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The solubility product of BaSO4 at 25 degree centigrade is 10POWER-9. What should be the concentration ofH2SO4 necessary to precipitate BaSO4 solution of 0.01 M Ba 2 ion?
Most Upvoted Answer
The solubility product of BaSO4 at 25 degree centigrade is 10POWER-9. ...
H2SO4 —> 2(H+) + (SO4)^-2
at t=0 s. 0 0
at complete dissociation
0 2s s
solubility product {ksp} = [ (Ba)^+2]×[ (SO4)^-2 ]
ksp = 10^(-9) { given }
[ (Ba)^+2 ] = 0.01 {given}

==>>>. 10^(-9) = 0.01 × s
=> s = 10^(-9) ÷ 0.01 = 10^(-7)
=> [ (SO4)^-2 ] = 10^(-7)

=> concentration of H2SO4 = 10^(-7)
Community Answer
The solubility product of BaSO4 at 25 degree centigrade is 10POWER-9. ...
Introduction
To determine the concentration of H2SO4 required to precipitate BaSO4 from a 0.01 M Ba2+ solution, we need to understand the solubility product (Ksp) of BaSO4.
Ksp of BaSO4
- The solubility product constant (Ksp) of BaSO4 is given as 10^-9.
- The dissociation of BaSO4 in water can be represented as:
BaSO4 (s) ⇌ Ba2+ (aq) + SO4^2- (aq)
Relation of Ksp to Ion Concentrations
- The Ksp expression for BaSO4 is:
Ksp = [Ba2+][SO4^2-]
- For BaSO4 to start precipitating, the product of the ion concentrations must exceed Ksp.
Calculating SO4^2- Concentration
- Given that [Ba2+] = 0.01 M, we can set up the equation:
Ksp = [0.01][SO4^2-] = 10^-9
- Rearranging gives:
[SO4^2-] = Ksp / [Ba2+] = (10^-9) / (0.01) = 10^-7 M
Concentration of H2SO4 Required
- H2SO4 dissociates in solution to provide SO4^2- ions:
H2SO4 → 2H+ + SO4^2-
- Therefore, to achieve a [SO4^2-] concentration of 10^-7 M, we need the same concentration of H2SO4:
[H2SO4] = 10^-7 M
Conclusion
- To precipitate BaSO4 from a 0.01 M Ba2+ solution, the necessary concentration of H2SO4 is 10^-7 M.
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The solubility product of BaSO4 at 25 degree centigrade is 10POWER-9. What should be the concentration ofH2SO4 necessary to precipitate BaSO4 solution of 0.01 M Ba 2 ion?
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