NaHCO3 and CaCl2 reaction

(too old to reply)

Can somebody give me an equation that explains the reaction between sodium
bicarbonate and calcium chloride? Thanks.
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Interesting discussion. We just mixed 20g CaCl2 (Driveway Heat) into about 100ml water. We also mixed 20g of NaHCO3 (washing soda?) 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? 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?

Washing soda is, in theory at least, sodium carbonate Na2CO3.

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

actually dissolving - the salts are already ionised in the solid state.

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?

yes - eg products can include solid CaCl(HCO3) as well.

*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?

If the concentrations are right the calcium carbonate (chalk) can form
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

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.

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

well, are you sure the two crystal lattices are so mutually
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)

Post by Peter Fairbrother
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 agree

i have doubts :-)
Imho it disproportionate as above, in ordinary conditions
(but perhaps ... under some tens or hundreds of CO2 partial
pressure, who knows)

Post by Peter Fairbrother
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.

why not trying to reason in terms of "mol" instead of grams ?

Post by Peter Fairbrother
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.
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.
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)

Can somebody give me an equation that explains the reaction between sodium
bicarbonate and calcium chloride? Thanks.
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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.

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.

That's common. They are called ionic solids. They tend to be
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?

The ions react with water molecules to form hydrated ions, which gives
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?

You may have gone from eg the tetrahydrate to the dihydrate when
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.

I think it goes from monohydrate to hexahydrate - at least that's what
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

A resurrected thread from 1999? Not much to do today, eh?