Well I'll send it to the list then.
The original heading was tech alert...so Marshalee, don't read this!
I've been sitting exams, which is why I have been scarce!
Another property of Lewis acids:
Some metal cations form Lewis acids (including silver), and the total
dissolved concentration of the metal depends upon the concentration of the
complexing ion...say Cl-... as the concentration of Cl- increases the
concentration of Ag+ decreases (common ion effect) until a point is reached
where Ag+ redissolves as variations of AgCl(aq) are formed eg Ag++ Ag2Cl and
so on.
I haven't looked at the pH of Lewis acids yet but will do so, now that I
have a bit of time.
Original post follows:
Ok, does anybody else want to play with this?
Quantitative Chemical Analysis 5th Ed. - Harris
has this worked example:
What is the maximum Cl- concentration at equilibrium in a solution in which
[2Hg2+] is somehow FIXED at 1.0E-9 M? Our concentration table looks like
this:
Hg2Cl2(s)<===>Hg(2)2+ + 2Cl-
---------------------------------------------
Initial concentration 0 1.0E-9 0
Final concentration solid 1.0E-9 x
---------------------------------------------
[Hg(2)2+] is not x in this example, so there is no reason to set [Cl-]=2x.
The problem is solved by plugging each concentration into the solubility
product:
[Hg(2)2+][Cl-]^2 = Ksp
(1.0E-9)(x)^2 = 1.2E-18
x = [Cl-] = 3.5E-5 M
Here is my thinking;
For every Ag+ ion released into the water an OH- ion is created (due to the
release of H2 gas) so the [OH-] (concentration) increases at the same rate
as the [Ag+].
Now if [H+] was to remain the same as found in the water (1.0E-7) then using
the template described above (except that x is not squared because there is
only 1 Ag+ in AgOH) the equation looks like this:
[OH-][Ag+] = Ksp
(1.0E-7)[Ag+] = 1.52E-8
[Ag+] = 0.152M or 16,416 mg/L
However, as the amount of OH- rises the solubility of Ag+ falls, but to
counter this effect hydrated Ag+ acts as a Lewis acid forming complex ions
with 4(?) water molecules, and frees some H+ ions into the water and taking
some OH- ions out of the water (so to speak). I'm not sure how to account
for these effects completely, but it seems clear that the solubility of
silver ions, if introduced as ions is far higher than some would
believe....I think they are using an equation that only applies if there is
AgOH solid present in the solution, and the solution is at equilibrium.
Or have I got this completely wrong?
BTW Roger, if you are out there? I think you asked why the colloid becomes
acidic...it is probably due to the Lewis acid effect of transition metals.
Ivan.
> -----Original Message-----
> From: [email protected] [mailto:[email protected]]
> Sent: Wednesday, 20 June 2001 22:56
> To: [email protected]
> Subject: Re: CS>Theorizing About the Species Produced at Electrodes
> During LVDC CS Pro...
>
>
> In a message dated 6/20/01 4:53:59 AM EST, [email protected] writes:
>
> << Subj: Re: CS>Theorizing About the Species Produced at Electrodes
> During LVDC CS Production
> Date: 6/20/01 4:53:59 AM EST
> From: [email protected] (Arnold Beland)
> Reply-to: [email protected]
> To: [email protected]
>
> If you guys go off list with this, please include me.
>
> Arnold >>
>
> Arnold: You're the second person to say that. So far I haven't heard from
> Ivan. When I do, my response will be to the list, so stay tuned. Roger
>
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