NaHCO3 and CaCl2 reaction
watte...@my-dejanews.com
bicarbonate and calcium chloride? Thanks.
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Uncle Al
watte...@my-dejanews.com wrote:
>
> Can somebody give me an equation that explains the reaction between sodium
> bicarbonate and calcium chloride? Thanks.
The spectator ions (which is all of them) do a square dance.
Equilbirium is pushed by precipitation of solids or evolution of gases
given solubility and equilibrium constants.
--
Uncle Al Schwartz
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Darrin Dailey
<7g76ka$5u4$1...@nnrp1.dejanews.com>...
:Can somebody give me an equation that explains the reaction between sodium
:bicarbonate and calcium chloride? Thanks.
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:http://www.dejanews.com/ Search, Read, Discuss, or Start Your Own
Looks like NaHCO3 (aq) + CaCl2 (aq) -----> NaCl (aq) + HCl (aq) + CaCO3 (s)
That's an expensive way to make chalk, although precipitated calcium
carbonate has certain coatings applications if you can control the particle
size.
Uncle Al
It's more interesting than that. HCl and CaCO3 do not coexist.
You've got multiple equlibria here, and the mechanism for karst
topograpy and pipe scale.
Peter Pommerenk
> Can somebody give me an equation that explains the reaction between sodium
> bicarbonate and calcium chloride? Thanks.
>
temperature?
pressure?
solvent, e.g. water, methanol, orange juice?
In water at standard temperature and pressure both salts dissociate
almost completely, yielding Na(+), Cl(-), HCO3(-), Ca(2+). One equation
will not be sufficient to describe what's gonna happen. Depending on
the final pH you may find Ca(2+), CaOH(+), CaHCO3(+), CaCO3(aq), CO2(aq),
H2CO3(aq), HCO3(-), CO3(2-), Na(+), NaHCO3(aq), NaCO3(-), Cl(-). And if
you add plenty of both you might precipitate some form of CaCO3(s) or
Ca(OH)2(s). When you do it, make sure you don't violate Gibbs phase
rule ;-)
Peter
Oscar Lanzi III
have this reaction:
2HCO3(-) --> CO3(2-) + CO2 + H2O
The reaction as written is not favored, but with calcium ions the
carbonate ions can be precipitated as CaCO3.
So, excluding spectator ions, we have:
Ca(2+) + 2HCO3(-) -->
CaCO3(s) + CO2 + H2O
--OL
Peter Pommerenk
On Sat, 1 May 1999, Oscar Lanzi III wrote:
[snip]
> But you could have this reaction:
>
> 2HCO3(-) --> CO3(2-) + CO2 + H2O
>
> The reaction as written is not favored
Why? Delta G0 = -214 kJ/mol ! Or am I missing something?
> but with calcium ions the carbonate ions can be precipitated as CaCO3.
> So, excluding spectator ions, we have:
>
> Ca(2+) + 2HCO3(-) --> CaCO3(s) + CO2 + H2O
... which would not proceed, if the first reaction did not occur. Or
where does the calcium get the carbonate ions (CO3(2-)) from? The driving
force for the latter reaction is not the precipitation of CaCO3
(Ca(2+) + CO3(2-) <-> CaCO3, Delta G0 = -47 kJ/mol).
I hope this illustrates the problems associated with lumping multiple
equilibria into one reaction.
Peter
Oscar Lanzi III
decomposition is so strongly favored without the Ca(2+) or any
co-reagent, why does baking soda not fizz when added to just water?
(I'm not referrng to what happens when you add an acidic substance).
And when Ca(2+) is added, can it not essentially displace a proton from
CO3(2-), the way fluorine displaces oxygen from water which would
otherwise not decompose?
You talk about multiple equilibria. But the net effect of all the
reactions that this implies must be an overall reaction that
precipitates CaCO3. The one proposed is the only one that makes any
sense (to me, anyway) as the net result of all the interactions.
--OL
Peter Pommerenk
> I strongly disagree with your argument. If the bicarbonate
> decomposition is so strongly favored without the Ca(2+) or any
> co-reagent, why does baking soda not fizz when added to just water?
Because the resulting solution may not be supersaturated with CO2.
Remember carbon dioxide is soluble in water! A water well equilibrated
with the atmosphere contains ~ 10^-5 M CO2(aq) and does it fizz?? Use a
vacuum pump and a suction flask and you will see bubbles (There are other
dissolved gases, too).
> (I'm not referrng to what happens when you add an acidic substance).
> And when Ca(2+) is added, can it not essentially displace a proton from
> CO3(2-), the way fluorine displaces oxygen from water which would
> otherwise not decompose?
Does it? Well, Ca(2+) won't displace a proton from CO3(2-), because there
are none.
> You talk about multiple equilibria. But the net effect of all the
> reactions that this implies must be an overall reaction that
> precipitates CaCO3.
No. If you're not supersaturated with respect to CaCO3, precipitation
will not occur. The calcite saturation horizon and the lysocline in the
ocean are at a depth of 2500m and 4000m, respectively. There are 10mM
calcium in surface seawater... think about it.
> The one proposed is the only one that makes any
> sense (to me, anyway) as the net result of all the interactions.
Fine.
au revoir,
Peter
Oscar Lanzi III
of -214 kJ/mol (almost 100 times RT) the reaction would go to such an
extent that the amount of supersaturation would be enough to cause the
mixture to fizz -- the partial pressure of CO2 would be way above
atmospheric.
--OL
chrispy...@gmail.com
> Can somebody give me an equation that explains the reaction between sodium
> bicarbonate and calcium chloride? Thanks.
>
> http://www.dejanews.com/ Search, Read, Discuss, or Start Your Own
Mixing the two solutions together: CaCl2(aq) + NaHCO3(aq)
The result has been difficult to understand.
I had believed that Ca(HCO3)2* + NaCl(aq) was the result. Doesn't this all depend on proportions, temperature, and grade reagents?
*Ca(HCO3)2 Is this correct? It was a solid white precipitate--in different concentrations it was often powdery, or this time: rather curdled/gel-like solid that was clear like a jello at first, but soon became less viscous and looked like Elmer's White Glue after about 5 or 10 more minutes.)
What's going on with these strange changes?
Peter Fairbrother
> On Wednesday, April 28, 1999 3:00:00 AM UTC-4,
> watte...@my-dejanews.com wrote:
>> Can somebody give me an equation that explains the reaction between
>> sodium bicarbonate and calcium chloride? Thanks.
>>
>> -----------== Posted via Deja News, The Discussion Network
>> ==---------- http://www.dejanews.com/ Search, Read, Discuss,
>> or Start Your Own
>
> Interesting discussion. We just mixed 20g CaCl2 (Driveway Heat) into
> about 100ml water. We also mixed 20g of NaHCO3 (washing soda?)
In practice products sold as washing soda may contain less than 60%
sodium carbonate.
Sodium bicarbonate NaHCO3 is readily available in fairly pure form as -
sodium bicarbonate, or bicarbonate of soda, used in cookery.
I will assume you are using sodium carbonate.
into
> another 100ml. Both grew quite hot and cloudy during the ionization
> process(audibly fizzing--but not visibly) for several minutes of
> occasional stirring.
>
> Mixing the two solutions together: CaCl2(aq) + NaHCO3(aq)
>
> The result has been difficult to understand.
>
> I had believed that Ca(HCO3)2* + NaCl(aq) was the result. Doesn't
> this all depend on proportions, temperature, and grade reagents?
> *Ca(HCO3)2 Is this correct?
yes, or CaCO3 if you used washing soda.
It was a solid white precipitate--in
> different concentrations it was often powdery, or this time: rather
> curdled/gel-like solid that was clear like a jello at first, but soon
> became less viscous and looked like Elmer's White Glue after about 5
> or 10 more minutes.)
>
> What's going on with these strange changes?
in very fine form, which remains as a suspension which doesn't seperate
easily.
If there is excess chloride then some CaCl(HCO3) particles can form too,
which will tend to "stick" to the water more than calcium carbonate
would and so be harder and slower to seperate.
Impurities in the calcium chloride and soda can also cause the
suspension to be slow to seperate.
Sometimes suspensions never seperate.
-- Peter Fairbrother
Peter Fairbrother
>> I had believed that Ca(HCO3)2* + NaCl(aq) was the result. Doesn't
>> this all depend on proportions, temperature, and grade reagents?
>
> yes - eg products can include solid CaCl(HCO3) as well.
concentrations which aren't whole numbers. like
Ca(HCO3(1.277)Cl(0.723).6.25 H2O. I don't know whether that particular
one actually can exist, but you get the idea.
There are several crystal structures which can form, and the solid
structures which form aren't always crystalline or regular - sometimes
where you would expect one ion it will be replaced by another.
Also, solid Ca(HCO3)2 does not exist in bulk, it tends to
disproportionate into calcium carbonate and carbon dioxide. I believe
hydrated forms can be precipitated as suspensions though.
It is also fairly soluble in water, and may not precipitate if the
initial solutions are dilute enough and the pH is high.
In general what tends to precipitate is calcium carbonate, along with a
lot of other stuff. Carbon dioxide can be given off too.
It can get verra complicated, Captain.
Hmm, just repeated your experiment.
Couldn't dissolve 20g sodium bicarbonate in 100 ml water, solubility is
only 9.1 g/l. Dissolved 20g in 250 ml water instead. It did not get hot
during dissolving.
20g CaCl2 dissolved in 100ml water just fine, but did not get hot.
Mixing the two, they got a little warm but not hot, went cloudy, and
gave of CO2 gas. Five minutes later, still giving off gas, now like
milk. Ten minutes later, still a little gas being given off, precipitate
settling.
I'd guess the main reaction is something like:
CaCl2 + 2NaHCO3 -> CaCO3 + 2 NaCl + CO2 +H2O
Also tried 20g/100ml calcium chloride with 23g/100ml potassium carbonate
(no sodium carbonate available). Both dissolved OK.
When mixed went like thick yoghurt, then like elmer's, then precipitate
separated within five minutes. No gas given off. Did not get hot.
The reaction will be something like:
CaCL2 + K2CO3 -> CaCO3 + 2 KCl
So I guess you used washing soda, not sodium bicarbonate. Washing soda
is notoriously impure, pretty much anything could have happened...
-- Peter Fairbrother
chrispy...@gmail.com
> Can somebody give me an equation that explains the reaction between sodium
> bicarbonate and calcium chloride? Thanks.
>
> -----------== Posted via Deja News, The Discussion Network ==----------
> http://www.dejanews.com/ Search, Read, Discuss, or Start Your Own
I was using some old pool chemical PH UP, labeled sodium carbonate! We can only assume that it was mOstly NaCO3 ... I does get hot when dissolving?
Why does dissolving these solids release so much heat?
This Calcium Chloride fascinates me to no end. I put together a fairly saturated solution and boiled it down to half the volume. It started to stick to the sides quite a bit, so I decided to stop heating while most of it was still fairly translucent. Seemed pretty viscous, like melted/boiing sugar. Suddenly it went opaque, as if all of it shifted from one structure to another.
I have read that Calcium Chloride hydrate has several forms, 2H2O to 10H2O. Might the sudden change in appearance have been a shift from one hydrate to another? I had the plate set to 500C although I doubt I ever got it higher than 175C... Thanks again.
Peter Fairbrother
> On Wednesday, April 28, 1999 3:00:00 AM UTC-4,
> watte...@my-dejanews.com wrote:
>> Can somebody give me an equation that explains the reaction between
>> sodium bicarbonate and calcium chloride? Thanks.
>>
>> -----------== Posted via Deja News, The Discussion Network
>> ==---------- http://www.dejanews.com/ Search, Read, Discuss,
>> or Start Your Own
>
> Thanks a million for the help. I was glad to learn (relearn!) that
> the crystals ARE hydrated, like sodium acetate crystals, but totally
> surprised to learn that the compound is even ionized in the crystal
> solid! Wow.
high-melting point, as the ions are tightly attracted to each other.
Other solids are held together by different means. like Van der Walls
forces, and tend to melt at lower temperatures.
> I was using some old pool chemical PH UP, labeled sodium carbonate!
> We can only assume that it was mOstly NaCO3 ... I does get hot when
> dissolving?
>
> Why does dissolving these solids release so much heat?
off heat. Something like Ca++ + H2O -> Ca.H2O++, though maybe not
exactly that.
It isn't just dissolving them which gives off heat - going from the
anhydrous solid to the solid hydrates usually gives off a lot more heat
than dissolving them, especially going from anhydrous to the lowest
hydrate.
For CaCl2 the lowest hydrate is the monohydrate, going from monohydrate
to dihydrate gives off less heat, and so on to the hexahydrate iirc -
can't remember them all, but there are a lot of hydrates.
>
> This Calcium Chloride fascinates me to no end. I put together a
> fairly saturated solution and boiled it down to half the volume. It
> started to stick to the sides quite a bit, so I decided to stop
> heating while most of it was still fairly translucent. Seemed pretty
> viscous, like melted/boiing sugar. Suddenly it went opaque, as if
> all of it shifted from one structure to another.
>
> I have read that Calcium Chloride hydrate has several forms, 2H2O to
> 10H2O. Might the sudden change in appearance have been a shift from
> one hydrate to another?
heating, but I doubt you got to the monohydrate. The dihydrate melts at
176 C ... if the heat of fusion is low then things can solidify very
quickly.
> I had the plate set to 500C although I doubt
> I ever got it higher than 175C... Thanks again.
wikipedia says.
http://en.wikipedia.org/wiki/Calcium_chloride gives the hydrates and
melting points.
and page 2 of www.cal-chlor.com/PDF/GUIDE-physical-properties.pdf
gives the heats of solution.
This may also be interesting:
http://en.wikipedia.org/wiki/Hydration_energy
-- Peter F
David Bostwick
news:e51a60d5-ebcd-4ea2...@googlegroups.com:
Soviet_Mario
> Bad form to reply to on's own posts, but:
>
>>> I had believed that Ca(HCO3)2* + NaCl(aq) was the
>>> result. Doesn't
>>> this all depend on proportions, temperature, and grade
>>> reagents?
>>
>> yes - eg products can include solid CaCl(HCO3) as well.
>
> That may be too simplified - the precipitated solids can
> have mixed concentrations which aren't whole numbers. like
> Ca(HCO3(1.277)Cl(0.723).6.25 H2O. I don't know whether that
> particular one actually can exist, but you get the idea.
compatible (dimensionally and structurally) as to mix freely
in non stoichiometric lattice ?
It is true in some cases (i.g. where coordination numbers
and radii ratios are alike), but it doesn't seem a general rule.
Chloride and hydrogen carbonate doesn't seem that similar.
----
does Ca(HCO3)Cl exist in the solid state ? Im not sure. It
is surely very soluble, indeed.
----
Ca(HCO3)2 is metastable (it resist in the cold, better
under some CO2 pressure, and not too conc. soln.).
If the mixed sol. were just warm, or better hot, surely it
disproportionates to CaCO3 and CO2/H2O for most part.
It begins to do it quickly over 60-70°-
This equilibrium is fairly sensible to minimal variation of
conditions (pH, CO2 pressure, T, common ions etc)
>
> There are several crystal structures which can form, and the
> solid structures which form aren't always crystalline or
> regular - sometimes where you would expect one ion it will
> be replaced by another.
>
>
>
> Also, solid Ca(HCO3)2 does not exist in bulk, it tends to
> disproportionate into calcium carbonate and carbon dioxide.
> I believe hydrated forms can be precipitated as suspensions
> though.
Imho it disproportionate as above, in ordinary conditions
(but perhaps ... under some tens or hundreds of CO2 partial
pressure, who knows)
>
> It is also fairly soluble in water, and may not precipitate
> if the initial solutions are dilute enough and the pH is high.
>
> In general what tends to precipitate is calcium carbonate,
> along with a lot of other stuff. Carbon dioxide can be given
> off too.
>
>
> It can get verra complicated, Captain.
>
>
>
>
> Hmm, just repeated your experiment.
>
> Couldn't dissolve 20g sodium bicarbonate in 100 ml water,
> solubility is only 9.1 g/l. Dissolved 20g in 250 ml water
> instead. It did not get hot during dissolving.
>
> 20g CaCl2 dissolved in 100ml water just fine, but did not
> get hot.
>
> Mixing the two, they got a little warm but not hot, went
> cloudy, and gave of CO2 gas. Five minutes later, still
> giving off gas, now like milk. Ten minutes later, still a
> little gas being given off, precipitate settling.
>
> I'd guess the main reaction is something like:
>
> CaCl2 + 2NaHCO3 -> CaCO3 + 2 NaCl + CO2 +H2O
>
>
>
>
>
> Also tried 20g/100ml calcium chloride with 23g/100ml
> potassium carbonate (no sodium carbonate available). Both
> dissolved OK.
>
> When mixed went like thick yoghurt, then like elmer's, then
> precipitate separated within five minutes. No gas given off.
> Did not get hot.
>
> The reaction will be something like:
>
> CaCL2 + K2CO3 -> CaCO3 + 2 KCl
>
>
>
> So I guess you used washing soda, not sodium bicarbonate.
> Washing soda is notoriously impure, pretty much anything
> could have happened...
>
>
>
> -- Peter Fairbrother
>
1) Resistere, resistere, resistere.
2) Se tutti pagano le tasse, le tasse le pagano tutti
Soviet_Mario - (aka Gatto_Vizzato)