temperature.  c = 3*8.31 J/mole K / (63.5 g/mole)= 0.392 J/(g K )= 0.392 J/(g º C), close to the experimental value of 0.386 obtained by dropping a heated, weighed piece of pure copper into the cool water of a calorimeter and measuring the maximum temperature attained by the water. ( Both the initial hot and cool temperatures also have to be measured).