I should not think so

It will crust up on the walls, completely unchanged, except in consistency.

Actually, I think you might get some, especially if it builds up some heat. I have hard water, and if I boil it, I see a fine, white floc, which I suspect is from calcium carbonate after losing CO2, making CaO or Ca(OH)2. If you are losing CO2, it will form NaOH, and your pH will rise. Are you, by chance, running a water electrolysis unit on a car with a bicarb electrolyte?

I'm playing with one on a diesel which runs mostly on recycled cooking oil. I saw improvement from 24-25 mpg to 28 mpg. Still trying to get it tuned right. My buddy runs these units on a delivery truck fleet, mostly in town, stop&go traffic and saw his fleet mileage go from 10-12 mpg to 16-18 mpg. I guess that makes sense- the unit puts out the same amount of Browns gas regardless of speed, and you get more supplement enrichment of the fuel/air mixture driving slow than fast.

CaCO3 is one of the few materials which is more soluble in cold water than hot.

When you boil your hard water, some of the limestone precipitates out.

Boil some vinegar in the appliance to clean it, as limestone is a very good thermal insulator and will rapidly reduce the efficiency of your pot.

Under normal conditions of electrolysis, the NaHCO3 will simply function as an electrolyte and will be unaffected, the products being O2 and H2.

Hydrogen carbonate salts can react with themselves if heated. My general chemistry students perform an experiment in which they distinguish between hydrogen carbonate salts versus carbonates salts by heating with a flame. The reaction is:

2 NaHCO₃ ---> Na₂CO₃ + CO₂ + H₂O

(by contrast, heating Na₂CO₃ at normal flame temperatures has no effect). Of course this reaction also applies to other Group I metal cations such as Li⁺ and K⁺.

Under electroysis conditions, it is possible that small amounts of CO₂ might be formed by this route, especially if the reaction conditions are extreme due to high temperatures, high voltages, or high hydrogen carbonate concentration. If you want to avoid CO₂ formation, simply use a carbonate salt as your electrolyte instead. Sodium carbonate (Na₂CO₃) is used to raise the pH of swimming pools, and should be readily available at hardware stores (I know that Lowes and Home Depot both have it).

Reaction mechanisms of gaseous reactions very often are different from the liquid phase reaction mechanism as a result of which often the products of even the same reactants occurring in the liquid phase are almost always different. often the possibility of such difference is ascertained through a thermodynamic analysis of the entropy generation status of postulated liquid-phase reaction mecahnism. So I am curious if the possibility you have proffered is based on a postulated reaction mechanism, and whether that postulated mechansim is further verified with a electrothermodynamic analysis.

Or is this just a hunch?

So I am curious if the possibility you have proffered is based on a postulated reaction mechanism, and whether that postulated mechansim is further verified with a electrothermodynamic analysis.

Or is this just a hunch?

Definitely not a hunch -- it's a standard experiment in undergraduate general chemistry. It works just as I described. I've seen it performed many times. But note that the reaction I described takes place in the solid state, at temperatures of a few hundred degrees. Since hydrogen bicarbonate is such a weak base, high heat is needed to initiate the reaction (to provide the activation energy). And the reaction is then assisted by the increase in entropy, as you mentioned (which is true for any reaction that leads to a net increase of gas molecules in an open system). I don't know how much CO₂, if any, could be produced this way in solution since the solution would by necessity be at a much lower temperature (i.e., below its boiling point). Hence my statement that electrolysis conditions (in aqueous solution) "might" produce a *small* amount of CO₂ under "extreme" conditions.

You are now clearer that the reaction most likely will not occur in the liquid-phase as the reaction temperature is much lower and the reaction mechanism would also be different - from the known gas-solid reaction - as the water molecules would most likely have a role in the mechanism. My curiosity was more on the liquid-phase reaction.

Thank you for the clarification.

You are now clearer that the reaction most likely will not occur in the liquid-phase...

Keep in mind that chemical reactions rates of exothermic are not "all or nothing", but rather occur at rates somewhere on the continuum from very slow to very fast. This results from the exponential dependence of reaction rate on the activation energy in the Arrhenius equation. So even in liquid solution, there might still be slow but measurable production of CO₂ from the self-reaction of NaHCO₃.

I know from experience that it's risky to rely on theory alone to make absolute statements about borderline chemical reactions cases like the one discussed here. So the only way to know for sure is to do some experiments. Much of synthetic chemistry practice involves answering the question "what the heck is this unexpected byproduct in my product mixture?" (especially true in organic synthesis). I would bet that a small amount of CO₂ will be formed during electrolysis, and that the solution pH will increase slightly (about 0.5 to 1 pH unit after a few days).

The basic mechanism should be similar in the two pathways mentioned (solid vs liquid). Either way, the unstable intermediate is carbonic acid, H₂CO₃, which spontaneously decomposes to CO₂ and H₂O. By the way, carbonic acid is the dissolved compound that provides the "carbonation" in beer (except Guinness) and soda pop -- it decomposes as soon as the high pressure inside the sealed can is released:

H₂CO₃ ---> CO₂ + H₂O

In the proposed reaction, carbonic acid forms from the acid-base reaction between two hydrogen carbonate ions, an energetically unfavorable and thus rate-limiting step (unfavorable because CO₃^2- is a much stronger base than HCO₃^-, therefore HCO₃^- does not "want" to transfer a proton to another HCO₃^- ion)

2 HCO₃^- ---> H₂CO₃ + CO₃^2- (unfavorable)

As you astutely mentioned, the mechanism of the solution-phase reaction undoubtedly involves solvent water molecules. And of course this should affect the energetics and rate of the reaction. But it might facilitate the reaction, or it might interfere -- it's hard to have much confidence either way based only on speculation. Someone should do experiments to find out (actually someone probably has published results about this scenario, but the experiment should be easier than the literature search).

It appears to me that the concept of ignition and extinction temperatures of reactions is lost in this discussion.

The possible change of the reaction mechanism in the liquid phase is what ultimately becomes the deciding factor. Look there are two cases here to contend with: The lower temperature and the reaction mechanism. You may have provided an argument, albeit tenuously plausible, for the temperature effect for the gas-solid reaction mechanism; but that argument does not extend to the case of the liquid phase reaction with its own reaction mechanism - which is yet to be determined - together with the corresponding ignition temperature of which you have no postulate. Really IMO to attempt to apply that argument that is only merely tenuously plausible even in the gas-solid phase reaction to pure liquid phase reaction, I submit is a stretch.

Reassert the temperature argument only after you have constrauted the reaction mechanism and the associated ignition and exttinction temperature, and may be we may begin to have a common basis for further consideration, I think.

What do you think?

Please define "ignition temperature" and "extinction temperature" -- I don't remember ever seeing these terms used in the context of chemical thermodynamics.

As I already stated at least twice, my suggestion that the liquid-phase reaction might produce a small amount of CO₂ is admittedly speculative. But the exponential dependence of reaction rate on temperature means that we can't categorically rule it out. Due to the presence of a range of collision energies, even at lower-than-normal temperatures some collisions might still have enough energy to reach the transition state, leading to product (of course if the temperature is very low, the reaction rate becomes too slow to detect). But how cold is too cold? I think the reaction rate should gradually approach zero as the temperature is lowered, as opposed to an abrupt change to zero at a specific temperature. Of course at lower temperatures chemical reactions take longer to reach equilibrium, but this is besides the point.

Maybe I don't understand the point of your most recent post. Do you suggest that we can know for sure that the liquid-phase production of CO₂ is impossible (or just too small to measure?) even without performing experimental tests? If so, would you predict the same for a boiling solution-phase?

First of all, let me make the"'words" correction - made bold - in my last post:

• Reassert the temperature argument only after you have constructed the reaction mechanism and the associated ignition and extinction temperature, and may be we may begin to have a common basis for further consideration, I think.

I had just returned from work and was somewhat spent.

Now then, the point of my contention is that the basis for your assertion of the possible liquid phase reaction is the same as the gas-solid phase reaction, and I do not believe this to be so.

First, the solid-gas phase reaction is a pure thermochemical reaction, while the liquid phase reaction is an electrochemical reaction, therefore the controlling thermodynamics are different. So were you to have asserted the situation in the liquid phase that causes the circumstance of the gas phase reaction to obtain then I would have accepted the applicability of your argument of the exponential temperature effect.

Second, the state of pure thermochemical reaction may be occurring in the gas phase exclusively unbeknown to us. Let me try to construct scenarios to explain my thrust here, as I am not always very obvious in these matters.Let us subsume that with respect to the solid-gas phase reaction one mechanistic scenario is that the molecules of the solid sublimate and then undergo a complex reaction of decomposition - and this of course makes sense because of the known/established need for high temperature for the reaction to occur - meaning that the temperature must be high enough to first effect a sublimation as well as then ignite the decomposition reaction. in this sense in order to apply the effect of exponential temperature dependence argument first you have to show that somehow the sublimation or "vaporization" of the salt is, in fact, occurring during the electrolytic process.

Second, the Equations of State of the liquid and the gas are not the same so I suppose that the reaction mechanisms in the phases of state would not be the same as well. Again my thoughts by another scenario: Let us suppose that the salt molecules do not sublimate, but that in some complex reaction mechanism, the oxygen in the air attacks specific atoms of the salts and by stripping the atoms initiates the decomposition reaction. So now you have to show how exactly the oxygen of the air above the electrolytic solution accomplishes the same initiation of the raction, in order for the gas-phase exponential temperature effect supposition have a plausible role.

Finally, and again because the equation of state of the liquid phase of state is different, the molecular level voids within which the salt molecules exist may be such as to cause to obtain entirely different inter-molecular associations as compared to the gas phase sublimated salt molecules, and consequentially different sets of reactions with an entirely different dominant reaction step. So the specific exponential temperature effect often manifest through the activation energy which may characterize the gas phase reaction may not apply for the liquid phase reaction.

I hope that I am clearer now, I often find it difficult to explain these concepts.

BTW, ignition and extinction temperatures are reaction-kinetic concepts often studied in Chemical-physics and Chemical Engineering, and not often considered in standard Chemistry that studies global/net results of reactions and the associated stoichiometric equations.

Since you took the time to write a long reply, I will respond to your main points. Those members not interested in the technical debate might still want to read the last few paragraphs about an experiment I conducted to test my hypothesis.

Reassert the temperature argument only after you have constructed the reaction mechanism and the associated ignition and extinction temperature, and may be we may begin to have a common basis for further consideration, I think.

I find your prerequisite to discussion puzzling since all I said was that the sodium bicarbonate self-reaction might occur in solution at room temperature, but if so probably at a slower rate than the rate observed for the solid-state reaction at flame temperature. This is not a controversial statement, rather simply speculation based on the Arrhenius equation (which states that reaction rate vs temperature is a continuous function that asymptotically approaches zero as the temperature is reduced). The Arrhenius equation is widely known to be accurate for reactions that have non-zero activation energies, which means the vast majority of reactions. And besides, I already gave the mechanism -- here it is again:

Step 1: Acid base-reaction (proton-transfer) between two aqueous hydrogen carbonate ions (possibly assisted by proton transfer via water molecules).

HCO₃^- + HCO₃^- ---> H₂CO₃ + CO₃^2-

Step 2: Decomposition of carbonic acid. This reaction is well known, and has a half-life of less than 10 seconds at room temperature.

H₂CO₃ ---> H₂O + CO₂

The net reaction is

HCO₃^- + HCO₃^- ---> CO₃^2- + H₂O + CO₂

That's it. Not particularly complex or controversial.

Now then, the point of my contention is that the basis for your assertion of the possible liquid phase reaction is the same as the gas-solid phase reaction, and I do not believe this to be so.

The basis for my assertion is that the acid-base reaction I mentioned above occurs both in the solid-state and in solution. This is established fact. I don't understand why you mention a "gas-solid phase reaction" since none is involved, and I never mentioned one. The solid-state reaction involves no gas as a reactant -- CO₂ is a product.

First, the solid-gas phase reaction is a pure thermochemical reaction, while the liquid phase reaction is an electrochemical reaction, therefore the controlling thermodynamics are different. So were you to have asserted the situation in the liquid phase that causes the circumstance of the gas phase reaction to obtain then I would have accepted the applicability of your argument of the exponential temperature effect.

Your argument is making progressively less sense. The solution-phase reaction is definitely *not* an electrochemical reaction, it is an acid-base reaction. Here, look again:

HCO₃^- + HCO₃^- ---> CO₃^2- + H₂O + CO₂

None of the elements involved undergo a change of oxidation state. This is not an electrochemical (reduction-oxidation) reaction! The original question asked whether CO₂ might be produced in an electrochemical cell using sodium bicarbonate as electrolyte. I said maybe, but this does not necessarily mean the CO₂-forming reaction is electrochemical. The hydrogen carbonate self-reaction might occur simply because hydrogen carbonate is present, independently of the water-splitting reaction (i.e., even if the cell has no input current). And heat caused by electrolysis (ohmic heating of the electrolyte) could speed up the hydrogen carbonate reaction.

Second, the state of pure thermochemical reaction may be occurring in the gas phase exclusively unbeknown to us. Let me try to construct scenarios to explain my thrust here, as I am not always very obvious in these matters.Let us subsume that with respect to the solid-gas phase reaction one mechanistic scenario is that the molecules of the solid sublimate and then undergo a complex reaction of decomposition - and this of course makes sense because of the known/established need for high temperature for the reaction to occur - meaning that the temperature must be high enough to first effect a sublimation as well as then ignite the decomposition reaction. in this sense in order to apply the effect of exponential temperature dependence argument first you have to show that somehow the sublimation or "vaporization" of the salt is, in fact, occurring during the electrolytic process.

There is no "gas-phase" reaction involved, so the rest of the above paragraph is a non sequitur.

Second, the Equations of State of the liquid and the gas are not the same so I suppose that the reaction mechanisms in the phases of state would not be the same as well. Again my thoughts by another scenario: Let us suppose that the salt molecules do not sublimate, ...

Sublimation? This is not involved. NaHCO₃ cannot sublimate -- it decomposes to Na₂CO₃ before it could even melt, let alone vaporize to a gas (even the more stable Na₂CO₃ decomposes to Na₂O + CO₂ before it could ever vaporize).

the oxygen in the air attacks specific atoms of the salts and by stripping the atoms initiates the decomposition reaction.

Oxygen is definitely not involved in the reaction. Flame-heating solid Na₂CO₃ inside a evacuated container will cause it to decompose even faster than it does in air.

So now you have to show how exactly the oxygen of the air above the electrolytic solution accomplishes the same initiation of the raction, in order for the gas-phase exponential temperature effect supposition have a plausible role.

Oxygen from the air is not even involved. If you think it is, it is up to *you* to provide evidence. But if you want citations for the reactions I gave at the top of this reply, I can supply some.

Finally, and again because the equation of state of the liquid phase of state is different, the molecular level voids within which the salt molecules exist may be such as to cause to obtain entirely different inter-molecular associations as compared to the gas phase sublimated salt molecules, and consequentially different sets of reactions with an entirely different dominant reaction step. So the specific exponential temperature effect often manifest through the activation energy which may characterize the gas phase reaction may not apply for the liquid phase reaction.

What "gas phase reaction"?! And regarding the "inter-molecular associations" (what I would call "inter-molecular collisions"), I stick to my suggestion that the rate-limiting step is the same in both the solid- and solution-phase: the acid-base reaction between to HCO₃^- ions.

I hope that I am clearer now, I often find it difficult to explain these concepts.

I agree that these are challenging concepts to explain (I know, I teach college chemistry), and I can see that you have some background understanding, but I think you have mixed in some misconceptions.

BTW, ignition and extinction temperatures are reaction-kinetic concepts often studied in Chemical-physics and Chemical Engineering, and not often considered in standard Chemistry that studies global/net results of reactions and the associated stoichiometric equations.

My understanding is that "ignition temperature" and "extinction temperature" apply to oxidation (combustion; reaction with oxygen) reactions. I do remember seeing the term "ignition temperature" referring to the minimum temperature at which a fuel/oxygen mixture will spontaneously ignite. If so, they are not relevant to this discussion which deals with an acid-base reaction.

Finally, I want to report the results of an experiment I did to test my hypothesis. I made a saturated aqueous solution of sodium hydrogen carbonate (NaHCO₃; aka "sodium bicarbonate", "baking soda"). This requires about 7.8 grams per 100 mL of water at room temperature. I refluxed 94 mL this solution (i.e., boiled it underneath a cooling coil which condenses the water vapor back to liquid which falls back into to the solution) inside of air-tight glassware attached to a rubber hose leading to a gas collector (an inverted graduated cylinder filled with water).

I initially thought that the CO₂ production (if any) would be very slow, and had planned to heat overnight. But as the solution started to warm (well before reaching the boiling point of 104 deg C), I observed gas bubbling into the collector. Note that this gas is not steam -- it passes through the condenser and exits the glassware below room temperature. Gas evolution rate reached its maximum rate at the solution's boiling point, then gradually tapered off over 1 hour 30 minutes, at which time the reaction seemed complete. I collected 1080 mL of a colorless odorless gas. I did not characterize it further, but it should be CO₂ since it is very unlikely that I split water into hydrogen and oxygen. Again note that this gas is at room temperature -- it cannot be hot air or steam.

Theory predicts that the proposed reaction should produce about 980 mL CO₂ gas, so the result matches the hypothetical prediction fairly well (my excess gas volume probably resulted from not measuring the mass of NaHCO₃ more precisely). But wait! I have one additional piece of supporting evidence. I measured the pH of the initial NaHCO₃ solution and of the final reaction mixture. The pH increased from 8.42 to 9.96, consistent with the conversion of NaHCO₃ to the more basic (alkaline) Na₂CO₃.

-------------

Theoretical Yield Calculation:

94.0 mL sat'd solution x (7.8 g NaHCO₃/100 mL sat'd solution) = 7.332 g NaHCO₃

7.332 g NaHCO₃ x (1 mol/84.007 g) = 0.08723 mol NaHCO₃

0.08723 mol NaHCO₃ x (1 mol CO₂ /2 mol NaHCO₃) = 0.04364 mol CO₂

0.04364 mol CO₂ x (22.4 L gas/1 mol gas) = 0.978 L CO₂ gas

------------

It was a very quick, easy, and satisfying experiment. Next I will try it without heating. Will the reaction show a measurable reaction rate at room temperature? Arrhenius says "maybe".

I encourage others to replicate my results. Here's a simplified version of the experiment: stir excess baking soda in water to make a saturated solution. Let it settle for an hour, then carefully decant some clear solution and half-fill an empty beer bottle. Secure an empty rubber balloon over the bottle's mouth. Lower the bottle into a pot of hot water for 1 hour. Observe. Pull the bottle out of the pot and let it cool to room temperature. Observe. Please share your results with the group.

Lets find out! Can you check the pH of your solution initially, after it runs an hour or so and warms up and run it a week or so and check again? NaHCO3 - CO2 ->NaOH, which will raise the pH.

Lets find out! Can you check the pH of your solution initially, after it runs an hour or so and warms up and run it a week or so and check again? NaHCO3 - CO2 ->NaOH, which will raise the pH.

Good point wcfloyd! Na₂CO₃ is definitely more basic than NaHCO₃, so the if the former is being produced by the reaction that I previously mentioned, then the solution's pH should go up. I proposed a hypothesis that makes a falsifiable prediction, and you suggested a straight-forward empirical test -- the scientific method in action!

All users of Browns Gas please review. http://ntrs.nasa.gov/archive/nasa/casi.ntrs.nasa.gov/19770016170_1977016170.pdf This would imply an improvement in exhaust emissions. So be a friend of the environment and possibly save money :)

Maybe ohmic heating in an electrolysis cell could cause some bicarbonate to decompose according to the reaction below?

I did the boiling-bicarbonate experiment before checking if others had already done it, thinking that performing the experiment would be easier doing a literature search. But it turns out that other experimenters have already done the solution-phase decomposition, and it didn't take long to find hits in a Google search, for example

http://newsgroups.derkeiler.com/Archive/Sci/sci.engr.chem/2005-07/msg00016.html

----

Sodium bicarbonate

From Wikipedia, the free encyclopedia

http://en.wikipedia.org/wiki/Sodium_bicarbonate#Thermal_decomposition

Thermal decomposition

Above 60 °C, it gradually decomposes into sodium carbonate, water and carbon dioxide. The conversion is fast at 250 °C:

2NaHCO₃ → Na₂CO₃ + H₂O + CO₂

>Calcium carbonate is one of the few salts that is more soluble in cold that hot water.

No it isn't! It isn't very soluble in either, but my chemical handbook says .0014 g/100ml at 25 C and .002g at 100 C. The deposition of calcium carbonate in kettles is due to the decomposition of calcium bicarbonate, which is considerably more soluble, into the carbonate and carbon dioxide. The classic limestone cave reaction.

As to the experiment with boiling sodium bicarbonate, the reaction is well known and is mentioned in many of the old chemistry texts. Partington (Text book of Inorganic Chemistry, 1950) , for example, says " By prolonged boiling practically all the bicarbonate is converted into carbonate, and if the crude bicarbonate from the ammonia-soda process is boiled with water, the ammonium salts present as an impurity are decomposed and ammonia is evolved. On recarbonating, by passing in carbon dioxide, nearly pure sodium carbonate is precipitated, and the commercial salt is made this way."

It is a shame that nobody has done the electrolysis experiment that was the origin of this thread.