More than a century ago, there was a famous chemical reaction that could supply (combustible) fuel gas to cities, which was based on first turning coal into coke, and then passing steam over white-hot coke:
C(s) + H2O -> CO + H2
The gas mixture that resulted, of carbon monoxide and hydrogen gas, was both a source of fascination before the year 1900, and a reason why some city blocks ended up blowing themselves up, over the first Historic gas leaks.
What some people may not know, is that this chemical reaction was already accompanied by another reaction:
CO + H2O -> CO2 + H2
That coincidentally resulted in the gas-mixture containing a higher amount of H2 than it would contain CO, just because there was already steam present, when the water-gas was made. This second reaction is called the Water-Gas Shift Reaction.
There is one Historical context in which this reaction was always useful (as the cited article already points out). There have always been industrial consumers of hydrogen, who could not afford hydrogen produced by hydrolysis, but who could afford hydrogen that is partially contaminated with carbon monoxide. BTW, the CO2 can be removed from the gas mixture, by passing it over quicklime, producing some form of calcium carbonate. The only problem with using quicklime in the modern era is the fact that quicklime itself is produced primarily, by heating calcium carbonate, thereby releasing the equivalent amount of CO2 in advance… (:1)
But there is a context in which this reaction is almost useless. It uses a large supply of coal as its energy source, and also produces CO2 as a byproduct, the latter of which needs to be controlled, if a population of cars is to be fuelled with hydrogen.
I suppose, though, that some countries could be so desperate for an alternative source of energy, or so determined, that they’d consider trapping all the CO2 thus produced, before selling hydrogen to fuel-cell-powered -car owners. Those would probably be countries that also have a lot of coal as a resource. Those countries might also need to refine their fuel-cell technology, into fuel-cells which won’t become poisoned by the carbon monoxide still present in the hydrogen, which quicklime would not trap.
I tend to think of this reaction as a trip down memory lane.
Carbon monoxide reacts with methanol to form acetic acid, which can in turn be trapped by quicklime, just as carbon dioxide once was. In this case, the precipitate will be calcium acetate. Therefore, special fuel cells might not be needed.
(Update 2/18/2020, 18h05 : )
(As of 2/12/2020 : )
It’s also possible to pass an electric current through solutions of NaCl or CaCl2, in order to obtain, among other reaction products, NaOH or Ca(OH)2. Because these last two substances are bases, they will also capture CO2 out of various mixtures of gasses. However, if the assumption is to be made that enough electrical energy is available to generate the relevant quantities of these bases, then it will probably be more efficient just to generate the H2 through hydrolysis.
In fact, when NaOH or Ca(OH)2 is being generated through the electrolysis of the respective brine, H2 is also generated, which suggests that the amount of energy needed to do so cannot be less, than what would generally be needed to generate H2. One reason is the fact that the electronegativity of Cl2 is approximately the same as that of O2, actually resulting in similar, positive electrode potentials. Another is the fact that the electrode potential of Na(m) or Ca(m) is negative with respect to H2.
! However, I think that an important observation to make about these suppositions would be, that The amount of H2 that would need to be generated ‘inefficiently’, through the generation of NaOH or Ca(OH)2, would only be half the amount of H2 generated via the water-gas.
Also, I’ve commented elsewhere in this blog, that NaOH is “thermally stable”, which means that no matter to what temperature it’s heated, it will stay NaOH, perhaps dissociating completely, but then returning as NaOH when the temperature is reduced again. AFAIK, the thermal behaviour of Ca(OH)2 is different, in that if heated to a sufficiently high temperature, it can be made to separate directly into CaO and H2O. Therefore, in general, another pathway does exist to manufacture quicklime. And yet, this quicklime will not perform better at capturing CO2, than Ca(OH)2 already would.
(Update 2/13/2020, 22h45 : )
The standard way I was taught NaOH is manufactured – and this was in the late 1970s / early 1980s – was, that NaCl brine was fed to an electrolytic cell, the anode of which was made of carbon, and at which Cl2 gas formed. The cathode of this cell was made of mercury, and metallic Na would dissolve in it, forming “sodium amalgam”.
What was special about this cathode was that its mercury was made to flow between this cell, and another cell that contained water, with which the sodium in the mercury would react spontaneously, to form both H2 gas and NaOH that would dissolve in the water.
There is no reason why a calcium homologue would not work.
If metallic sodium was desired as an end-product, obtaining it was a question of extracting amalgam from the cathode, and boiling off the mercury – in the absence of oxygen, of course.
(Update 2/14/2020, 16h30 : )
If the reader has been paying close attention to all the implications of this posting, then he will notice that in order to succeed, at what the subject line suggests, a good, working “Carbon Sink” would be needed, even though, to the best of my knowledge, none has been discovered so far, in the real world. And the carbon sink would need to be an artificial one.
The next question becomes, ‘Would it be better to dump calcium carbonate as a waste product, or sodium carbonate?’ Because, what this posting proposes is, that the carbon which fuelled the water-gas, should be dumped in one of those two forms, so that clean hydrogen fuel can be derived on a large scale. At first glance, calcium carbonate would seem an obvious choice, but this posting also tries to point out, that an arbitrary choice of reagents is not always available, those reagents potentially being derived products of other reagents, derived in a way that emitted CO2 again.
The advantages of using CaCO3 include that this is a non-toxic solid at ambient temperatures, and insoluble in water. But, the ability to generate CaCO3 as a waste product would depend on the salt being used as an input, that is commonly known to the population as CaCl2, or, as calcium chloride.
CaCl2 is derived in practical Chemistry, as a byproduct in the production of ‘baking soda’, through the “Solvay Process”, which in turn consumed calcium carbonate, and stored its carbon in (sold) ‘NaHCO3‘. It would be safe to assume that, after it has been purchased, 100% of all baking soda re-releases its CO2 into the ecosystem as such. (:2)
The advantages of aiming to use Na2CO3 as the carbon waste-product include, that ordinary NaCl could be used as an input, which could even be derived from sea water, and that therefore, a virtually unlimited amount of water-gas could be generated, leading to the required quantities of H2, and a virtually unlimited quantity of Na2CO3 could be dumped. Yet, if the reader imagines wide and long dunes of sodium carbonate being dumped, that is highly soluble in water, it would only follow that, as soon as rain falls on those dunes, the CO2 is on its way back into the ecosystem.
In fact, household instructions on how to dispose of Na2CO3 involve, eventually flushing it down a water drain, which again, leads to this same result…
It would be ideal if a large quantity of (powdered) Na2CO3 could be mixed with ‘a small quantity of something else’, so that an insoluble product would form, which could be dumped on a large scale, which would not decompose, yet which could be left that way safely. I have yet to find that ‘something else’.
I can be slightly more creative in my thought process and suggest the following:
When the NaCl was electrolyzed, in order to generate NaOH and H2, Cl2 (chlorine gas) was also generated, which might be a useful reagent on a smaller scale, but which was of no use so far, in the hypothetical water-gas process being considered. At the same time, coal was converted into coke, which is a process that drives out the volatiles in the coal in the form of “coal tar”, being a sticky residue when back at ambient temperatures.
Theoretically, the Cl2 could be forced to react with the coal tar, especially since coal tar is largely a mixture of unsaturated hydrocarbons. In bulk, this would lead to an ugly, smelly, toxic mixture of unpredictable composition, yet, a mixture of chemicals that are generally electrophiles, that should react readily with Na2CO3. Hence, if this mixture was added to the sodium carbonate, some sort of amorphous substance should form, that is also a polymer.
Only, such a polymer could eventually cause toxic residues to form, when left buried, which could turn it into a long-term Eco-Hazard.
(Update 2/18/2020, 14h20 : )
According to a WiKiPedia article entry, the U.S.A. has perfected a method of extracting purified CaCl2 from sea-brine. It was one of those WiKi-articles tagged with “Citation Needed”. If this is true, then the process which I proposed in this blog posting is closer to being viable.
However, the way an anonymous Industrial Chemist has replied to this concept was, that “The concentration of Ca2+ ion in sea-water is limited, by the solubility of CaCO3 (since there is also some CO2 in the sea-water), and this solubility is close to zero.”
If somebody had asked me a month ago, I’d have said that there is bountiful magnesium in the sea-water, but little calcium. Logically, one can design an ion-exchange process to be as clever as possible; it will only extract calcium if calcium is present.
In order to understand this question properly, one needs to have learned what the concept of “an equilibrium constant” is. But the question can be rephrased as, ‘Why can Ca(OH)2 be used to scrub CO2 out of a solution, if the solubility product of Ca(OH)2 is only on the order of 5.5·10-6?’ And the answer could be, ~Because the solubility product of CaCO3 is on the order of 3.3·10-9.~ The Ca(OH)2 succeeds because its alkalinity increases the concentration of (OH)- ions, so that the concentration of Ca2+ ions can become as high as the cube root of 5.5·10-6, which is, 1.7·10-2, if nothing neutralizes the alkalinity. If the solution remained neutral, then the concentration of (OH)- -squared would remain 3.1·10-11 (Moles2/Litres2), which, being smaller than 5.5·10-6, means that the level of Ca2+ ions could become arbitrarily high. In that situation, solid Ca(OH)2 would never precipitate. However, if the concentration of Ca2+ was truly, ‘arbitrarily high’, then the reason would be, the possibility that it went into solution due to more than one reagent, not only, due to Ca(OH)2 being dissolved. And so, this observation really explains why, if ?CaCl2? was dissolved in large quantities, solid CaCl2 will also precipitate, according to its own solubility product, but not solid Ca(OH)2. (:3)
The way in which equilibrium constants generally work is, they state the product of concentrations, of each reaction product, divided by the product of concentrations, of each reagent, and will be valid at one temperature. We were taught how to apply them for gasses, because those examples are easiest to work with, but they also apply for any non-gaseous concentrations eventually… Even though additional chemical reactions can change the concentration of one reagent or reaction product, the equilibrium constant stays the same. Only, the concentrations of the other reagents and reaction products will adapt.
Solubility constants are similar to equilibrium constants, except for the fact that they state the constant of the applicable dissociation reaction.
Because the equilibrium constant with which H2CO3 converts into 2H+ + [CO3]2- remains constant (2.1·10-17), depleting the concentration of H+ will increase the possible concentration of [CO3]2-. Specifically, if the concentration of H+ becomes as low as 1.8·10-9 (Moles/Litre, corresponding to a pH of 10.5), then this value squared will become 3.2·10-18, which is smaller than 2.1·10-17, and the concentration of [CO3]2- can become as high as 7.8·10-3. At that point, the solubility constant of CaCO3 being very low, will in fact complete the job, of removing the [CO3]2- ions, and thus also, of depleting the supply of CO2.
If the reader needs to solve the riddle, of [CO3]2- ions having ‘an arbitrarily high concentration’, when
the maximum concentration of anything is (1), then the way to do that would be, to recompute what the real concentration of the reagent, CO2, must be, from a known [CO3]2- concentration, as well as from the known H+ concentration / acidity level. Their product, taking care to square the concentration of H+, can be divided by 2.1·10-17, to arrive at the real concentration of H2CO3. If this was greater than 1.2·10-3, then CO2 gas would come bubbling out.
Well, an oversimplification of the concentration of CO2 in sea-water, or the concentration of Ca2+ ions, should not occur. The constant with which any quantity of CO2 converts into H2CO3 is only 1.2·10-3, (1.7·10-3 for fresh water), and the real amount of it present in sea-water as [CO3]2- is only 5.6·10-11. Because this last number is ~59x smaller than 3.3·10-9 (the solubility constant of CaCO3), in principle, the level of Ca2+ could again, be any level (this time, without solid CaCO3 precipitating)…
According to one source, the real level of Calcium in sea-water is 1·10-2 !
When working with equilibrium constants, I assumed that all concentrations are in Moles / Mole. This is also referred to as a “Molar Fraction”. I supposed that was why, when I was taking CSCS Chemistry courses, we restricted ourselves to working with gas volumes, when studying equilibrium constants. When working with other measures of concentrations, the appropriate conversions would need to be calculated, for example, to Moles / Litre, and then to Grams / Litre.
1 Mole of H2O, having ~18g of mass, takes up ~0.018 Litres of volume. This also means that 1L contains 5.55·10+1 Moles of H2O. Small Molar Fractions can be multiplied by that number, in order to arrive at Moles / Litre, which is also known as “the Molarity of a solution”. What can make this calculation messier is the fact that the solute displaces the solvent. I.e., if the Molar Fraction of Ca(OH)2 was 1.0, then the result would no longer follow that 5.55·10+1 Moles of the solute are present… What this means is that the Molar Fraction of the solvent must be derived from those of the solutes(, unless the computation is only an approximation for what happens at low solute- concentrations / Molar Fractions). And then, this solvent- Molar Fraction needs to be multiplied by 0.018L (for H20), to which the volume of the solutes would need to be added… After that, I think, the Molar Fraction of each solute must be divided by this calculated total volume in Litres, to recompute how many Moles / Litre of each solute are present. From that, Grams / Litre can be computed relatively easily.
In order to compute the volume of a pure reagent or reaction product, its Molar Quantity needs to be multiplied by its Molar Mass, and the product divided by its density. And a problem which my Chemistry courses did focus on was, that a mass of reagent had been measured on a scale, while that reagent was assumed to be buoyant, displacing a corresponding volume of air. The mass of this volume of air would need to be added to the measured weight, to arrive at the assumed mass of reagent, but not necessarily in that order. What this meant was that, in addition to performing Gram-Molar conversions ad nauseam, we were also required to compute volumes from time to time.
Most of the casual computations in our Chemistry courses were in Moles / Litre.
(Erratum 2/18/2020, 16h20 : )
As it turns out, to avoid the problems that I wrote above, the standard practice in Chemistry on Planet Earth is, to state solubility products on the assumption that all the concentrations are in Moles / Litre, but then, to omit any statement of a unit which would follow.
While this greatly simplifies the Math, it also has as side effect that the maximum concentration of a species in aqueous solution can become greater than 1 ! This is because, as I wrote, pure water has an equivalent concentration of 5.55·10+1 Moles / Litre. Theoretically, the concentration of dissolved ions could become equally high, if the solubility products of the salts they form did not cause their precipitation first, thus limiting the actual concentrations dissolved.