• Is it true that precipitates dissolve upon addition of diluted acids?
  • If a precipitates is dissolve , does it disappear from solution , or any thing else happened out there?

I am just curious about this. Because I have heard that precipitates can be dissolved if you add a diluted acid. But why it would happen? And is it valid for any diluted or strong acid ?.Do all the precipitates dissolve upon addition of any acid ?.I would like to know more about this.


Another example of a precipitate dissolving in acid is the reaction $\ce{AgNO3 + HCl -> AgCl (v) + HNO3}$. At first, the silver chloride precipitates out as a white powder (if kept away from direct light). Adding more acid, though, forms ligands with the additional chlorine ions: $\ce{AgCl(s) + Cl- (aq) -> AgCl2- (aq)}$. See also Physics Forums.

  • $\begingroup$ Then if ligands is formed upon addition of excess dil. acid does the precipitate dissolve? $\endgroup$ Jun 15 '15 at 1:12
  • $\begingroup$ Yes; see the references. $\endgroup$ Jun 15 '15 at 19:06

I would like to answer your second question first:

When a precipitate dissolves, it basically reacts to form some other soluble compound.

For your first question:

In many cases, a precipitate is soluble in a dilute acid solution. For example:

$$\ce{BaSO3 + 2HCl (dilute) -> BaCl2 (soluble) + SO2 (evolve) + H2O}$$

However is not necessarily true that a precipitate will dissolve in a dilute acidic solution. For example, calcium carbonate on reacting with dilute $\ce{H2SO4}$ evolves $\ce{CO2}$:

$$\ce{CaCO3 + H2SO4 -> CaSO4 (insoluble) + H2O + CO2 (evolve)}$$

It depends on the hydration enthalpy of the (would be) product after dissolution of a precipitate, whether the precipitate will be soluble or nor. In these examples, hydration enthalpy of $\ce{BaCl2}$ is −8.8 kJ/mol and that of $\ce{CaSO4}$ is −73 kJ/mol.

  • $\begingroup$ How can I relate hydration enthalpy to this? $\endgroup$ Jun 14 '15 at 5:11

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