MrPhlopper
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Calcium Hydroxide and Hydrochloric Acid Reaction
For fun, I was trying to make a small quantity of Calcium Chloride according to this reaction:
Ca(OH)2 + 2HCl (aq) -> 2H20 + CaCl2
I added 4mL of 20% HCl onto 1 gram of dry Ca(OH)2. Unexpectedly, a small amount of a whiteish pungent smelling gas evolved from the reaction beaker.
I terminated the reaction due to this unknown foul smelling gas. On my next attempt, I added an excess of water to make a solution of Ca(OH)2 before
adding my 20% HCl, and everything went as planned with no gas evolution. I titrated to pH 7 and then partially evaporated it to let CaCl2 crystals
form at the bottom.
My question is what was this gas and why did it form? My best guess is that it was HCl gas, though it’s possible it could’ve been Cl2 or H2 given
the atoms involved in the reaction. The only explanation that I could come up with was that the dry Ca(OH)2 was more hydrophilic and pulled some of
the water from the hydrochloric acid before the reaction could even proceed. The anhydrous HCl then went up as gas. I also know that HCl gas can
also be produced by dropping hydrochloric acid on dry calcium chloride. Perhaps there is a similar reaction with dry calcium hydroxide? Or is it
possible that calcium chloride, formed in the reaction, desiccated the HCl? I would’ve thought that the excess water in the nonconcentrated HCl
solution would’ve been enough to prevent gas formation but perhaps I needed to use a lower concentration of HCl.
What do you guys think happened?
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DraconicAcid
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I suspect the heat evolved from the reaction would have boiled off some HCl.
Please remember: "Filtrate" is not a verb.
Write up your lab reports the way your instructor wants them, not the way your ex-instructor wants them.
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MrHomeScientist
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It's also possible that dropping the liquid on dry powder kicked some dust up, which then made it to your nose. Often when opening my bottle of NaOH,
I get a whiff of something acrid - that's hydroxide dust that was disturbed when I scooped out what I needed.
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blogfast25
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Adding a fairly concentrated, strong [in the Bronsted sense] acid like 20 % HCl to a solid strong alkali like Ca(OH)2 is nutty beyond
description.
The Enthalpy of Neutralisation:
H3O+ + OH- === > 2 H2O
of - 57 kJ/mol (of acid) is enough to cause your acid to start boiling immediately, possibly almost explosively ('steam explosion').
And there's more enthalpy to be released too: with so little water present, the formed CaCl2 has no choice but to crystallise immediately, releasing
more energy (heat). In the case of CaCl2 that may not have happened because it is highly soluble but there are always pockets of high concentration,
where mixing is inefficient.
Anyone who wants to do neutralisation reactions using concentrated solutions should at a minimum estimate the end temperature of the mixture to avoid
nasty surprises, like over-boiling and severe splattering. Trust me, I do this EVERY time when in doubt regarding the end temperature being possibly
too high.
What you smelled was hydrochloric acid vapours. Very pungent and not great for your lungs either.
This bit:
Quote: Originally posted by MrPhlopper  | | The only explanation that I could come up with was that the dry Ca(OH)2 was more hydrophilic and pulled some of the water from the hydrochloric acid
before the reaction could even proceed. The anhydrous HCl then went up as gas. |
... is simply nonsense.
[Edited on 23-5-2013 by blogfast25]
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woelen
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@blogfast25: Please be not so harsh. This action might be not the smartest thing, but MrPhlopper did this reaction on a small scale with just a single
gram, so nothing serious happened. I agree that mixing strong bases with concentrated strong acids is not the wisest thing to do, but in such small
amounts the effects will remain manageable. I'm quite sure that in your young and uneducated years you also sometimes did less wise things  I certainly did 
@MrPhlopper: I indeed think that you smelled HCl-gas, whch escaped from the acid due to high temperatures. This gas is quite toxic, but the single
whiff you obtained is too little to worry about. It is good that you terminated the experiment after you perceived the bad gas.
You can also learn something from this. Experimenting with chemicals is something which must be done with care and it is good to think about what
might happen. Using small quantities also is a very good precaution. I also sometimes mix chemicals, even when I expect an extremely violent reaction,
but if I do so, then I never use more than a few 100's of milligrams. At the moment I work with experiments on ClO2 (a highly explosive and quite
toxic gas), but I only do these experiments with very small amounts. I did have a few explosions already, but these never are worse than loud WHOOSH
sounds, due to the small scale of the experiments.
Another very important thing is that you never look into a test tube from above when a violent reaction may occur. If the content erupts into your
eyes, then the experience can be disastrous. So, always point the open end of a test tube away from you.
Keep this warning in mind: The goddess of chemistry is a beautiful one, but you have to treat her with care. She can bite seriously.
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blogfast25
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| Quote: Originally posted by woelen | | @blogfast25: Please be not so harsh. This action might be not the smartest thing, but MrPhlopper did this reaction on a small scale with just a single
gram, so nothing serious happened. |
And choosing that small scale was wise. Trouble is that some wouldn't have taken that precaution and could have suffered the consequences. Mixing
Ca(OH)2 with 20 % HCl is like pouring it over KOH.
Runaways and unexpected exotherms are a common cause of accidents or at a very minimum experiments ruined. Not so long ago very experienced member
'Magpie' suffered a runaway due to an unaccounted exotherm when oxidising propylene glycol with nitric acid to malonic acid. There were no
consequences but it could have ended very badly. Recently I suffered a small fire when very foolishly and absentmindedly treating potassium with
methylated spirits. With unexpected heat one has to be careful, especially around potentially corrosive/toxic chemicals.
Maybe I was a bit too harsh but it's a labour of love, really...
[Edited on 23-5-2013 by blogfast25]
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amazingchemistry
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When estimating exothermicity, would it be as simple as looking up the enthalpy values of the reagents and of the expected products? What if we don't
really know what the "expected" products should be?
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Fantasma4500
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when you say foul i think H2S smell.. unless if there was impurities (in which could easily form H2S if there was sulfides present) then i dont see
reason for H2S to form..
what you can try some other time is to put some ammonia next to the whole thing as described, in which you smelt something.. if its the HCl it will
make a nice thick grey smoke of NH4Cl
i believe that there isnt much danger to adding an acid such as HCl (not even being 30 or 37% in this case) to a very insoluble hydroxide such as
Ca(OH)2
now if we took 25% NH4OH and added 98% H2SO4 it would get pretty intense, but thats much more concentrated than HCl can get as store bought (:
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blogfast25
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| Quote: Originally posted by amazingchemistry | | When estimating exothermicity, would it be as simple as looking up the enthalpy values of the reagents and of the expected products? What if we don't
really know what the "expected" products should be? |
You need to really know what is reacting with what and what is being formed (and in what state: solid, liquid, gaseous or in solution) to be able to
predict the change in enthalpy ΔH (heat released or absorbed during reaction) and in many cases that's not easy to do. 'On paper', knowing the
Enthalpies of Formation of all reagents and reaction products allows you to calculate ΔH using Hess Law, as a simple addition/substraction. In
practice these values aren't always known and several complications can arise from changes in state (melting, evaporating, solvation, crystallisation,
etc). Often the value of ΔH is determined experimentally, using calorimetry, instead of from theory.
In the case of aqueous neutralisation reactions, one generally assumes only H3O+ + OH- === > 2 H2O proceeds and that most of the other ions present
are just 'spectators' that don't influence ΔH. The calculation then becomes very simple and the end-temperature determined from ΔH = m
Cp ΔT (*), assuming no heat losses or external heat sources.
(*) ΔH = m Cp ΔT: m mass of water present, Cp specific heat capacity of water, ΔT change in temperature of the
solution.
In short, sometimes it's quite straightforward, sometimes it's rather challenging...
[Edited on 24-5-2013 by blogfast25]
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MrPhlopper
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Thanks for your replies everyone! The HCl gas as a result of the highly exothermic neutralization reaction seems plausible and consistent with what I
observed.
I'm an inexperienced chemist and I agree this wasn't my best idea - I didn't consider the heat that would be released by the reaction. But
fortunately nothing bad happened and I'll know better next time.
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chloric1
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You can use the heat of neutralization to your advantage. For example, strong ammonium hydroxide neutralize with strong HCl can make a strong ammonium
chloride solution near its boiling point from room temperature acid and base. This would be quite handy for preparing slightly soluble ammonium salts
such as the dichromate, vanadate, or perchlorate via saturated solutions of the sodium salts. Or if you want to reduce time to crystallize a salt.
For example, neutralizing maybe a 40% sulfuric acid solution with strong solution of NaOH then putting in ice bath to crystallize the decahydrate.
[Edited on 7/2/2015 by chloric1]
Fellow molecular manipulator
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Dangle89
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Hello MrPhlopper
I had a possibly similar experience making a small amount of MgCl2 from Mg and HCl.
Got distracted by the postman and didn't realize I hadn't diluted my 37% HCL 
Had about 5g of Mg in a test tube and only added maybe 1/4-1/2ml of HCl before similar foul smelling white smoke started pouring out!
Continence of the test tube never came close to coming out of the top but made me crap my pants for half a second before I realized what I had done!
Ran away and waited for the smoke to clear and repeated the experiment with sufficiently diluted acid
Glad you are OK and be safe :-)
[Edited on 29-7-2015 by Dangle89]
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pneumatician
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| Quote: Originally posted by woelen | @blogfast25:
Using small quantities also is a very good precaution. I also sometimes mix chemicals, even when I expect an extremely violent reaction, but if I do
so, then I never use more than a few 100's of milligrams. At the moment I work with experiments on ClO2 (a highly explosive and quite toxic gas)
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wow, check the MMS people, most totally ignorant of what they have at hand and I only read some minor domestic "issues" of 1 or 2 persons. so what you
are doing with ClO2 my friend?
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Zyklon-A
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pneumatician lolwut? What are you doing on a computer?
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Dangle89
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| Quote: Originally posted by pneumatician |
wow, check the MMS people, most totally ignorant of what they have at hand and I only read some minor domestic "issues" of 1 or 2 persons. so what you
are doing with ClO2 my friend? |
Sorry this might be a stupid question but what is the MMS? Is that another country's version of the MSDS?
Googled it and only got the long name for picture messaging 
Cheers
Dangle
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Bot0nist
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Miracle Mineral Supplement / or Solution.
https://en.m.wikipedia.org/wiki/Miracle_Mineral_Supplement
A particularly dangerous snake oil remedy. A solution of sodium chlorite. It is activated by addition of a food grade acid and taken internaly by
unwitting fools or victims of fools.
"MMS is falsely promoted as a cure for HIV,malaria, hepatitis viruses, the H1N1 flu virus,common colds, autism, acne, cancer, and much more."
[Edited on 4-8-2015 by Bot0nist]
U.T.F.S.E. and learn the joys of autodidacticism!
Don't judge each day only by the harvest you reap, but also by the seeds you sow.
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kecskesajt
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My explanation is also:neatrulisation or acid-base reactions are quite exothermic.Calcium oxide hydrolysis/dissolution is exothermic(can reach water
to boil)
The reaction between sodium hydroxide solution (20%) and sulphuric acid (50%)is so exotherm that this is splattering the solution everywhere.
Back tho your question.Use Ca(OH)2 suspension and add the HCl INTO the Ca(OH)2.
The smell is HCl vapour and H2S.Commercial grade lime almost always have sulphide/polysulphide impurities.
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pneumatician
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and you, what are doing with your soul?
lolwut?
http://theactionelite.com/site/wp-content/uploads/2013/11/lo...
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pneumatician
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| Quote: Originally posted by Bot0nist | Miracle Mineral Supplement / or Solution.
https://en.m.wikipedia.org/wiki/Miracle_Mineral_Supplement
A particularly dangerous snake oil remedy. A solution of sodium chlorite. It is activated by addition of a food grade acid and taken internaly by
unwitting fools or victims of fools.
"MMS is falsely promoted as a cure for HIV,malaria, hepatitis viruses, the H1N1 flu virus,common colds, autism, acne, cancer, and much more."
[Edited on 4-8-2015 by Bot0nist] |
try to edit this article in wikishit, in a moment is again the official shit.
http://g2cforum.org/
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deltaH
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Commercial lime contains significant amounts of calcium carbonate, particularly if it's old and has been stored poorly (reacting with the CO2 from the
atmosphere).
I strongly suspect what you observed was simple the vigorous production of CO2 from the acidification of calcium carbonate impurity which carried with
it HCl aerosols... very nasty on the nose (and everything else) 
If so, the reaction is simply:
CaCO3(s) + 2HCl(aq) => CaCl2(aq) + CO2(g) + H2O(l)
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