Permafacture
2009-07-04 06:04:27 UTC
Hi all,
I had some chemistry in college but am trying to get a more intuitive
grasp on the subject now that I am graduated.
I've heard of using sodium/potassium/calcium hydroxide solutions as an
absorbent of carbon dioxide, since the carbonate salt is insoluble.
These salts are not easily decomposed however. using a metal with a
less energetic reaction would be more easily reversible. For
instance, Copper carbonate or silver carbonate are easily decomposed
by heat.
The most obvious ways I can think of using copper or silver ions to
absorb gaseous co2 is 1) to have the reaction happen at low temp and
high pressure, such that the carbonium ion causes some acidity and the
metal oxide is somewhat dissolved, and then reacts to form the
carbonate or 2) to use the nitrate salt of these metals.
in the second case, The precipitation of the carbonate would cause the
solution to become acidic (nitric acid)? (Is this true?) Transition
metal carbonates are soluble in acid, and so the reaction would reach
an equilibrium well before all the transition metal ions are
consumed. But I'm confused about how to set this up to figure out at
what point (pH) will the carbonate stop precipitating, or become so
slow as to be not worth while. It isn't a common ion situation, as
N03^- does not appear in the precipitation equation:
M^+ + CO3^2- -> MCO3 (s)
How do I go about this?
In practice, leaving some unreacted metal oxide (silver or Copper(I)
in this case) in the container would counter act this, adding more M+
as the solution becomes more acidic so that an arbitrary ammount of
CO2 gas could be sequestered with a given amount of nitrate salt. is
this true? (i guess that's a synthesis between my two possible
approaches).
Also, to get an idea of what pressure of CO2 could be developed in a
closed, heated container of this salt, i believe I would use the Gibbs
free energy. i tried doing this for calcium carbonate but did not
figure it out. is this the right path? If so i'll just keep plugging
away at it, perhaps check a different textbook. Any pointers to a
good reference would be appreciated.
much thanks,
elliot
I had some chemistry in college but am trying to get a more intuitive
grasp on the subject now that I am graduated.
I've heard of using sodium/potassium/calcium hydroxide solutions as an
absorbent of carbon dioxide, since the carbonate salt is insoluble.
These salts are not easily decomposed however. using a metal with a
less energetic reaction would be more easily reversible. For
instance, Copper carbonate or silver carbonate are easily decomposed
by heat.
The most obvious ways I can think of using copper or silver ions to
absorb gaseous co2 is 1) to have the reaction happen at low temp and
high pressure, such that the carbonium ion causes some acidity and the
metal oxide is somewhat dissolved, and then reacts to form the
carbonate or 2) to use the nitrate salt of these metals.
in the second case, The precipitation of the carbonate would cause the
solution to become acidic (nitric acid)? (Is this true?) Transition
metal carbonates are soluble in acid, and so the reaction would reach
an equilibrium well before all the transition metal ions are
consumed. But I'm confused about how to set this up to figure out at
what point (pH) will the carbonate stop precipitating, or become so
slow as to be not worth while. It isn't a common ion situation, as
N03^- does not appear in the precipitation equation:
M^+ + CO3^2- -> MCO3 (s)
How do I go about this?
In practice, leaving some unreacted metal oxide (silver or Copper(I)
in this case) in the container would counter act this, adding more M+
as the solution becomes more acidic so that an arbitrary ammount of
CO2 gas could be sequestered with a given amount of nitrate salt. is
this true? (i guess that's a synthesis between my two possible
approaches).
Also, to get an idea of what pressure of CO2 could be developed in a
closed, heated container of this salt, i believe I would use the Gibbs
free energy. i tried doing this for calcium carbonate but did not
figure it out. is this the right path? If so i'll just keep plugging
away at it, perhaps check a different textbook. Any pointers to a
good reference would be appreciated.
much thanks,
elliot