Why does heat reduce solubility

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Marshall Dudley

Posted: Mon Dec 29, 2003 11:13 pm

Guest

Normally a warm liquid will increase the solubility of almost any
solute.

Why is the lime in hard water the opposite? Is it because the heat
drives off dissolved CO2, increasing the ph which then reduces the
solubility?

Thanks,

Marshall

Marvin Margoshes

Posted: Tue Dec 30, 2003 2:16 pm

Guest

"Marshall Dudley" <mdudley@execonn.com> wrote in message
news:3FF0FB52.FC7EDBB7@execonn.com...

Quote:

water than in cold.

rsm109@york.ac.uk

Posted: Tue Dec 30, 2003 9:07 pm

Guest

Marshall Dudley <mdudley@execonn.com> wrote in message news:<3FF0FB52.FC7EDBB7@execonn.com>...

Quote:

It's because CO2 gas has a greater entropy than CO2 solution, so the
formation of the solution is less favourable at high temperatures. The
same is true of any gas.

Bouillote

Posted: Wed Dec 31, 2003 6:43 am

Guest

rsm109@york.ac.uk wrote:

Quote:

Marshall Dudley <mdudley@execonn.com> wrote in message
news:<3FF0FB52.FC7EDBB7@execonn.com>...
Normally a warm liquid will increase the solubility of almost any
solute.

Why is the lime in hard water the opposite? Is it because the heat
drives off dissolved CO2, increasing the ph which then reduces the
solubility?

Thanks,

Marshall

It's because CO2 gas has a greater entropy than CO2 solution, so the
formation of the solution is less favourable at high temperatures.
The
same is true of any gas.

According to Van't Hoff law it's seems to be a question of enthalpy but not
entropy, no?

Rachmanonov

Posted: Wed Dec 31, 2003 8:39 am

Guest

This reaction in question is a chemical reaction. If this were a simple
solubility change, then the condition would be dictated by the change in
combinatorial entropy with temperature and any enthalpic changes that would
arise during the temperature change. For example, increasing the temperature
would increase the combinatorial entropy of the system. However, elevated
temperatures can disrupt solvent-solute interactions (hydrogen bonding, ionic
pairing etc) and induce precipatation.This is not an example of this situation.

This is the decompostion of calcium carbonate to calcium oxide and carbon
dioxide. Increased temperatures are going to increase the rate of the reaction
by increasing the population of reactive molecules(think boltzman distribution)
This is further complicated by the release of CO2 (Henrys law) and the
precipation of Calcium oxide which drives the system away from its equilibrium
position. The non-equilibrium condition of the system renders the Vant hoff
equation invalid, since it describes the effect of temperature on equilibrium
constant of the system.

polychem

rsm109@york.ac.uk

Posted: Wed Dec 31, 2003 9:12 am

Guest

"Bouillote" <superbouillote@yahoo.fr> wrote in message news:<3ff2b611$0$20142$a3f2974a@nnrp1.numericable.fr>...

Quote:

rsm109@york.ac.uk wrote:
Marshall Dudley <mdudley@execonn.com> wrote in message
news:<3FF0FB52.FC7EDBB7@execonn.com>...
Normally a warm liquid will increase the solubility of almost any
solute.

Why is the lime in hard water the opposite? Is it because the heat
drives off dissolved CO2, increasing the ph which then reduces the
solubility?

Thanks,

Marshall

It's because CO2 gas has a greater entropy than CO2 solution, so the
formation of the solution is less favourable at high temperatures.
The
same is true of any gas.

According to Van't Hoff law it's seems to be a question of enthalpy but not
entropy, no?

Ooops. I was thinking of the Gibbs function by itself without
considering the exact relation between it and the equilibrium
constant.

ln k = -deltaG/RT = -deltaH/RT + deltaS/R

though a positive entropy change does mean that the process of CO2
coming out of solution is favourable at the limit of high T. It is
endothermic and involves an entropy increase, so like all such
processes it is favourable at high T and unfavourable at low T. An
endothermic process involving an entropy *decrease* is unfavourable at
any temperature (less so at high T, but still unfavourable).

Apologies for my confusion.

Robert

Mark Tarka

Posted: Wed Dec 31, 2003 9:30 am

Guest

"Bouillote" <superbouillote@yahoo.fr> wrote in message news:<3ff2b611$0$20142$a3f2974a@nnrp1.numericable.fr>...

Quote:

I don't know about that Van't Hoff dude, but, yes.

Mark (Please do not snicker as I say...greetings from the
University of Texas at Austin Smile

Bouillote

Posted: Wed Dec 31, 2003 10:13 am

Guest

Mark Tarka wrote:

Quote:

"Bouillote" <superbouillote@yahoo.fr> wrote in message
news:<3ff2b611$0$20142$a3f2974a@nnrp1.numericable.fr>...
rsm109@york.ac.uk wrote:
Marshall Dudley <mdudley@execonn.com> wrote in message
news:<3FF0FB52.FC7EDBB7@execonn.com>...
Normally a warm liquid will increase the solubility of almost any
solute.

Why is the lime in hard water the opposite? Is it because the
heat
drives off dissolved CO2, increasing the ph which then reduces
the
solubility?

Thanks,

Marshall

It's because CO2 gas has a greater entropy than CO2 solution, so
the
formation of the solution is less favourable at high temperatures.
The
same is true of any gas.

According to Van't Hoff law it's seems to be a question of enthalpy
but not
entropy, no?

I don't know about that Van't Hoff dude, but, yes.

In France we call "Van't Hoff law" this relation

d(lnK°)/dT = DeltarH°/RT²

Bouillote

Posted: Wed Dec 31, 2003 10:18 am

Guest

Quote:

Ooops. I was thinking of the Gibbs function by itself without
considering the exact relation between it and the equilibrium
constant.

ln k = -deltaG/RT = -deltaH/RT + deltaS/R

That's true but you must look at the variation of the constant k so when you
take the "d/dT" the entropy disapears because d(DeltarH)/dT = Cp and
d(DeltarS)/dT= Cp/T

J.Pelmont

Posted: Fri Jan 02, 2004 4:52 am

Guest

Marshall Dudley <mdudley@execonn.com> wrote:

Quote:

Normally a warm liquid will increase the solubility of almost any
solute.
......

Not true for potassium chloride, cesium chloride... Dissolving these
salts in water cools down the mixture. The reversal is true for lithium
chloride. The solubility of sodium salt does not change much with
temperature, maybe with a small increase at elevated temperature. I
believe these differences are due to ion solvatation energies.

--
J. Pelmont
http://perso.wanadoo.fr/pelmont/index.html
for mail remove "aky"

Mark Tarka

Posted: Fri Jan 02, 2004 9:26 am

Guest

"Bouillote" <superbouillote@yahoo.fr> wrote in message news:<3ff2e73b$0$20141$a3f2974a@nnrp1.numericable.fr>...

Quote:

Mark Tarka wrote:
"Bouillote" <superbouillote@yahoo.fr> wrote in message
news:<3ff2b611$0$20142$a3f2974a@nnrp1.numericable.fr>...
rsm109@york.ac.uk wrote:
Marshall Dudley <mdudley@execonn.com> wrote in message
news:<3FF0FB52.FC7EDBB7@execonn.com>...
Normally a warm liquid will increase the solubility of almost any
solute.

Why is the lime in hard water the opposite? Is it because the
heat
drives off dissolved CO2, increasing the ph which then reduces
the
solubility?

Thanks,

Marshall

It's because CO2 gas has a greater entropy than CO2 solution, so
the
formation of the solution is less favourable at high temperatures.
The
same is true of any gas.

According to Van't Hoff law it's seems to be a question of enthalpy
but not
entropy, no?

I don't know about that Van't Hoff dude, but, yes.

In France we call "Van't Hoff law" this relation

d(lnK°)/dT = DeltarH°/RT²

I have been denied access to the books for several
years, but IIRC, the lnK(superscript 0) would
be the equilibrium constant at STP (?), but the
"DeltarH...".

Mark (Treat me like a doormat, hero Smile

Douglas Scot Gillman

Posted: Sun Jan 04, 2004 7:09 am

Guest

ÈIon solvation energiesÈ is a new key phrase to
sci.chem
Dream on nacelles' wind, Caught in tepid Tome, an urge,
Gin in from the cold.

JRP

Posted: Mon Jan 05, 2004 5:40 am

Guest

Douglas Scot Gillman <cstwtch8@aol.comginin> wrote:

Quote:

ÈIon solvation energiesÈ is a new key phrase to
sci.chem
Dream on nacelles' wind, Caught in tepid Tome, an urge,
Gin in from the cold.

Do your home work
(he reads this group and does not know that water molecules are bound to
ions in water...)

--
J. Pelmont
http://perso.wanadoo.fr/pelmont/index.html
for mail remove "aky"

GeneralChemTutor

Posted: Thu Jan 08, 2004 11:01 pm

Joined: 08 Jan 2004 Posts: 41

Remember that CaO is a basic anhydride. It reacts with water to form CaOH; thus the volume of the solid actually increases during the progression of the reaction. Ca and O do not stay solvated by water, Ca further reacts.

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