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.

  • 1
    $\begingroup$ Your assumption is correct. $\endgroup$ – 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:

$\ce{ Ag+ + Cl- <=> AgCl }$

Originally you dissolve a certain amount of a chloride in distilled water. The concentration of chloride is wholly determined by this dissolved salt. If you do this in tap water, the amount of chloride will be higher as you stated, due to chlorination.

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 $\ce{AgCl}$.


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