# Precipitation of AgCl from the tap water solution of the group 2 chloride

I had this question about an experiment where the group 2 chloride was dissolved in distilled water and excess $$\ce{AgNO3}$$ was added to the solution to form $$\ce{AgCl}$$ precipitate. One of the sub-questions read

State how the amount of $$\ce{AgCl}$$ will change if tap water is used to dissolve the group 2 chloride instead of distilled water (will the amount of $$\ce{AgCl}$$ formed increase or decrease) and explain why.

Can someone please explain how the amount of $$\ce{AgCl}$$ formed will change if tap water is used? I assumed that tap water is chlorinated and so the amount of $$\ce{AgCl}$$ formed will increase. I don't know if this is right.

• Your assumption is correct. – andselisk Feb 18 '19 at 9:46

Yes the amount of $$\ce{AgCl}$$ formed will be more due to more $$\ce{Cl-}$$ ions present in the tap water. They are mainly because of some salts already present in water, and due to chlorination.
The amount of $$\ce{AgCl}$$ formed will increase. This is simple to justify if you look at the reaction:
$$\ce{ Ag+ + Cl- <=> AgCl }$$
Thus, if you take a look to the equilibrium above, the higher chloride concentration in the second experiment causes the equilibrium to "displace to the right", which ends up forming more $$\ce{AgCl}$$.