Sciencemadness Discussion Board - Lead(II) halides

Lead(II) halides

LanthanumK - 29-11-2011 at 05:13

I had a solution of lead acetate and tried to precipitate the chloride and bromide using sodium chloride and sodium bromide, respectively. No precipitate was formed. Why was that? Was it complexing with leftover acetic acid in the lead acetate or something like that?

ScienceSquirrel - 29-11-2011 at 05:17

Maybe your solution was not concentrated enough.
Lead bromide and chloride are slightly soluble in water so very dilute solutions will not produce a precipitate.

Adas - 29-11-2011 at 06:42

Sodium has lower electronegativity than lead, it will not react, am I right?

blogfast25 - 29-11-2011 at 07:00

Quote: Originally posted by Adas

Sodium has lower electronegativity than lead, it will not react, am I right?

Electronegativity has nothing to do with anything here. Sodium (here Na+

is just a spectator ion here.

In sufficient concetrations what happens is: Pb2+(aq) + 2 X-(aq) === > PbX2(s) with X = Cl, Br or I. Sodium ions do not take part.

Adas - 29-11-2011 at 07:28

Quote: Originally posted by blogfast25

Sodium ions do not take part.

Why is this?????? And how can I know whether an ion is spectator or not?

sxl168 - 29-11-2011 at 08:00

The concentration of NaCl also matters for solubility as per Chart

So don't use too much Cl or Br and chill the solution after mixing. If still nothing, probably started with store vinegar which is very dilute and you should evaporate the water from the Lead Acetate before trying to precipitate.

If you are just after a little PbCl2, I have found that boiling Pb in 10-15% HCl for 2 hours, or until you see PbCl2 crystals form works well. Just decant off the hot solution and let it cool. You end up with a bunch of PbCl2 after cooling and you can then decant off the acid from the cooled crystals back into your boiling Pb container to make more PbCl2 until the acid runs out by repeating the process over and over. When boiling, just have it so that the acid is just barely at the boiling point and cover your container for reflux, or else you will lose a lot of acid and generate lots of fumes.

Nicodem - 29-11-2011 at 09:19

Indeed, lead does dissolve in hydrochloric acid. It does so very slowly at room temperature, but if you leave it long enough to reach saturation (weeks!) you will see small crystals of PbCl2 growing.

Quote: Originally posted by Adas

Quote: Originally posted by blogfast25

Sodium ions do not take part.

Why is this?????? And how can I know whether an ion is spectator or not?

"Spectator ion" is a term to label ions that do not take part in the ionic reactions. In this particular precipitation reaction the Na+ ions have no specific role. They just stay there happily solvated doing nothing but entertaining the solvent (which also influences the Ksp equilibrium, but almost insignificantly at low concentrations). If these ions could anyhow interfere by a second reaction that would reduce the concentrations of the ions taking part in the Ksp equilibrium, then you they would not be spectator ions anymore (not ignorable!). For example, if CH3COO- would be able of coordinating with Pb(II) species or forming a covalent bond with it, then it could not be considered a spectator ion. And indeed the acetate anion does just that (DOI: 10.1021/ja01866a030 ; DOI: 10.1021/ja01984a013 ; pictures ).

Panache - 29-11-2011 at 13:59

Quote: Originally posted by Nicodem

They just stay there happily solvated doing nothing but entertaining the solvent

How do they do this? Maybe they start with a few jokes, a witty ancedote, then they serve some drinks and nibbles, compliment a few molecules individually on their dress sense and finally they might take the solvent out to a movie.

[Edited on 29-11-2011 by Panache]

LanthanumK - 30-11-2011 at 07:54

Back to Pb; ScienceSquirrel was partially correct. When I added HCl in place of the NaCl, a white precipitate was formed. It did appear that concentration was the problem.

Nicodem - 30-11-2011 at 09:31

Quote: Originally posted by LanthanumK

Back to Pb; ScienceSquirrel was partially correct. When I added HCl in place of the NaCl, a white precipitate was formed. It did appear that concentration was the problem.

That is a very unscientific conclusion! What were the respective concentrations anyway? And what makes you believe it was not due the acetate which is already reported in the scientific literature to increase the solubility of Pb(II) salts? Given that hydrochloric acid reduces the concentration of the acetate to insignificant levels and thus allows the precipitation (see the pKa of acetate anions) it is not prudent to blame it on the concentration like that.

ScienceSquirrel - 30-11-2011 at 09:42

Really it is all a bit hand wavy as we do not know the concentration of any of his solutions.
Suppose he had a very dilute solution of lead acetate and added a few drops of a fairly dilute solution of sodium chloride. A precipitate may or may not form depending on concentration and the presence of acetate ion that you mention.
Then suppose he adds a few drops of 30% hydrochloric acid to the second tube. I agree with you that this will protonate the acetate ion but he has also added vastly more chloride which will tend to force the precipitation of lead chloride anyway.

LanthanumK - 30-11-2011 at 09:52

The concentration of Cl(-) ions affects the PbCl2 solubility. When a large amount of Cl(-) is added, PbCl2 precipitates. When there is less Cl(-), PbCl2 does not precipitate. Of course, when more Cl(-) ions were added, the concentration of Ac(-) became more dilute and its effect was weakened. This seems to be something that would have reasonably happened.