Sciencemadness Discussion Board - Remarkable reaction with adjustable delay

Remarkable reaction with adjustable delay

woelen - 7-11-2013 at 00:04

By accident I found a very interesting reaction, not because of its reaction products, but because of a really spectacular transition from seemingly no reaction at all to near explosive violence. The experiment involves mixing two solutions and then for tens of seconds or even a few minutes nothing seems to happen and after that, suddenly a very violent reaction starts, which completes in a fraction of a second

The experiment is very simple.

Take appr. 100 mg of NaBrO3 and dissolve this in just enough water. Add water drop by drop and try to dissolve this small amount of NaBrO3. Assure that all solid is dissolved, no solid particles may be left behind. It is easiest to use a dry test tube, put the NaBrO3 in that and then add water drop by drop, each time swirling until no more NaBrO3 dissolves and continue doing this until no solid is left.

Take appr. 100 mg of NH2OH.HCl (hydroxylammonium chloride) and dissolve this in just enough water in the same way as you did with the NaBrO3. Dissolving the NH2OH.HCl requires less water than dissolving the NaBrO3.

Mix both solutions. Simply add them to each other in a test tube, swirl and then set the test tube aside and step back a meter or so.

Wait (this can take a minute, or even a few minutes) and then suddenly within a fraction of a second a near-explosive reaction occurs. The water boils away and a small amount of brown vapor is produced. The entire reaction takes less time than a blink of the eye (actually, you can't follow it with your eyes, it goes very fast, maybe a tenth of a second and then all is over).

You can play around a little with quantities. E.g. take 150 mg of NH2OH.HCl and 100 mg of NaBrO3 and you get NO/NO2 after the reaction. Take just 50 mg of NH2OH.HCl and 100 mg of NaBrO3 and you get bromine vapor after the reaction. In all cases, there is a delay, which can range from appr. 1 minute to well over 5 minutes. If you hand-warm the liquids while dissolving the solids, then you need less water and the concentration of the solutions is somewhat higher, and waiting time is shorter. If you dilute each solution by a factor of 1.5 before mixing, then you get a much longer waiting time (between 5 and 10 minutes) before the reaction starts.

An interesting variation is to dissolve 100 to 200 mg of NH2OH.HCl in water and add some solid NaBrO3 (50 to 100 mg) to the solution. Again, for a minute (or even a few minutes) nothing seems to happen, the NaBrO3 just sits there at the bottom, under the liquid. Then there is a POP sound, due to a small explosion and a plume of vapor and brown gas escapes from the test tube.

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I think this is very remarkable, because of the long delay, relative to the duration of the reaction. This is a very special type of dynamics. Such violent reactions are not that special (e.g. 65% HNO3 with isopropylalcohol can be equally violent when mixed), but it is the ratio

delay : (duration of reaction)

which surprises me. The delay is hundreds or even thousands times as long as the duration of the reaction.

With all other reactions of this violence, the delay is at the same time scale as the duration of the reaction.

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

Finally, a warning: This reaction can be carried out safely with the indicated quantities, but do not scale up!!

[Edited on 7-11-13 by woelen]

woelen - 7-11-2013 at 04:18

In the meantime I googled "bromate hydroxylamine". A lot of links appear which talk about complex dynamics and one link talks about clock reactions:

http://www.readcube.com/articles/10.1002/kin.550260305

I also did the experiment with (NH2OH)2.H2SO4 instead of NH2OH.HCl and this shows exactly the same type of behavior. I also did experiments with KBrO3 instead of NaBrO3. This shows similar behavior, but this only works with solid KBrO3 (this is much less soluble than NaBrO3 and solutions of this probably are too dilute).

I did not read complete articles (no access), but the search makes clear that there is something interesting about this reaction.

bbartlog - 7-11-2013 at 04:23

Any speculation about the nature of the delay (or indeed the overall net reaction)? Presumably there is some initial slow reaction that ultimately creates conditions for the fast reaction to take place, but what is it? A gradual increase in the pH as the HCl is oxidized?

woelen - 7-11-2013 at 04:43

I did not yet speculate (nor read) anything on the nature of the reaction. I want to think that over tonight, but I did not yet find the time to do that.

blogfast25 - 7-11-2013 at 05:46

Quote: Originally posted by woelen

I did not yet speculate (nor read) anything on the nature of the reaction. I want to think that over tonight, but I did not yet find the time to do that.

Is it possible that this is just a plain old runaway? Reaction proceeds very slowly at RT but generates heat. Slow temperature rise slowly increases reaction rate, thereby increasing heat output and further accelerating reaction rate and so on and so on until reaction rate become extremely high?

You could test that hypothesis by adding the second reagent to the first but while the first is hot (or warm) (and by some form of remote control!)

froot - 7-11-2013 at 06:01

Or maybe see if placing the test tube into a ice water bath after mixing affects the delay.

woelen - 7-11-2013 at 06:25

@froot: Heating certainly affects the delay. I tried that. I took the saturated solutions (which are around 15 C, the same as my lab ambient temperature) and mixed them. I did the same experiment again, using samples from the same solutions, but then I warmed the liquids in my hands before mixing them. This is a difference of 15 degrees or so. In the latter case, there still is a delay, and the characteristic of the reaction is the same, but the delay is shorter. So, temperature affects the delay.

@blogfast25: An ordinary runaway is different. As you say already, in a runaway you have accelerating rate and it keeps on accelerating on and on. I have seen runaways quite a few times (e.g. with nitric acid and organics, but also while making peroxochromates from H2O2 and chromates) and then you see a slow reaction, which gradually becomes more violent, e.g. faster bubbling, until it runs out of control. In this reaction there is no visible change at all and then suddenly there is production of a lot of gas (I think most of it is N2) and the liquid becomes hot. The rate of production of the gas, however, is not increasing visibly, it just goes POP. During the entire delay, nothing visible happens at all! And then suddenly it does POP, and then the reaction is over. I hope to be able to make a movie of this next weekend in the daytime with good light at 200 frames per second. Then I get more info on the duration of the reaction. It certainly is much less than a second, faster than a blink of the eye.

deltaH - 7-11-2013 at 07:00

I believe that clock reactions were greeted with great skepticism when they first appeared because many scientists argued that such chemical kinetics went against known chemical behaviour (of the type blogfast describes)... particularly the ones that 'tick' forwards and backwards... was considered total pseudoscience lol

Clear cut case of 'nature not understanding the theory'

Now here's a real challenge for you woelen. Can you demonstrate a clock reaction coupled to chemiluminescence? Preferably a ticking one would be much more spectacular over a off then on one, but even the latter would be wonderful nevertheless.

[Edited on 7-11-2013 by deltaH]

blogfast25 - 7-11-2013 at 11:01

Woelen: you're probably right but a very steep exponential style increase in reaction rate could still be confused visually with a hockey stick type phenomenon... At the end of the day, even during the delay, something (invisible or poorly visible) has to be going on.

DubaiAmateurRocketry - 7-11-2013 at 11:34

Have you tried making Hydroxyl ammonium nitrate ?

bbartlog - 7-11-2013 at 12:10

Quote:

a very steep exponential style increase in reaction rate could still be confused visually with a hockey stick type phenomenon

Except that the overall reaction in this case has a very steep inflection point that doesn't match any single reaction rate function. Anyway, the obvious way to test this idea is to measure the temperature progression of the reaction, see what the temperature is just before it goes poof, and then *start* the reagents at that temperature. If it's really just a single stage, then it should take almost no time at all; if there's some unperceived priming reaction, then there will still be an appreciable initiation period. I assume that's why woelen warmed his reagents: to demonstrate this test.

Ethylene glycol and bleach (6% aqueous sodium hypochlorite) also react in a delayed, initiation-and-then-reaction way. But it's quite boring compared to this.

woelen - 7-11-2013 at 12:11

Quote: Originally posted by DubaiAmateurRocketry

Have you tried making Hydroxyl ammonium nitrate ?

Let's try to stay on-topic. Discussing the synthesis of hydroxylammonium nitrate is interesting, but not in this topic.

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I tried whether this clock-type kinetics of this reaction is specific for bromate or not. And yes, it is.
1) When KIO3 or NaIO3 is added to a solution of NH2OH.HCl, then immediately a violent reaction occurs. A plume of iodine vapor escapes from the liquid immediately.
2) When NaIO4 is added, the behavior is very similar to the case of adding NaIO3 or KIO3.
3) When NaClO2 is added, then you see a more classical runaway. First, there is gentle bubbling and foaming for a few tens of seconds, but slowly the bubbling and foaming intensifies and at a certain point in time it becoms quite rapid and then in second or so, the reaction is over in a final climax, being a puff of white fume and gas (water vapor and most likely N2).

Another direction of research is to try to find a set of differential equations, which shows clock-type behavior. This certainly will be a non-linear set of equations. The set also must be simple. Only two reactants are involved, so the equations cannot have many states. The final property of these equations is that the involved quantities must be non-negative, as they should represent concentrations of chemicals.

Unfortunately, my search on internet does not yield very much more info than I already have. Quite a few articles exist, which mention special kinetics in the hydroxylamine/bromate system, but much more information is not available freely.

@bbartlog: I have been thinking about measuring temperature of the reaction mix, but I only dare perform the experiment on a very small scale and for decent temperature measurements you need at least a few ml of solution. Next weekend I want to try the reaction with lower concentrations of chemicals in somewhat larger volumes. I already did this with small amounts and then I had a waiting time of more than 7 minutes (!!) and after that it did *whoosh* and the reaction was over. After the whoosh, the test tube was hot.

[Edited on 7-11-13 by woelen]

blogfast25 - 7-11-2013 at 12:29

bbart: nowhere did I imply a single reaction rate.

Woelen: do you have access to the full article?

What do you mean by "so the equations cannot have many states"?

Differential equations could be the way to go but with sets of simultaneous non-linear equations, computer iterative programs would almost certainly be needed.

There are some really simple clocks too: KI/H2O2/Na2S2O3/starch is a classic. No non-linearities involved.

Metacelsus - 7-11-2013 at 12:44

I found an interesting article that may explain it:

Quote:

A kinetic study of oxidation of hydroxylamine by bromate ion in acid sulfate solution using spectrophotometric and potentiometric methods is reported. Oxidation of hydroxylamine to nitrate is quantitative and followed competitive, consecutive, and auto catalytics steps characterized by induction periods. In the slow rate limiting step, hydroxylamine on reaction with HOBr (kmath image) forms an intermediate I, which further reacts fast with second molecule of HOBr (kmath image) giving nitrite. Nitrite reacts with HOBr (kmath image) yielding the final product nitrate. Nitric acts as an autocatalyst also and its initial addition decreased the induction periods. In excess of hydrogen ion concentration all the reaction steps follow second-order kinetics. All the second-order rate constants are reported and the reaction mechanism is proposed.

http://onlinelibrary.wiley.com/doi/10.1002/kin.550161011/abs...

Sadly, I don't have access to the full text.

woelen - 7-11-2013 at 12:46

I do not have access to the full article. I'll look further, but I have no real hope for that.

With 'many states' I mean many independent state variables, i.e. the dimension of the system.

I know of the simple clock, but that is not representative for this. The clock involves Na2S2O8, Na2S2O3, KI and starch and works like this:
Na2S2O8 is a strong but slow oxidizer. It oxidizes I(-) ion slowly to iodine, while not oxidizing S2O3(2-). As long as S2O3(2-) ion is present, the iodine is reduced at once to I(-) again and the S2O3(2-) is oxidized to S4O6(2-). As soon as all S2O3(2-) is used up, the iodine remains in solution and then gives a dark blue complex with starch. So, the liquid remains colorless for a while and then at once turns dark blue. The dynamics of this is described by highly non-linear equations, but the behavior of these equations is simple. There is no sharp inflexion point in the reaction, everything runs smoothly. It looks a sharp reaction, due to the sensitive nature of the starch indicator, but the underlying dynamics are tame and predictable.

The reaction I describe in this thread, however, must be of quite a different nature. There is a very sharp rise in reaction rate in an amazingly short amount of time. The solution of the differential equation system hence switches from minute's time scale to millisecond's time scale and that makes the system so special for me.

vulture - 7-11-2013 at 12:47

Some thinking points:

Did you try the other reactions with the same quantities of water? Maybe the boil off of the water is required before the reaction really takes off? Does bromate show anomalous solubility compared to the other oxidizers?

PHILOU Zrealone - 7-11-2013 at 12:48

@Woelen,
To keep things relative to each other you should have tested NaClO3 and not NaClO2.
I think you would also have a delayed reaction with chlorate.

The reaction is for sure due to the formation of HBrO3 and resultant NH2OH.HBrO3
Hydroxylamine bromate contains a strong reducer and a strong oxydant in the same molecule... just like NH4ClO3 it is on th edge of stability.
Probable cause of the delay of explosion is the slow building up of concentration of NBr3.

woelen - 7-11-2013 at 13:06

I tried the NaClO3 as well, but this is not interesting at all. It does not react and this is exactly what I expected. My excuse for not mentioning that. I considered it common knowledge that chlorate ion is sluggish in aqueous solution and only acts as serious aqueous oxidizer at low pH or at elevated temperatures, which cannot be achieved in aqueous solution at normal pressure. I can imagine that finely powdered solid KClO3, mixed with finely powdered solid NH2OH.HCl can make a very energetic mix, but this is far from the conditions of the special reaction which I observed.

I asked access to the paper, mentioned by Cheddite Cheese. It looks promising. I hope to be able to derive the set of differential equations if I have access to the paper and have insight in the reaction mechanism. Simulating a set of differential equations is not a problem for me. If they are in the form dx/dt = f(x), with x a vector of state variables then things are really simply (e.g. Runge-Kutta with variable step size detection to find the sharp inflexion point), if they are implicit algebraic differential, i.e. of the form dx/dt = f(x,y); g(x,y) = 0 with algebraic state y, then I have more of a challenge. This may require DASSL or a similar piece of software.

I do not think that stuff like NBr3 or NH2OH.HBrO3 is involved. There must be a much more intricate mechanism.

@vulture: I did quite a few different tests, with solid KBrO3 and solid NaBrO3 and solutions of that. All show clock-type behavior (except too dilute solutions). The solubility of the bromates is not really anomalous compared with the other oxidizers. The sodium salt dissolves more easily than the potassium salt.

[Edited on 7-11-13 by woelen]

blogfast25 - 8-11-2013 at 06:02

Quote: Originally posted by woelen

I asked access to the paper, mentioned by Cheddite Cheese. It looks promising. I hope to be able to derive the set of differential equations if I have access to the paper and have insight in the reaction mechanism. Simulating a set of differential equations is not a problem for me. If they are in the form dx/dt = f(x), with x a vector of state variables then things are really simply (e.g. Runge-Kutta with variable step size detection to find the sharp inflexion point), if they are implicit algebraic differential, i.e. of the form dx/dt = f(x,y); g(x,y) = 0 with algebraic state y, then I have more of a challenge. This may require DASSL or a similar piece of software.

Well, I certainly look forward to that treatment!

woelen - 8-11-2013 at 14:57

A little bit of the mystery is resolved. The reaction is not a true clock reaction, but it is a so-called branching chain reaction (see link for a very short, but clear qualitative description):

http://www.britannica.com/EBchecked/topic/77570/branching-ch...

A branching chain reaction is a chain reaction, where one of the products in the chain of reactions catalyses the conversion of one of the initial reactants. It has aspects of an autocatalytic reaction, but from initial reactant(s) to catalysing product is not through a single step, but through a chain of steps.

In the case of my reaction, the formation of bromide ions and H(+) ions catalyses the reaction.

The net reaction (when excess NH2OH.HCl is present, which is split into ions NH3OH(+) and Cl(-) in solution) is:

2BrO3(-) + 6NH3OH(+) ---> 2Br(-) + 3N2O + 9H2O + 6H(+) + heat

This net reaction occurs through many steps, this equation only gives the final result.

Initially, when there is no bromide and only a little amount of H(+), due to splitting of NH3OH(+) in NH2OH and H(+), bromate reacts with NH3OH(+) very slowly, giving intermediate species HBrO2 and [NH2(OH)2](+). Both species react further very rapidly, giving final products Br(-), N2O, H(+), water.

With the increase of the concentration of Br(-) and H(+) another reaction occurs, much more rapidly, but not so fast that it can be considered momentaneous:

BrO3(-) + 5Br(-) + 6H(+) --> 3Br2 + 3H2O (simplified, through a chain of steps, involving HBrO2 and HOBr as transient species).

The Br2 in turns reacts rapidly with NH3OH(+), forming N2O, water, bromide and H(+). This reaction can be considered momentaneous (quasi-static, non-differential).

Many steps go so fast, that they can be considered non-differential and can be considered quasi-static. A few steps have a perceptible range and these cannot be described as quasi-static and lead to a differential equation.

The type of equations for a branching chain reaction lead to non-linear equations, exhibiting super-exponential behavior. Such systems can have solutions which remain close to 0 for a long time and then suddenly the solution 'explodes'. In technical terms, such a solution is called a non-thermal explosion.

I'll try to find a set of differential equations, based on real chemical reactions, which on simulation reproduces the effect which I observe.

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

Another thing, which makes the system even more extreme is that in the short time of the non-thermal explosion the liquid heats up considerably (close to the point of boiling) and this accelerates the reaction even more. The produced heat also works "autocatalytic" in some sense. In a normal runaway, the heat is only one factor, which leads to the fast reaction and this can lead to exponential runaway. Here, in this reaction, the non-thermal explosion, described above, combined with the strong heating of the reaction, gives the effect of the sudden explosion.

A nice article, describing some of the concepts is attached to this post.

If I find a good set of equations, which describe this reaction and reproduces my experimental results, then I certainly will make a web page about this interesting phenomenon, with references, pictures and a Java-program which demonstrates the behavior.

Attachment: clock_reactions.pdf (69kB)
This file has been downloaded 556 times

[Edited on 8-11-13 by woelen]

woelen - 9-11-2013 at 07:24

I made a few movies of this reaction:

Excess bromate: http://www.homescience.net/chem/exps/hydroxylamine_bromate/e...
Excess hydroxylamine: http://www.homescience.net/chem/exps/hydroxylamine_bromate/e...
Slow motion: http://www.homescience.net/chem/exps/hydroxylamine_bromate/s...

bfesser - 9-11-2013 at 07:40

<del>woelen, I've taken the liberty of uploading your videos to YouTube (unlisted), so that they can be embedded here. Some users don't have native AVI support on their devices; this is a way around that. As I said, they're unlisted, so they're only really visible from these links, but I can remove them if you like; just say the word.</del>

[edit] Amazing work, by the way! The slow motion video in particular shows some surprising and unexpected (for me) phenomena.

[edit] Videos removed.

[Edited on 9.11.13 by bfesser]

deltaH - 9-11-2013 at 07:47

Spectacular! I enjoyed the comparison of the excess one or the other. Lovely work, well done woelen!

[Edited on 9-11-2013 by deltaH]

kmno4 - 9-11-2013 at 08:24

I tried this on test tube-scale, ~50mg reactants + ~1cm3 water
Violent reaction , however without bromine, only colourless gases, but it is a matter of proportion I think. In my hands it took few seconds from mixing reactants to the spectactular end.
I also tried N2H4•H2SO4 instead of NH2OH•HCl.
Reaction similar, but it speeds up much longer (with increasing amount of bubbles), and the end is not so violent. This time I got at the end orange solution with strong Br2 smell.

When some solid KBrO3 was added to spent solution (with excess of hydroxylamine), no delay was observed but sudden reaction starts at once.

[Edited on 9-11-2013 by kmno4]

woelen - 9-11-2013 at 08:29

For this time it is OK to me to leave them on Youtube, but generally I am not charmed of it at all. The movies are screwed, all detail is lost and video quality is crappy. I want people to watch the original movies and if they are on Youtube, then people do not watch the originals. Especially the slow motion video loses a lot of detail.

Another issue with Youtube is that the material becomes owned by Youtube as soon as you upload. If somewhere in the future I see my movies (or frames of it) used (e.g. in a commercial), then I cannot do anything against that. For that reason I have my own webspace. If people do not have AVI support on their PC, I recommend them to install VLC media player, which is available for Windows and Linux (not sure about Apple OS-X) without any cost.

watson.fawkes - 9-11-2013 at 08:53

Quote: Originally posted by woelen

A branching chain reaction is a chain reaction, where one of the products in the chain of reactions catalyses the conversion of one of the initial reactants.
[...]
The type of equations for a branching chain reaction lead to non-linear equations, exhibiting super-exponential behavior.

A chain reaction always has an expanding propagation step with more than one unit of output per unit input. Rapid growth in a chain reaction occurs when the expansion rate is greater than the extinction rate. They can also be modeled with a first-order, linear differential equation. Chain reactions exhibit exponential growth, not super-exponential. Another physical example with a similar mathematical model is the ionization shower created by an energetic charged particle.

Now these differential equations are generally in more than one variable, not the single variable ones most commonly encountered early in calculus courses. "It's complicated" is not at all the same "it's nonlinear". Multidimensional linear equations can have rather surprising behavior, and these include sudden state changes. Large classes of reaction rates can be modeled with continuous-time Markov chains, and it takes some argument to claim that this class of equations cannot model the system in question.

PHILOU Zrealone - 9-11-2013 at 11:52

Good examples of chain reactions are:
H₂ + Cl₂ --> 2 HCl
2 H₂ + O₂ --> 2 H₂O

In a chain reaction there are 3 phases:
-Initiation
-Propagation
-Termination
If the propagation reaction is fast enough or induces more and more reactive species then the reaction can be explosive.

Sometimes initiation of the reaction happens because of:
-the glass surface
-the temperature
-the pressure
-the light
Hydrogen and oxygen can be made to explode at ambient T° in a glass tube, simply by increasing slowly pressure of the gas mix in the tube.

Delay reaction are typical of a mix of many reducer-oxydiser couples:
All are temperature and concentration dependant:
-Formol + HNO₃
-Aceton/propanon + HNO₃
-Glycerol/propantriol + KMnO₄

Being such a mix I stil think that the mix of NH2OH.HCl and NaClO3 would form a mix that is spontaneously explosive with a given delay because it would form NH2OH.HClO3 what is very close and related to NH4Cl and NaClO3 mix what forms the infamous NH4ClO3.

@Woelen
In your videos in real time, the reaction is very fast and starts between two frames.

In your slow motion it is funny to see the reaction starts in the bottom right and goes up to the left like an explosion.
Also the crystals seems to take rocket propelling behaviour making funny loopings

Very noticeable are the NOx colour in the exces hydroxylamonium movie and the Br2 droplets in the excess bromate movie.

Very nice videos!

[Edited on 9-11-2013 by PHILOU Zrealone]

woelen - 9-11-2013 at 12:02

I do not agree with you, watson.fawkes. The description of this type of reaction definitely is non-linear and I perfectly know the difference between complicated and non-linear! Multidimensional linear equations have nothing really surprising for me (at least the time-invariant ones which can be described as dx/dt = Ax, where x is the state vector and A is a matrix), simply determine the eigenvalues and you know the characteristics of the solution.
Actually, nearly all chemical reactions must be described by non-linear equations, because you have products of concentrations in the equations. Only if a concentration can be considered (nearly) constant, behavior may be approximated by linear equations, but for adequate explanation of branched chain reactions (and also of the related oscillating and chaotic reactions) you need to have a non-linear model.

I know that the state-vector of the system I am studying is multi-dimensional. It is both non-linear and multi-dimensional and that can be quite an interesting situation. From dimension 3 and up the system can exhibit chaotic behavior, but even from dimension 1 you can get clock-like behavior.

In the meantime I have done some more reseach already and have done some math. The basic idea behind the dynamics of the system I have is the behavior of the system dx/dt = x*x (or somewhat more generic: dx/dt = x^a, with a > 1). Such systems have solutions which remain less than 1 for a long time and within one second they flash from less than 1 to infinity.
Of course, this is an oversimplification of the system described here, but it captures the non-thermal explosion perfectly. In reality, there are not unlimited amounts of reactants, and the equation is more complicated, but the basis is there. I'll try to get a better set of differential equations and keep the order of the system as low as possible (reaction steps which proceed very fast can be regarded as quasi-static and only add algebraic equations, not additional differential equations).

What is special about this chain reaction is not the fact that it is a chain, but the fact that one of the products, produced in the chain enormously speeds up the initial reaction. It catalyses the reaction. The catalyst is bromide ion and another catalyst is H(+) ion. I tried this, by adding NaBrO3 to a solution of the hydroxylammonium salt to which some KBr is added as well. The addition of KBr has a marked effect on the induction time. It becomes much shorter. Addition of H(+) (e.g. a single drop of 2M H2SO4) before adding the NaBrO3 also has a strong effect. When this is done, then a violent reaction starts nearly immediately after adding the NaBrO3, but the reaction is less spectacular. It looks more like a normal violent runaway.
Temperature also has a strong effect. As I wrote before, when I handwarm the solution before adding NaBrO3, then the induction period is much shorter than when I use cold tap water (which is appr. 15 C this time of the year).

[Edited on 9-11-13 by woelen]

blogfast25 - 9-11-2013 at 14:09

Woelen:

Very interesting that you were able to prove bromide and H<sub>3</sub>O<sup>+</sup> (you didn't mean H<sup>+</sup> literally, right? There is water present) have a catalytic effect.

Re the equations, I'll reserve judgement until someone comes up with the most correct set (I assume some simplifying assumptions will have to be made) but I agree that unless some reagent concentrations are large with respect to others (and thus quasi-constant), most equations governing the system will be non-linear. They have to be, relying on non-linear reaction equilibrium constants as they do.

Have you got any links up on your website to these vids?

woelen - 10-11-2013 at 00:28

@blogfast25: I wrote H(+), but of course this is a simplification. Even H3O(+) is a simplification, albeit a much better one already.

I have not yet made a webpage about this reaction. I made the videos, so that I can show some results already, but the math must be much more rigorously established. I also want to show a simulation, which shows behavior, similar to the real reaction.

deltaH - 10-11-2013 at 02:07

woelen, I'm curious, have you tried this reaction with an iodate as oxidant by chance?

[Edited on 10-11-2013 by deltaH]

watson.fawkes - 10-11-2013 at 08:10

I spoke too hastily above about the reaction rate equations being linear. I was principally objecting to the hyperbole involved in describing them with phrases such as "highly non-linear". Too inflamed to think through the situation, I posted too soon. In a mathematical sense, reaction rate equations are anything but highly non-linear. They are, indeed, about as simple as its possible to be and not be linear. "Highly nonlinear" would be equations that include terms such as (∂f/∂y)², but there's no hint of that here.

Generically, reaction rate equations are in a class of equations called "semi-linear". This is one class of equations in a sequence: linear --> semi-linear --> quasi-linear --> nonlinear. (The buzzwords are all there for searching, but it's some serious mathematics behind them.) Semi-linear equations are defined for PDE's, not just the simple temporal ODE's we've been considering here. (Simple because we assume no spatial difference in concentrations of reagents.) These are first order (only the first derivative) and the derivative coefficients are constant. Furthermore, the nonlinear terms are simple polynomials in the concentration coefficients, not arbitrary functions. And in most cases, the equations are even component-wise linear, say, under the assumption of dilute solutes.

First-order semi-linear equations have solutions by the so-called "method of characteristics". This method applies to PDE's of arbitrary dimension, so it can handle proper chemical engineering problems, where concentrations vary with position. The essence of the technique is find a coordinate transform that "straightens out" some of the coordinates. The characteristic curves are well-behaved and it's straightforward to get numerically stable computational solutions. The point I'm making is that these reaction rate equations are mathematically relatively simple even in the PDE case, much less the ODE case.

"Super-exponential" is another phrase I object to. The fastest that any solutions grow (locally) is bounded by an exponential of some polynomial function. That's in distinction to shock formation, where you do get super-exponential growth at the shock in the neighborhood where the solutions ends. As I understand it, this class of reaction rate equations here cannot form shocks, though I'd need to sit down and prove it to be sure. (Shock formation is at the center of the Clay prize about the Navier-Stokes problem; it's not a particularly well-understood area of mathematics.)

woelen - 10-11-2013 at 10:35

Quote:

The fastest that any solutions grow (locally) is bounded by an exponential of some polynomial function.

Local analysis is not the solution to problems like this. You need global analysis. Actually, global analysis also is a term from mathematics, it is about describing dynamical systems with (possibly) non-euclidian state space (sometimes quite complicated manifolds), but I do not specifically refer to that branch of mathematics. The use of differential geometry is not needed for solving this kind of problems.
The term "highly non-linear" refers to the type of solutions, needed to describe systems like this. Even a simple first order non-linear equation like dx/dt = a*x + x*x already has quite interesting behavior for certain initial conditions and certain values of a, which I would describe as "highly non-linear".

blogfast25 - 10-11-2013 at 11:06

Quote: Originally posted by woelen

Even a simple first order non-linear equation like dx/dt = a*x + x*x already has quite interesting behavior for certain initial conditions and certain values of a, which I would describe as "highly non-linear".

Do you mean: dx/dt = ax + x² ? Like x(t)^' = ax + x² ?

t = (1/a) ln[x/(x +a)] + C (C constant)

Yup, that could get interesting for some a.

Or am I missing some notation here?

[Edited on 10-11-2013 by blogfast25]

woelen - 10-11-2013 at 11:26

Yes, I meant what you write. I am a software nerd and I am spoiled so much that I use * for multiplication, even in non-software contexts

. On the other hand, software engineers are fond of notations, using simple ASCII characters only on a single line, for even the most draconian mathematical expressions

Let me rewrite the thing: dx/dt = a*x + x^2

[Edited on 10-11-13 by woelen]

watson.fawkes - 10-11-2013 at 13:27

Quote: Originally posted by woelen

Well, then, I will just continue to protest and view this use of language with disdain. "It's complicated" is just not the same thing as "highly non-linear".

There are highly non-linear equations out there, with far more structure, such as the recursion operators in certain integrable systems such as the KdV (Korteweg-deVries) equation that map solutions to other solutions. Such operators devolve to trivial (identity) operators in simpler situations. Even here, though, inverse scattering transforms can linearize these equations, in a certain sense, so they are not as non-linear as others out there.

In short, non-linearity is not about the external behavior of the solutions, but about internal structure. Chemical rate equations are just barely non-linear.

Pok - 10-11-2013 at 17:40

Fantastic videos, woelen!! I first watched the slow motion video with these wonderful "bubble rockets". I did't imagine that in real time it happens soooo suddenly.

The reaction reminds me of a similar experiment which I did. The colour change and the sudden evolution of gas:

<iframe sandbox frameborder="0" width="480" height="270" src="http://www.dailymotion.com/embed/video/x14acdf"></iframe><br />

It's part of this experiment: a luminol clock reaction. Unfortunately, this reaction is not nearly as "shocking" as your one:

<iframe sandbox frameborder="0" width="480" height="270" src="http://www.dailymotion.com/embed/video/x149yl0"></iframe><br />

[Edited on 11-11-2013 by Pok]

deltaH - 11-11-2013 at 01:48

I'm sorry, those luminol clock reaction's are simply too beautiful for words...

[Edited on 11-11-2013 by deltaH]

woelen - 11-11-2013 at 05:01

I found an interesting paper, which describes how one can determine in general when a dynamical system, described by polynomial differential equations, shows "blow-up" in its solution:

http://www.math.bme.hu/~csikja/files/v26-43.pdf

This is exactly the type of phenomenon which is observed in this reaction. I also made some progress in my own determination of suitable equations:

There are two rate-determining steps in the reaction:

1) BrO3(-) + NH3OH(+) + H(+) --> HBrO2 + NH2(OH)2(+) (slow reaction)
2) BrO3(-) + Br(-) + 2H(+) --> HBrO2 + HOBr (quite fast reaction, but not instantaneous)

The first reaction has rate [BrO3(-)]*[NH3OH(+)]*K, where K can be written as K = k*[H(+)]/(A + B*[H(+)]), with k, A, and B constants.
The second reaction has rate c*[BrO3(-)]*[Br(-)]*[H(+)]², with c being some constant.

Further reaction from HOBr and HOBr2 with excess hydroxylamine to bromide and N2O are very very fast.
The intermediate species NH2(OH)2(+) decomposes to N2O and water and H(+). This reaction also is very fast.

The reason for the complicated order for H(+) in equation (2) is because of the presence of a very fast equilibrium reaction between NH2OH+H(+) and NH3OH(+). Such equilibria lead to algebraic (non-differential) constraints.

I have references to papers for all of these reactions, unfortunately I do not have them with me here. Lateron in subsequent posts, I will post the references, so that other can check my findings.

[Edited on 11-11-13 by woelen]

deltaH - 11-11-2013 at 05:56

I have some humble questions about these kinetics woelen.

Firstly, I am confused about this statement that two reactions are rate limiting, by definition can there not be only one? Do you mean to say that two are candidates and it's unclear which actually is?

Secondly, these reactions do not appear elementary steps as written, is it not therefore incorrect to call either a rate limiting step (the true rate limiting step being an elementary reaction step in a sequence of steps which those equations describe)?

Thanks.

[Edited on 11-11-2013 by deltaH]

woelen - 11-11-2013 at 06:21

My wording indeed is not entirely clear. I meant to say that there are two steps in the reaction which add dynamics on such a timescale, that it cannot be described as an instantaneous reaction. These reactions lead to differential equations, while all other steps in the total net reaction lead to algebraic equations (examples of the latter are the equilibrium equations for acid-base reactions, involving the Kz of acids).

The first reaction seems indeed to be a really elementary step. It is described in two of the papers I used in my research and in both the rate is described as c*[BrO3(-)]*[Br(-)]*[H(+)]². This fits the stoichiometry of the reactants, used up in the reaction, the coefficients are 1, 1, 2 and these are exactly the exponents in the reaction rate.

The second reaction is somewhat more involved. There is interaction with an acid-base equilibrium between NH3OH(+) and NH2OH+H(+). This leads to the somewhat more complicated form, presented in my previous post, due to some algebraic constraints on the concentration of H(+). Precise details will follow in my write-up.

There are many more steps, the total net reaction is bromate plus hydroxylammonium ion gives bromide plus N2O plus water plus acid. Because all these steps use up hydroxylamine and finally produce bromide ions, H(+) ions and N2O, there is a positive feedback to the two rates, given above. Especially the formation of bromide leads to a spectacular speedup. This is something I confirmed with a little experiment in which I added bromide before adding bromate. The experiment of kmno4 also is instructive. He added bromate to the spent solution of hydroxylamine (which contains excess hydroxylamine) and when the bromate is added to this, then there immediately is a violent reaction. This is because a lot of bromide and acid is present in the spent liquid and then there is immediate fast production of HBrO2 and HOBr, which in turn rapidly react with excess NH2OH.

[Edited on 11-11-13 by woelen]

deltaH - 11-11-2013 at 07:05

Thanks.

I think there is some confusion here with nomenclature then and wording.

To help clear it up, I'll paraphrase the IUPAC definition that an elementary reaction step can only consist of one transition state.

As a result, they are very simple things, generally of the form A + B => C or D => E + F. Your equation has three reactants, which tells me that it is a combination of elementary steps as written.

In other words, all three reagents do not need to simultaneously collide to form the two products by a single reaction step (and don't, the probability of this happening statistically is insignificant).

Much more likely, two need to collide to form some hidden intermediate and then this collides with your third reactant to form the products, or even more steps than this, but you get the picture.

I was also under the impression (maybe wrong) that only an elementary step could be a rate limiting step, i.e. the slowest step causes all others to be assumed to be at equilibrium.

Links:
http://en.wikipedia.org/wiki/Elementary_reaction

[Edited on 11-11-2013 by deltaH]

woelen - 11-11-2013 at 07:41

Quote:

[...]the slowest step causes all others to be assumed to be at equilibrium.[...]

This cannot be true, because if this indeed were the case, then every chemical reaction would behave like a first order system (only one differential equation, all other equations being algebraic, because all other steps are in equilibrium).

What I have understood about elementary reactions is that they can involve more than 2 species. The chance that a certain species is in a certain small piece of space is proportional to the concentration of that species. Assuming independent motion of all species, this is true for all involved species. So, the chance that all required species for a reaction are in a certain small volume of space is the product of all concentrations. If N molecules (or ions) of the same species are needed, then the product for that species is simply its concentration to the power N. So, the chance of three entitities (or in my case, even 4 entities) is not negligible, although at low concentrations of all entities it can be very low. If, however, the species are reactive, then the constant k for the reaction may be high and this may compensate.

From experimental data it is known that the rate of the reaction between bromate, bromide and H(+) ions is as described in my previous post and this is in perfect agreement with what I described above. I can only imagine a process where all 4 ions are close to each other in a small space, but if you see another mechanism, which could lead to the same rate exponents (1 for bromate, 1 for bromide, 2 for H(+)), then I would like to know that. For the discussion and the explanation of the observed behavior, however, this is not important. What is important is the observed dynamics from experimental data and the equations describing this dynamics accurately. It is nice to know the precise physics behind this, but if this is not possible/unknown, then I still can do the math and try to describe the total reaction. It then becomes more descriptive, but that is a valid approach in natural sciences.

deltaH - 11-11-2013 at 08:19

Okay, I concede that describing chemical kinetics by short posts is a nightmare! Ok there is too much to have to describe here and I am too poor a describer to pull this off concisely.

[Edited on 11-11-2013 by deltaH]

PHILOU Zrealone - 11-11-2013 at 09:37

Quote: Originally posted by deltaH

woelen, I'm curious, have you tried this reaction with an iodate as oxidant by chance?

[Edited on 10-11-2013 by deltaH]

See post of Woelen posted on 7-11-2013 at 21:11 in this tread

watson.fawkes - 11-11-2013 at 09:42

Quote: Originally posted by woelen

I found an interesting paper, which describes how one can determine in general when a dynamical system, described by polynomial differential equations, shows "blow-up" in its solution:

"Blow up" is essentially the same thing that I referred to above as shock formation (the detailed differences don't matter for the present topic). There's a rather useful section for the present discussion therein:

Quote:

Suppose we have a mass conserving reaction endowed with mass action type kinetics. Then, no solution of the induced kinetic differential equation with nonnegative initial condition blows up.

Thus, given a particular class of equations that looks a lot like the right class of equations to model a system, we may well find that it contains behaviors that don't actually occur. In other words, given a class of first-order semi-linear differential equation that varies from linearity only by a quadratic form (this is the class of equations at the start of this paper), we see that some blow up and some do not. Making a further restriction (see the quote) eliminates the blow up. What we should conclude is that the unrestricted class of equations is the wrong class to model the phenomenon of interest, even though it looked like a good candidate at the outset. Such is the progress of science.

Differential equations are chock full of this kind of thing. Things that seems close to each other turn out not to be very similar at all. Overall, the reason for this is that there are many incompatible notions of "close to". Mathematically, "close to" is formalized as a topology. Closeness in, say, ordinary 3-dimensional space is (mostly) unique, but utterly non-unique in infinite dimensional spaces, and even more so in spaces with more than one kind of infinite extent such as differential equations. If anyone would like a taste of this field, just take a gander at the Wikipedia page on jet bundles; jet bundles are the appropriate structure in which to define coordinate-invariant differential equations (the only kind expected to model physical reality). That page is mostly definitions; it doesn't even get to topological issues.

deltaH - 11-11-2013 at 09:43

aw... missed that, thanks PHILOU.

PHILOU Zrealone - 11-11-2013 at 10:05

Quote: Originally posted by woelen

I tried the NaClO3 as well, but this is not interesting at all. It does not react and this is exactly what I expected. My excuse for not mentioning that. I considered it common knowledge that chlorate ion is sluggish in aqueous solution and only acts as serious aqueous oxidizer at low pH or at elevated temperatures, which cannot be achieved in aqueous solution at normal pressure.

I think that water solutions of HONH3Cl are relatively acidic just like solution of NH4Cl are.
NH4Cl is already enough acidic to ensure strong oxydising properties of chlorate mixed with it in water solution and favourise the slow NCl3 formation...
In the case of HONH3(+) it must be a stronger acid than NH4(+), just like H2N-NH3(+) is.

woelen - 11-11-2013 at 10:49

Quote: Originally posted by watson.fawkes

Here you have a very important point! The type of equations, you mention indeed cannot show blow-up, because they describe the reaction of finite amounts of reagents. In my experiments, the very violent reaction looks like a blow-up, but it is not. Although it is very violent, it is not infinitely violent and does not involve infinite amounts of reactants. The total amount of each of the elements in the system is constant.
So, I need to search for strong splike-like behavior instead of true blow-up. Another approach may be to use approximate systems, where blow-up is an approximation of a very strong spike, but if such a solution is obtained, then one knows that it is not a true description of reality, but only a (coarse) approximation.

Thanks for providing this piece of insight!

blogfast25 - 11-11-2013 at 12:17

Like Delta, I'm sceptical about elementary reactions that require collisions between more than two species.

deltaH - 11-11-2013 at 14:06

Quote: Originally posted by woelen

Quote:

[...]the slowest step causes all others to be assumed to be at equilibrium.[...]

Ok, I'll give this a shot and probably make a total ass of myself... I hated chemical kinetics, nevertheless, here goes my attempt at refuting this

I thought the best way to illustrate my point concisely is by a 'simple' example (turned out not to be simple at all, but it's as concise as I could make it).

Consider the hypothetical reaction with the same form as yours:

a + b + c => d + e

Let us say that this consist of the set of elementary reaction steps where rxn 1 is rate limiting (thus others being in eqbm.):

(1) a + b => f
(2) f + c <=> g
(3) g <=> d + e

Note: Specie 'f' and 'g' are therefore 'reactive intermediates'. I am not saying your equation works like this, this is simply made up for illustrative purposes and to keep things as simple as possible.

We can write rate equations for each:

r1 = dCf/dt = kf1.Ca.Cb – kb1.Cf
r2 = dCg/dt = kf2.Cf.Cc – kb2.Cg = 0: K2.Cf.Cc – Cg = 0 by using K2 = kf2/kb2
r3 = dCg/dt = kb3.Cd.Ce – kf3.Cg = 0: Cd.Ce – K3.Cg = 0 by using K3 = kf3/kb3

Eliminating for the transient specie ‘g’: Cd.Ce – K2.K3.Cf.Cc = 0
Rearranging: Cd.Ce/(K2.K3.Cc) = Cf
Using this in r1 to eliminate the transient specie ‘f’:
d[Cd.Ce/(K2.K3.Cc)]/dt = kf1.Ca.Cb – kb1[Cd.Ce/(K2.K3.Cc)]

The rest of the maths is simply tedious (chain rule differentiation, pulling out constants and grouping constants to make new constants), but my point is you don't land up with a first order equation.

What is incorrect here please?

[Edited on 11-11-2013 by deltaH]

watson.fawkes - 11-11-2013 at 21:54

Quote: Originally posted by deltaH

I'll [...] probably make a total ass of myself... I hated chemical kinetics, nevertheless, here goes my attempt at refuting this
[...]
What is incorrect here please?

You have certainly succeeded at your originally stated goal.

Incorrect? This is in the not-even-wrong category. "First-order" means something different in chemical kinetics than it does with respect to differential equations. woelen's statement is perfectly meaningful because it refers to a first order chemical reaction, not to a first-order differential equation.

This link and this link should help.

deltaH - 11-11-2013 at 22:35

watson.fawkes, you need to understand the meaning of the insult before you can use it effectively. The 'not even wrong' category refers to statements that are not falsifiable. For example, a statement such as "invisible gnomes adept at hiding all traces of themselves frolic in my garden at night" is clearly not falsifiable, so that would indeed be in the category of 'not even wrong'.

However, I have provided a mathematical derivation which certainly can be falsified if incorrect, but you have not done so.

Implying that I do not understand the difference between the order of a rate law and the order of a differential equation does not change the order of the derived rate law.

This is a classical strawmanning strategy.

The fact remains that my derived rate law is not first order because the power of concentration terms are not one and so I have demonstrated by exception that woelen's assertion, that setting all fast steps to be in equilibrium results in the reduction of the rate law to first order, to be incorrect.

Well, that is at least what I tried to do

[Edited on 12-11-2013 by deltaH]

woelen - 12-11-2013 at 00:37

Please, no quarreling in this thread, I don't like that at all!

@deltaH: Your equations assume equilibrium, but my equations all are about far from equilibrium situations, except the NH3OH(+)<-->NH2OH+H(+) equilibrium. Algebraic constraints also come from reactions which are very fast compared to other ones. Such reactions simply can be propagated to their final products and the rate at which these final products appear are algebraic expressions of the concentration of intermediate products.

@blogfast25: I myself also have the feeling that the 4-particle reaction is not the best description of the reaction between bromate, bromide and hydrogen ion. So, I actually agree with you and deltaH. But I see no better explanation. And observed kinetics also is according to this model, so for the time being, I use this at least as a mathematical model, describing the kinetics quite well, realizing that although this mathematical model does describe the reaction quite well, it probably is not a true reflection of the underlying physics. Most likely, the system has more state variables, but the other equations are very stiff (have much smaller time scales) so that the observed behavior within the accuracy of the measurements is like the equation I presented in earlier posts.

Even now, while I still have no complete describing model, I learned already quite a few new things about chemical kinetics. It is a highly descriptive branch of science, which heavily relies on macroscopic observations and for which it is amazingly difficult to elucidate all intermediate steps. The real mechanisms behind the transfer of atoms between reactants is a fascinating and largely unknown area of science.

@PHILOU Zrealone: I have done some more experiments with KClO3, but it really does not show any interesting reaction. I even provoked the goddess of chemistry a little, but she did not bite this time

:
I (carefully) tried boiling a fairly concentrated solution of (NH2OH)2SO4 with KClO3. No visible reaction. To that I added a little 2M H2SO4, while still hot. Still no visible reaction. Finally, I added a little NaCl. When this is done, there is some slow production of gas, but nothing interesting. Most likely the chloride with chlorate at low pH slowly gives some Cl2 and/or ClO2 which immediately react with hydroxylamine to give N2 and/or N2O.

[Edited on 12-11-13 by woelen]

Marvin - 12-11-2013 at 04:00

Quote: Originally posted by woelen

Quote:

[...]the slowest step causes all others to be assumed to be at equilibrium.[...]

Actually it does make sense, all the other reactions have time to get close to equilibrium. It doesn't mean the equilibrium point doesn't change throughout the reaction. If the products of the reaction affect the equilibrium, or catalyse the rate determining step then the reaction overall will be high order dynamically. I think the general feeling is that during the induction phase of this reaction the pH drops, leading to increased concentration of the species involved in the rate determining step. It's an Autocatalytic reaction.

Quote: Originally posted by blogfast25

Like Delta, I'm skeptical about elementary reactions that require collisions between more than two species.

I'm not. If nothing else, a third body is often needed just to take away the energy. Any reaction that happens on human time scales must be pretty unlikely on the molecular level.

watson.fawkes - 12-11-2013 at 08:11

Quote: Originally posted by deltaH

The 'not even wrong' category refers to statements that are not falsifiable.

Excuse me. I managed to have forgotten that you live in a Humpty-Dumpty land where everything you say means exactly what you choose it to mean, neither more nor less.

watson.fawkes - 12-11-2013 at 08:37

Quote: Originally posted by woelen

Algebraic constraints also come from reactions which are very fast compared to other ones. Such reactions simply can be propagated to their final products and the rate at which these final products appear are algebraic expressions of the concentration of intermediate products.

Just because a reaction is fast doesn't mean it generates a constraint. As above, this is a question of topology of the equation. For a fast reaction to generate a constraint, that means that every solution of the original equation has to lie near the solution set defined by the constraint. This may be a property of specific equations, but it is not a property of these systems in general.

Indeed, with regard to the original reaction, the "instant" phase change is evidence of a very fast rate constant somewhere in the system. If you were to convert that immediately into a constraint, you'll end up with a class of candidate functions for solutions that do not contain that rapid phase change behavior.

What I am seeing here is that you have two very fast rate constants in the system, one for a reaction that goes forward in some sense, and another that goes in reverse. The rate constant for the reverse reaction is faster than for the forward one. The combination of these two looks like nothing is happening for a while, until some input of the reverse reaction is exhausted, and suddenly the forward reaction proceeds without inhibition.

morganism - 14-11-2013 at 12:43

Are you doing this experiment at a window ?

It is possible that this could be a nanoparticle creation/reaction.

Lots of studies of water purification with nanoparticles are happening for developing countries.
I may have posted some up here.

They have found that nanoparticles absorb energy, creating a vapor bubble around themselves, preventing heat from escaping. this can lead to vigorous boiling, but havn't heard of anything this rigorous from that model yet...

Sunlight is enough to run these reactions, don't know if a halide would.

[Edited on 14-11-2013 by morganism]

blogfast25 - 14-11-2013 at 13:42

Quote: Originally posted by morganism

I may have posted some up here.

Like where precisely?

woelen - 16-11-2013 at 09:35

I made quite some progress. I derived a set of equations, which wonderfully captures the observed behavior. Initial addition of bromide ion shortens the induction period, but the reaction remains sharp and sudden. Initial addition of H(+) ions also shortens the induction period, but the reaction also becomes more like a normal runaway. This behavior occurs in reality, and the equations also capture that behavior.

The model has 5 state variables (concentration of bromate, bromine, bromide, hydrogen ion and hydroxylammonium ion).

I also made a few more movies. One movie with sound, so that you can hear the POP sound and a movie with mixed liquids instead of added solid to liquid. The mixed liquids actually seem more impressive than solid added to liquid. Most of the liquid simply evaporates immediately after start of the reaction.

More will follow tomorrow or the day after tomorrow. I also want to write a webpage about these phenomena.

Mailinmypocket - 19-11-2013 at 05:39

Came across a somewhat similar experiment, albeit not as violent or complex. Thought I would share it here in case there is interest

http://www.versuchschemie.de/topic,17874,-chemischer+Geysir....

woelen - 19-11-2013 at 07:24

I finally have finished the webpage about this experiment. It has a lot of pictures, a few videos and a mathematical model of the reaction dynamics, which explains the observed behavior quite well:

http://woelen.homescience.net/science/chem/exps/hydroxylamin...

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

Your link from versuchschemie is quite interesting as well. Something I certainly will try myself. This reaction, however, is not nearly as violent as the one I have (re)discovered. I also want to warn people not to scale up my reaction to the size of the reaction, described on versuchschemie! That would almost certainly lead to really nasty accidents!

blogfast25 - 19-11-2013 at 10:08

It's an interesting one. It seems to be autocatalytic (on superficial inspection)

woelen - 21-11-2013 at 12:18

I tried a few other similar reactions.

- The one on versuchschemie.de (Na2S2O5 with solid KClO3): It works as described. The reaction becomes faster and faster over time and ends with a climax in which the liquid becomes boiling hot.
- N2H4.2HCl with solid NaBrO3: This gives a violent, but not explosive, reaction at once. No induction time at all.
- N2H4.2HCl with a drop of 30% hydrazine added and then addition of solid NaBrO3: No reaction occurs at all, not even after ten minutes. When a drop of 30% HCl is added, then there is slow increase of reaction speed and a runaway occurs.
- (NH3OH)2.H2SO4 with solid NaBrO3: There is an induction period, followed by a small explosion (POP sound), which is only slightly less violent than with NH3OH.HCl. This most likely is due to the somewhat lower solubility of the sulfate salt, so that the concentration of the solution is somewhat lower.

It appears that NH3OH(+) ion has some specific property which the other reductors do not have: N2H6(2+) immediately seems to react quickly with bromate, while NH3OH(+) only reacts slowly with bromate.

woelen - 24-10-2018 at 12:25

After 5 years I revisited this reaction again. I now have a camera which can make movies with 1000 frames per second at good resolution. The movies are really really stunning.

Have a look at these two videos:

http://www.homescience.net/chem/exps/hydroxylamine_bromate/N...
http://www.homescience.net/chem/exps/hydroxylamine_bromate/N...

Five years ago I already made a webpage about this very special reaction. The videos are not yet in the webpage, I will add links to them soon. The webpage is the following:

http://woelen.homescience.net/science/chem/exps/hydroxylamin...

I am still struggling with finding good video editing software. I now had to reduce the resolution and bitrate a lot to make videos of acceptable dowbload size (still appr. 15 MByte), but this reduces the quality a lot. The videos are fairly small (360x640), the originals from the camera are 1080x1920, but the original files are almost 1 GByte, not suitable for web publishing at all. I think, however, that the videos, I have uploaded, show the stunning effect of the reaction sufficiently well.

Each video covers appr. 2.5 seconds of real time. The start of the reaction already is very fast on these super slow motion videos, in reality it is in the milliseconds time scale.

Waffles SS - 24-10-2018 at 12:57

Woelen,
I get excited when i read your posts.

That is great,
How you find out this reaction mechanism?you spent a lot of time for providing dynamics of this reaction on your website.
after 5 year you are still interested to this reaction and want to complete information about it.You are really talented chemist.

[Edited on 24-10-2018 by Waffles SS]

woelen - 25-10-2018 at 02:17

In the derivation of the equations I used a few papers. They can be downloaded on my webpage (there are links in the webpage to these papers). By combining the equations of the papers I come to a complete set, which can be simulated with a numerical method.

I expect to revisit more reactions on my website. With the new camera I can make new recordings and some reactions may exhibit interesting phenomena when recorded at high speed.

wg48 - 25-10-2018 at 04:27

Quote: Originally posted by woelen

In the derivation of the equations I used a few papers. They can be downloaded on my webpage (there are links in the webpage to these papers). By combining the equations of the papers I come to a complete set, which can be simulated with a numerical method.

I expect to revisit more reactions on my website. With the new camera I can make new recordings and some reactions may exhibit interesting phenomena when recorded at high speed.

FANTASTIC job on the experiments, simulation and web report woelen.