[Rabn]

How is the state of the equilibrium affected after adding NaCl?

[Mitch]

It'll push solutes out of the water phase.

[Hunt]

Mitch I did not get your point.

How to know whether a complex molecule is soluble in water or not ? For binary salts, there are certain solubility rules that serve for qualitative purposes. Is there anything like that for coordination compounds?

For clarity my question should have been,

Does precipitation occur at all between Cl- and [Cu(NH3)4]2+ ?

But what if NaCl is added, would the ppt be  [Cu(NH3)4]Cl2(s)  ?

[Mitch]

in general, salt pushes other solutes out of water. I have no clue whether your complex is soluble in water, I'm just stating a common principle of adding salt to solutions.

[Rabn]

I wonder if that complex with Cl(-) would even exist. If that complex does exist I'm sure that there is a Ksp associated with it. You should do a literature to see if it does exist. I say this because Cu's electron configuration with the four NH3 coordinated to it has a favorable half full D-shell. Cl has such a high electronegativity that it would form an ionic type bond to the complex. But as we know ionic Cl bonds usually dissociate in aqueous solutions.  So I don't think the Cl(-) complex would actually exist as a ppt. You would definitely need to do some search through literature. My guess is that adding NaCl would precipitate CuCl2 but the amount would be affected by the pH of the solution, which would become basic.

[Borek]

Copper (II) chloride solubility is in the range of 45 g (mass of anhydrous salt) per 100 g of H₂O (solution density well over 1.2). You may convert that to K_so, although saturated solution has a high ionic strength (around 16), so take the number obtained with a grain of salt. That's not a sparringly soluble salt, thats why K_so is not listed in tables. What was amount and concentration of HCl added?

My guess is that adding NaCl would precipitate CuCl2 but the amount would be affected by the pH of the solution, which would become basic.

Please elaborate - why do you expect pH to rise after NaCl addition?

[Rabn]

I suspect a pH rise because some of the released NH₃ ligands would surely become NH₄⁺ which would increase the [OH^-].

[Borek]

That'll be possible only if Cl^- will either remove NH₃ from the complex (with stablility constants of Cu²⁺/Cl^- complexes around 1 it aint gona happen) or if the CuCl₂ precipitates - not very likely with so high solubility.

[Rabn]

Thanks for making me double check myself Borek. I mistakenly looked at the Ksp of Copper(I) Chloride. Which in effect made me contradict myself. I'm curious about these stability constants you mentioned. I haven't run into those in my studies(I haven't taken any advanced inorganic/organic courses yet).  I'm going to do a search for some info, if you know of any good sites describing those would you mind posting a link?

[Borek]

I have several books with different aqueous equilibrium constants collections; however, these are mostly published in Eastern Europe back in seventies/eighties so they will be mostly likely out of reach for you.

[Hunt]

Mitch , nevermind
Borek , the experiment was qualitative.

What was amount and concentration of HCl added?

[AWK]

Stability constants for copper(II) ammonia system are given in Journal of Chemical Education 82(3) 408-14 (2005)
A.R. Johnson, T.M McQuinn, K.T. Rodolfa

[Borek]

Borek , the experiment was qualitative.

Still, there is completely different picture between adding 1mL of 1M HCl and 5mL of 35% HCl... I am not asking about exact numbers, rather about orders of magnitude.

[Hunt]

It wasn't more than 1 ml of HCl in a test tube, its concentration might have been 1 or 2 M. I cant really remember.

Is there any possibility that hydrogen gas was emitted during the process?

[AWK]

_{Is there any possibility that hydrogen gas was emitted during the process?}

Did you try writing down a balanced reaction?