Discussion:
Solubility of AgCl
(too old to reply)
Marshall Dudley
2004-12-22 15:19:42 UTC
Permalink
Silver chloride has a solubility of about .89 ppm in cold water.
Accoding to Le Chatelier's Principle the solubility will decrease in the
presence of additional chloride ions. So the addition of HCl or NaCl
will cause the solubility to drop.

However at an excess of chloride of ions a dichloroargentate complex is
formed, which is highly soluble.

Does anyone know at what point the solubility turns around and starts
increasing again with respect to the concentration of Cl ions, or the pH
of HCl?

I found one reference (
http://www.sampleprep.duq.edu/dir/Chapter2/Chapter2.htm ) that indicates
that this occurs around a molar concentration of 3, but that seems
awfully high (pH of -.48?).

Thanks,

Marshall
f***@hotmail.com
2004-12-22 15:39:52 UTC
Permalink
Post by Marshall Dudley
Silver chloride has a solubility of about .89 ppm in cold water.
Accoding to Le Chatelier's Principle the solubility will decrease in the
presence of additional chloride ions. So the addition of HCl or NaCl
will cause the solubility to drop.
However at an excess of chloride of ions a dichloroargentate complex is
formed, which is highly soluble.
Does anyone know at what point the solubility turns around and starts
increasing again with respect to the concentration of Cl ions, or the pH
of HCl?
I found one reference (
http://www.sampleprep.duq.edu/dir/Chapter2/Chapter2.htm ) that
indicates
Post by Marshall Dudley
that this occurs around a molar concentration of 3, but that seems
awfully high (pH of -.48?).
The website mentions 3M Cl(-) ion not the hydrogen ion concentration.
Did you mean pCl of -0.48? (pH is defined to be a positive number from
0 to 14; I don't think negative pH is of any use except as negative log
of a number greater than 1)
Marshall Dudley
2004-12-22 16:02:31 UTC
Permalink
Post by Marshall Dudley
Post by Marshall Dudley
Silver chloride has a solubility of about .89 ppm in cold water.
Accoding to Le Chatelier's Principle the solubility will decrease in
the
Post by Marshall Dudley
presence of additional chloride ions. So the addition of HCl or NaCl
will cause the solubility to drop.
However at an excess of chloride of ions a dichloroargentate complex
is
Post by Marshall Dudley
formed, which is highly soluble.
Does anyone know at what point the solubility turns around and starts
increasing again with respect to the concentration of Cl ions, or the
pH
Post by Marshall Dudley
of HCl?
I found one reference (
http://www.sampleprep.duq.edu/dir/Chapter2/Chapter2.htm ) that
indicates
Post by Marshall Dudley
that this occurs around a molar concentration of 3, but that seems
awfully high (pH of -.48?).
The website mentions 3M Cl(-) ion not the hydrogen ion concentration.
Did you mean pCl of -0.48? (pH is defined to be a positive number from
0 to 14; I don't think negative pH is of any use except as negative log
of a number greater than 1)
I thought so too, but http://members.aol.com/profchm/pfactor.html
indicates:

For strong acid molar concentrations equal to or less than 1, the pH value
would have a value from 0-14.
One can have a pH that is a negative value in for example strong acid
solutions greater than 1 mole/liter.

So I am really confused. What I really want to know is if the Cl(-)
concentration of stomach acid (pH of about 1.5) is sufficient to allow any
significant solubility of AgCl.

Marshall
f***@hotmail.com
2004-12-22 17:24:22 UTC
Permalink
Post by Marshall Dudley
Post by f***@hotmail.com
The website mentions 3M Cl(-) ion not the hydrogen ion
concentration.
Post by Marshall Dudley
Post by f***@hotmail.com
Did you mean pCl of -0.48? (pH is defined to be a positive number from
0 to 14; I don't think negative pH is of any use except as negative log
of a number greater than 1)
I thought so too, but http://members.aol.com/profchm/pfactor.html
For strong acid molar concentrations equal to or less than 1, the pH value
would have a value from 0-14.
One can have a pH that is a negative value in for example strong acid
solutions greater than 1 mole/liter.
I wanted to say that we can have a negative pH mathematically; but note
that it was just a convineant scale of expressing hydrogen ion
concentration so if we have a concentration greater than 1 M [H(+)]; we
can not ignore acitivity coefficients and thus a 3 M HCl would not be
necessarily -log[3] but a number say 0.988 multpilied by 3. Also the pH
of 1 x 10^-7 M HCl would not be 7.
Post by Marshall Dudley
So I am really confused. What I really want to know is if the Cl(-)
concentration of stomach acid (pH of about 1.5) is sufficient to allow any
significant solubility of AgCl.
Perhaps not significant but we can not ignore Cl(-) from our diet which
would be present in the stomach and the organic complexing agents in
our foods may increase the solubility silver.
This data might be helpful.
http://faculty.uccb.ns.ca/~dkeefe/chem200/math/1998/mathtutorial.htm:

AgCl(s) <--> Ag+(aq) + Cl-(aq) pKsp = 9.75
AgCl(s)<--> AgCl(aq) pK0 = 6.70
AgCl(s) + Cl-(aq) <--> AgCl2- (aq) pK1 = 4.70

Substitute the value of Cl(-) in stomach in K1 =
[AgCl2(-)]/[AgCl][Cl(-)]
and substitue AgCl by manipulating the first two equations. This might
give us a rough idea of solubility of silver chloride in stomach acid.
Wilco Oelen
2004-12-22 21:48:07 UTC
Permalink
Post by f***@hotmail.com
Post by Marshall Dudley
Post by f***@hotmail.com
The website mentions 3M Cl(-) ion not the hydrogen ion
concentration.
Post by Marshall Dudley
Post by f***@hotmail.com
Did you mean pCl of -0.48? (pH is defined to be a positive number
from
Post by Marshall Dudley
Post by f***@hotmail.com
0 to 14; I don't think negative pH is of any use except as
negative
Post by f***@hotmail.com
log
Post by Marshall Dudley
Post by f***@hotmail.com
of a number greater than 1)
I thought so too, but http://members.aol.com/profchm/pfactor.html
For strong acid molar concentrations equal to or less than 1, the
pH
Post by f***@hotmail.com
value
Post by Marshall Dudley
would have a value from 0-14.
One can have a pH that is a negative value in for example strong acid
solutions greater than 1 mole/liter.
I wanted to say that we can have a negative pH mathematically; but note
that it was just a convineant scale of expressing hydrogen ion
concentration so if we have a concentration greater than 1 M [H(+)]; we
can not ignore acitivity coefficients and thus a 3 M HCl would not be
necessarily -log[3] but a number say 0.988 multpilied by 3. Also the pH
of 1 x 10^-7 M HCl would not be 7.
Post by Marshall Dudley
So I am really confused. What I really want to know is if the Cl(-)
concentration of stomach acid (pH of about 1.5) is sufficient to
allow any
Post by Marshall Dudley
significant solubility of AgCl.
Perhaps not significant but we can not ignore Cl(-) from our diet which
would be present in the stomach and the organic complexing agents in
our foods may increase the solubility silver.
This data might be helpful.
AgCl(s) <--> Ag+(aq) + Cl-(aq) pKsp = 9.75
AgCl(s)<--> AgCl(aq) pK0 = 6.70
How should I interpret these two numbers? The first one, I assume is a
measure, useful as solubility product, -log10([Ag(+)][Cl(-)])? Is the
other one a solubility "product" also (containing just one factor,
expressed as -log10([AgCl]))?

If I try to solve the equations for AgCl(s), added to pure water, then
I need a third equilibrium constant for
AgCl(aq) <---> Ag(+)(aq) + Cl(-)(aq),
otherwise I have too many unknowns. If I'm missing something, could you
please explain that to me?

My unknowns are:

[Ag(+)(aq)]
[Cl(-)(aq)]
[AgCl(aq)]

The species AgCl(s) is not part of the set of equations, see definition
of Ksp.
Post by f***@hotmail.com
AgCl(s) + Cl-(aq) <--> AgCl2- (aq) pK1 = 4.70
Probably you also should have AgCl(aq) + Cl(-) <---> AgCl2(-)(aq). How
is pK1 defined here? Is this -log10([AgCl2(-)(aq)]/[Cl(-)(aq)])?
Post by f***@hotmail.com
Substitute the value of Cl(-) in stomach in K1 =
[AgCl2(-)]/[AgCl][Cl(-)]
See above, is this the definition of K1? I think that you are 'mixing
up' AgCl(s) and AgCl(aq).
Post by f***@hotmail.com
and substitue AgCl by manipulating the first two equations. This might
give us a rough idea of solubility of silver chloride in stomach acid.
Using this, I come up with four equations and four unknowns, not quite
easily solved by hand, but with some software it should not be that
difficult. My unknowns are:

[Ag(+)(aq)]
[Cl(-)(aq)]
[AgCl(aq)]
[AgCl2(-)(aq)]

Because of the fairly large concentration of Cl(-), the equations can
be simplified considerably safely, by assuming [Cl(-)] constant. Then
the system reduces to a set of three equations with the three aquated
Ag-species as unknowns.

Wilco
Borek
2004-12-23 21:35:18 UTC
Permalink
Post by Wilco Oelen
Post by f***@hotmail.com
AgCl(s) <--> Ag+(aq) + Cl-(aq) pKsp = 9.75
AgCl(s)<--> AgCl(aq) pK0 = 6.70
AgCl(s) + Cl-(aq) <--> AgCl2- (aq) pK1 = 4.70
(...)
Post by Wilco Oelen
If I try to solve the equations for AgCl(s), added to pure water, then
I need a third equilibrium constant for
AgCl(aq) <---> Ag(+)(aq) + Cl(-)(aq),
otherwise I have too many unknowns. If I'm missing something, could you
please explain that to me?
I would assume that for the reaction

AgCl(aq) <---> Ag(+)(aq) + Cl(-)(aq)

dissociation constant is

pK = pKsp - pK0

Best,
Borek
--
BPP Marcin Borkowski, ul. Architektów 14, 05-270 Marki
If you know someone with dyslexia take a look at http://www.bpp.com.pl
Remove your.pants to email me directly :)
Wilco Oelen
2004-12-24 07:34:36 UTC
Permalink
Post by Borek
Post by Wilco Oelen
Post by f***@hotmail.com
AgCl(s) <--> Ag+(aq) + Cl-(aq) pKsp = 9.75
AgCl(s)<--> AgCl(aq) pK0 = 6.70
AgCl(s) + Cl-(aq) <--> AgCl2- (aq) pK1 = 4.70
(...)
Post by Wilco Oelen
If I try to solve the equations for AgCl(s), added to pure water, then
I need a third equilibrium constant for
AgCl(aq) <---> Ag(+)(aq) + Cl(-)(aq),
otherwise I have too many unknowns. If I'm missing something, could you
please explain that to me?
I would assume that for the reaction
AgCl(aq) <---> Ag(+)(aq) + Cl(-)(aq)
dissociation constant is
pK = pKsp - pK0
Yes! Simple log-math! I have been thinking how to solve this problem
quite some time. Probably my brains already were in Xmas-mode!

PS: Farooq, this is the missing point. Using a relation, like the one
stated here makes solving the equations easier. In fact, the different
equilibria relations are not independent (mathematically speaking).
Post by Borek
Best,
Borek
Thanks!

Wilco

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