Where to look up vaporization points of metal salts?

[billysundays]

A google search for the vaporization points of metal compounds, such as strontium carbonate, or just pure strontium, doesn't turn up any relevant info. Neither does looking them up on Wikipedia. Anybody know where I can find this info?

[OldMarine]

Carbonates don't vaporize, they just decompose. Try going to Sigma Aldrich and looking up the MSDS on the chems you're wondering about. Their sheets are usually more detailed than most.

[billysundays]

Carbonates don't vaporize, they just decompose. Try going to Sigma Aldrich and looking up the MSDS on the chems you're wondering about. Their sheets are usually more detailed than most.

I'll try that, thanks. So if carbonates don't vaporize, how do they contribute metal in gaseous form to the flame to effect its color?

[OldMarine]

Decomposition can contribute all of the components as well as vaporization would. The decomposition temps are listed in most MSDS sheets.

I don't claim to know shit about chemistry but I have to read the sheets often.

[greenlight]

Oldmarine is right, carbonates like strontium carbonate decompose upon melting to give the corresponding metal oxide (SrO) and CO2.

[billysundays]

I see, thanks for further clarifying guys. As you can tell, I only have a rudimentary understanding of chemistry, but my assumption has been that with all the chemical compounds used as colorants, when heated to a high enough temperature, the metal disassociates with the "salt", be it carbonate, oxide, chloride, what have you, at which point it becomes a vapor/gas, and in the case of strontium, copper, and barium, recombines with any free chlorine gas to create a (still-gaseous) monochloride salt.

 

It sounds like what you guys are saying is that in the case of carbonates, there's an extra step whereby it first decomposes to the solid oxide form, and from there disassociates from the oxide and becomes gaseous. Would that be correct?

 

The other question this brings up for me is, why would the metals become gaseous either way? Don't metals melt down to liquid form when heated high enough. Seems like it should take much higher temperatures to vaporize metal than what pyrotechnic compositions could provide, let alone a campfire. My guess is that being in "salt" form lowers their vaporization point, is that right?

Edited September 15, 2017 by billysundays

[greenlight]

It depends on the anion (negative charged ion). Strontium nitrate decomposes to the metal oxide too but also liberates oxygen which will react with the fuel to sustain burning making it an oxidizer and colourant and therefore different from just the carbonate colorant. Heres the reaction:

 

Sr(NO3)2 ---> SrO + NO + NO2 + O2

 

At higher temps than melting it just produces SrO, N2 and O2.

 

Once enough heat is generated to melt and decompose the oxidizer, it will mix intimately in liquid form with the fuel and then the composition will ignite. The ignition starts in a small part of the mix and spreads through the entire portion via the hot gasses and flame.

 

Pure metal melts down when heated at high temperatures yes. When finely powdered, there is alot more surface area to react with the oxidizer. Most metal-fueled comps ignite around the melting point of the metal used. Melting point of Mg is around 600.C which is around the ignition temp of flash using it as the fuel.

 

Metal salts have totally different properties than the metals they come from. Pure lead is a grey hard metal insoluble in water. Lead nitrate is white crushable water soluble salt. Nothing like the initial pure metal.

The difference between the properties of the salt and pure metal form varies. Lead nitrate has a higher melting point by about 150C than pure lead wheras other salts are lower than the metal it was formed from. Alot of things in chemistry trip you up like this I have noticed. You think ypu have found a trend and then bam it changes.

Edited September 15, 2017 by greenlight

[billysundays]

Love the info greenlight, thanks. So how is the process different with other salts,for comparison, like strontium chloride? Does it vaporize straight to strontium and chlorine, to recombine as as monochloride? What exactly are the practical implications for the differences in how carbonates and chlorides vaporize? Edited September 15, 2017 by billysundays

[greenlight]

Strontium chloride once again forms Strontium oxide because it reacts with the atmosphere like the others and also chlorine. No oxygen so it will not function as an oxidizer like the nitrate form.

 

It would be a better colourant than the carbonate because of the chlorine which enhances colour purity but it is highly hygroscopic so the carbonate is usually used instead for colouring purposes.

 

A lot of salts just can't be used because of hygroscopicity issues even though they would function well otherwise.

Potassium is a very popular cation as you would notice because of its very low hygroscopicity. Li, Mg salts mostly no good and sodium is hygroscopic and colours flames yellow so good luck making other colours. Makes great illuminating flares though.

[OldMarine]

Just had a recollection. Drill ¼" holes in a 3" log. Make a slurry of your chosen color producing chem and press with the fingers into the holes. Wrap in plastic for a day and place in the sun to force the chem into the wood. Remove the plastic and dry for a week in the shade with moving air.

Don't dry them too long or they burn way too fast no matter whether you use nitrates, carbonates or oxides.

Sorry I didn't remember sooner but drugs back in the day.....

[billysundays]

Lol, OldMarine, sounds like you've always had a colorful time, with or without pyrotechnics. So would the holes be drilled all the way thru? What colorants did you add to the logs when you tried this or was this an idea you've read about?

 

Greenlight, it seems like what your saying is that the same process is true for all the colorants: the metal basically trades its anion for oxygen to becomes an oxide, then disassociates from the oxide to become gaseous, then combines with any available chlorine gas to become a gaseous monochloride. Is that right? I think I misunderstood OldMarine when he said strontium carbonate decomposes instead of vaporizes to mean "unlike what's common with most other metal salts".

 

Well anyway, the reason I was asking about vaporization points to begin with was because I had read in a thread here that barium chloride was useless for pyrotechnics because "it vaporizes at 2,840 degrees F, which is way hotter than you can achieve with a pyrotechnic mixture", and I was just about to add barium chloride to my purchase, so I thought "if pyrotechnic compositions don't get hot enough to get it to work, a bonfire sure as hell won't either, maybe I should ask about the vaporization points for the compounds I'm considering, to try and avoid wasting money on duds". Now I'm wondering why Firefox would even carry barium chloride if its a dud.

 

So yeah, got lost in the weeds here a bit, but can't complain. This is all intriguing. Thanks for letting me pick your brain Greenlight.

Edited September 16, 2017 by billysundays

[greenlight]

Most of the common colorants decompose to the corresponding oxide and monochloride is ususally responsible for the colour. Once in the gaseous state the reactions products can vary depending on the formula and color being emitted.

Chlorine can be evolved from the colorants but mainly comes from the oxidizers.

It starts getting more technical and confusing with all the factors when you go further into the colour production. You have to get the electrons excited and hot but not too hot or the effect is destroyed like the infamous perfect blue. High temperatures will also decompose the monochloride further and lose the color.

 

Barium chloride is not very popular in pyrotechnic compositions due to hygroscopicity issues again I think.

The decomp temp isn't too bad it melts at around 900 I just looked it up but I haven't seen any green comps with it. Your better off buying barium carbonate or barium nitrate for green colour comps preferably barium nitrate I personally think.

Edited September 16, 2017 by greenlight

[billysundays]

I looked again at the post I quoted before about the vaporization point of barium chloride(here), and I'm realizing that he referencing the boiling point to indicate vaporization point, which makes sense; its true of water after all, and if you look up the boiling points of many compounds in wiki, such as barium sulfate or cuprous chloride, you'll see it even specifies that they decompose at this temp. You seem to believe melting point is the relevant property. Could you explain?

 

Interestingly enough, copper(II) carbonate hydroxide seems all but ready to decompose, melting at 200 °C, and boiling at 290 °C! But if what ultimately matters is the melting or boiling point of the oxides of these metals, as you've described, then:

 

Copper(I) oxide
Melting point 1,232 °C
Boiling point 1,800 °C

Copper(II) oxide
Melting point 1,326 °C
Boiling point 2,000 °C

 

That are much higher temps that what people say pyrotechnics can reach, and would indicate that there should be no way that they work to color the flames of pyrotechnics, let alone a humble campfire, neither of which is true, so there's obviously more to the picture than what we've discussed so far. What are your thoughts. Maybe Mumbles can weigh in as well?

Edited September 16, 2017 by billysundays

[greenlight]

Decomposition starts at the melting point of the oxidizer/colourant and there is a difference between the amount of heat needed to decompose the pure chemical compared to when it is in a pyrotechnic mixture.

 

Once the liquid phase from melting is reached the next step is combustion of the composition which will be a large temperature spike and the full decomposition will be witnessed. The liquid phase allows the fuel and oxidizer to mix intimately and soonafter ignite.

 

If you are trying to decompose the chemical in its pure form over a flame, a higher temp would be needed.

With an actual pyrotechnical composition, once ignition is achieved, there is more than enough heat to reduce the chemicals to their decomposition products.

 

I worded it confusingly before but melting is the start of the whole decomposition process. Once both the fuel and oxidizer are in the liquid (melted) state, the compositikn will usually ignite.

Edited September 16, 2017 by greenlight

[billysundays]

Jinx! Seems like I posted my edit at the same moment you posted your reply. Could you take another look at my last post? Or, wait, did you read it and type out a response in less than 20 seconds?? Cuz that would be impressive!

Edited September 16, 2017 by billysundays

[greenlight]

Hahaha I try explain it thoroughly here.

Once ignited the temperature can get much hotter than the melting point and then the vaporization stage takes place on its own. But to start off this reaction, a temperature not much higher than the melting point of fuel/oxidizer is needed.

The temperature needed is where the oxidizer is hot enough to release its oxygen and the fuel is in a state where it can bind with the oxygen. After ignition, the temperature will spike rapidly decomposing the rest of the chemical ingredients into their corresponding products.

This is my understanding of the process.

 

Have a look at this chart I have copied from a military pyrotechnics manual listing the ignition temperature of magnesium flash with different oxidizers:

 

NaNO3/Mg -------- 635C ignite temp

Ba(NO3)2/Mg -------- 615C ignite temp

Sr(NO3)2/Mg -------- 610C ignite temp

KNO3/Mg -------- 650C ignite temp

KClO4/Mg -------- 715C ignite temp

 

Magnesium's melting point is 650C. Notice how all the mixtures have an ignition temperature around this melting point temperature. The oxidizer has melted and given it oxygen to the partly liquid state Mg and then poof, ignition which spreads to the rest of the comp by vaporization from the high burn temp.

 

Here is another chart from another book and you can see once the actual composition has ignited and the burning is propagating on its own, the temperature is high enough to meet the values you listed in your post:

 

Photoflash. ------ 2500 to 3500C max flame temp

Solid rocket fuel. ------ 2000 to 2900C max flame temp

Coloured flame mixs.------ 1200 to 2000C max flame temp

 

The melting plays a huge part in the ignition and is the start of decomp and upon ignition vaporization from the burn temp decomposes the rest.

Hopefully I have cleared it up properly now haha.

Edited September 16, 2017 by greenlight

[billysundays]

Well, you've given me quite a morsel to digest, lol! I'm gonna take some time later to visualized what you've described to try and understand it better.

 

If "Coloured flame mixs" means "fireworks", then that temp range would disprove the other poster's (Peret) assertion that the reason barium chloride wasn't working for the OP's composition was because 1,560 °C is hotter than what fireworks can achieve. Its likely a hygroscopicity issue, as you said.

 

So the last question this leaves me then; If copper oxides and pure copper have a boiling point between 1,800 °C and 2562 °C, it shouldn't be possible to get, for example, blue flames from cupric chloride when introduced to a campfire, but that's not the case. It works in a campfire all on its own, as well with several other colorants, and even pure copper gives off green. How can that be explained?

Edited September 16, 2017 by billysundays

[greenlight]

Coloured flame mixs is quite a broad heading for such a large range of different compositions hahaha I agree but thats whats written down in the table.

Those temperatures can definitely be reached especially if metal fuels are thrown in.

 

All I can bring to the table for the copper compounds is that copper chloride melts at about 500 C and decomposes at about 900C so the red embers in a fire pit should be hot enough to bring out the colour in the flames. Anything over the melting point and you will start noticing it I think because it will still be decomposing but at a slow rate.

 

The copper oxide is quite a puzzling one because it has a very high MP and BP like most other oxides and yet it is used to make blues which require a cooler burn for their color purity.

Have you ever seen a campfire coloured with copper oxide before? Would be interesting to see if you get much color saturation.

Edited September 16, 2017 by greenlight

[billysundays]

I plan on testing a sample of all the compounds I get by pouring some in a small metal tray and placing it in the fire to verify what happens for each by themselves, including the oxide. I'll be post results in my main thread.

 

I think we can consider this thread thoroughly answered and closed. Greenlight, thanks again for taking the time and letting me pick your brain, this was interesting.

 

OldMarine, I'll ask you the log idea question again in my main thread so we can continue the conversation there.

[greenlight]

Thats okay billysundays, I am interested to see the results from the fire test.

[Boophoenix]

I can't recall off the top of my head about copper oxide, but all of the other coppers seemed to work fine including fine copper shavings. I had an abundance of PVC as the chlorine donor. The only problem I experianced was burn rate duration. The fire was fairly large though. We were feeding the fire with a 20 ton excavator and I was tossing pounds of mixtures in.

 

An interesting tidbit after the color expires though introducing more chlorine alone can bring the color back. Inhibiting the release of the chlorine should most likely extend the duration, but I've not dabbled with that any yet.

 

Sorry I can't offer anything other that my simple experiment information though. I would also think a calmer fire would be much more rewarding, I just used what was available at the time my curiousity struck.

[billysundays]

A 20 ton excavator?! That sounds like a interesting night. Tell me more.

 

Yeah, I can say I had the same observation; Just adding more PVC was enough to bring back the colors even when I removed the copper wire, which leads me to some more ideas to test later this weekend.