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Precipitation reactions with AgNO3 and NaCl

  1. Jun 14, 2013 #1
    When we add them together, AgCl precipitates as AgCl is insoluble in water so their interactions would cause them to form a solid. However, isn't the interaction between Na+ and Cl- greater than in Ag+ and Cl-? As NaCl has a stronger lattice energy than AgCl. So why is AgCl insoluble while NaCl is soluble?

    I thought if the ionization energy is less for AgCl it should be more soluble than NaCl as less energy is able to break the AgCl into ions. So when placing AgCl and NaCl into 2 different containers of water, shouldn't more AgCl be able to break up into Ag+ and Cl-?

    Thanks for the help :)
     
  2. jcsd
  3. Jun 14, 2013 #2

    Borek

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    Staff: Mentor

    You can't ignore water and cation hydration.
     
  4. Jun 15, 2013 #3
    Oh why is that so? Actually what goes on when i dissolve them? The NaCl just breaks into Na+ and Cl- then it gets solvated?

    So using the Hess's Law, NaCl(s)->Na+(g)+Cl-(g)
    then Na+(g)+aq->Na+(aq) and Cl-(g)->Cl-(aq)? Seems weird to me because I feel that these 2 processes occur simultaneously. And it feels funny because when they turn into gases I feel that a lot of heat should we supplied to it (more than what the water has at rtp)

    What the right way to think about it? And what's wrong with how I'm looking at it?

    Thanks Borek!
     
  5. Jun 15, 2013 #4

    Borek

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    No, they are not gaseous in between, dissolution means ions get surrounded by water molecules, these are electric dipoles, so they are attracted to the ion and in a way separate it from other ions. That happens in the solution, and is covered by the

    NaCl(s) -> Na+(aq) + Cl-(aq)

    (aq) means just that - solvated ion.

    That in turn means the interactions between cation and anion are at least partially replaced by the interaction between the ions and water molecules. Water molecules get ordered around ions, lowering solution entropy, at the same time they usually give off heat, so you have plenty of additional thermodynamic effects.
     
  6. Jun 15, 2013 #5
    Oh! So actually what's the reason why AgCl is less soluble than NaCl? Because it should be easier to remove Ag+ from the Cl- ion so shouldn't the AgCl(s) ->Ag+(aq)+Cl-(aq) enthalpy change have a smaller number than for the NaCl case?
     
  7. Jun 15, 2013 #6

    Borek

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    You would have to compare these numbers precisely. Ag+ is a large cation, so it doesn't attract water molecules so strongly (it has identical charge as Na+, but larger radius, so the Coulomb attraction is much smaller). That means it is not that strongly stabilized in teh solution as Na+ is.

    But that's not all - so far we were comparing both salts as if they were just ionic - they are not. Compare Na and Ag electronegativities - AgCl is much more covalent, so the comparison is not that easy.
     
  8. Jun 15, 2013 #7
    Ohh so even though its easier to seperate the Ag+ and the Cl- it's harder for the water to solvate the Ag+ due to it having a small charge density, so less water actually tries to seperate them?
     
  9. Jun 15, 2013 #8

    Borek

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    You can put it this way.
     
  10. Jul 3, 2013 #9
    Ohh but actually, why would it precipitate as AgCl? Because alone AgNO3 is ok being aqueous as well as NaCl. But when I place them together, why should it suddenly precipitate?

    Even though Ag+ is unstable its fine when NO3- is the only anion. But when Cl- is present then why would the Ag+ become more unstable suddenly?

    Thanks :)
     
  11. Jul 6, 2013 #10
    Also, consider that AgCl is less ionic and more covalent than NaCl, so the the reaction:

    AgCl + (n+m)H2O --> [Ag(H2O)n]+ + [Cl(H2O)m]-

    is less favoured than:

    NaCl + (r+s)H2O --> [Na(H2O)r]+ + [Cl(H2O)s]-.
     
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