Show, with equations, why the solution from which BaSo4 is precipitated cannot be allowed to be more then 0.01-0.05M in HCl.

To understand why the solution from which BaSO4 is precipitated cannot have a concentration of more than 0.01-0.05M in HCl, we need to consider the chemical reaction involved.

When BaSO4 (barium sulfate) is precipitated, it can be represented by the following balanced chemical equation:

BaCl2 (aq) + H2SO4 (aq) → BaSO4 (s) + 2HCl (aq)

From this equation, we can see that barium chloride (BaCl2) reacts with sulfuric acid (H2SO4) to form barium sulfate (BaSO4) and hydrochloric acid (HCl).

Let's assume we have a solution of barium chloride (BaCl2) with a concentration of [BaCl2] M, and we add sufficient sulfuric acid (H2SO4) to completely react with all the barium chloride. According to the stoichiometry of the reaction, 1 mole of barium chloride reacts with 1 mole of sulfuric acid to form 1 mole of barium sulfate and 2 moles of hydrochloric acid.

Therefore, if the initial concentration of barium chloride is [BaCl2] M, the concentration of hydrochloric acid formed will be twice that, which is 2[BaCl2] M.

Now, the maximum concentration of hydrochloric acid (HCl) that can be tolerated without interfering with the precipitation of barium sulfate is typically around 0.01-0.05 M. This concentration range ensures that the excess hydrochloric acid does not react with the formed barium sulfate to reform barium chloride.

So, to prevent the interference of hydrochloric acid with the precipitation of barium sulfate, the initial concentration of barium chloride (BaCl2) should not be allowed to exceed (0.01-0.05)/2 M, which is equal to 0.005-0.025 M.

In summary, the solution from which BaSO4 is precipitated cannot have a concentration of more than 0.01-0.05 M in HCl to prevent the reformation of barium chloride during the reaction.

To understand why the concentration of HCl should not exceed 0.01-0.05M when precipitating BaSO4, we need to consider the solubility product constant (Ksp) of BaSO4 and the concept of precipitation.

BaSO4 is an ionic compound that dissociates in water to form ions:
BaSO4(s) ⇌ Ba2+(aq) + SO42-(aq)

The solubility product constant (Ksp) expression for BaSO4 is:
Ksp = [Ba2+][SO42-]

The Ksp value for BaSO4 is approximately 1.1 x 10^-10 at 25°C.

When a compound is considered insoluble, such as BaSO4, it means that its concentration in solution is very low and virtually constant. This happens because the dissociation process in the forward direction is very weak compared to the backward recombination process.

Now, let's consider the effect of adding HCl to a BaSO4 solution. HCl is a strong acid that dissociates completely in water to form H+ ions:
HCl(aq) ⇌ H+(aq) + Cl-(aq)

The addition of HCl increases the concentration of H+ ions in the solution. According to Le Chatelier's principle, this increase in H+ concentration will shift the equilibrium of the solubility reaction of BaSO4 to the left, favoring the precipitation of more solid BaSO4.

However, if the concentration of H+ ions becomes too high, it can cause the reverse reaction to occur, resulting in the dissolution of BaSO4. This would lead to the redissolution of the precipitated BaSO4, which is undesirable.

To prevent the reverse reaction and ensure the precipitation of BaSO4, the concentration of H+ ions (or HCl) should be kept relatively low. By maintaining a lower HCl concentration (between 0.01-0.05M), the equilibrium is shifted towards precipitation and the precipitation process is favored.

In summary, the concentration of HCl should not exceed 0.01-0.05M when precipitating BaSO4 to prevent the reverse reaction and ensure successful precipitation based on the solubility product constant (Ksp) and the principles of equilibrium.