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Let the solubilities of AgCl in H2O,0.01MCaCl2; 0.01MNaCl and 0.05MAgNO3 be S1, S2, S3 and S4 respectively. What is the correct relationship between these quantities? (1) S1>S2>S3>S4 (2) S1>S2=S3>S4 (3) S1>S3>S2>S4 (4) S4>S2>S3>S1

Short Answer

Expert verified
The correct relationship is: S1>S2=S3>S4.

Step by step solution

01

- Identify the Solubility Product (Ksp) of AgCl

Identify the solubility product constant, Ksp, of AgCl. AgCl dissociates as AgClAg++Cl. The Ksp is given by Ksp=[Ag+][Cl].
02

- Solubility of AgCl in Pure Water (S1)

For AgCl in pure water, consider that there are no common ion effects. Assuming 'S' is the solubility, we have [Ag+]=S1 and [Cl]=S1.
03

- Solubility of AgCl in 0.01 M CaCl2 (S2)

Add a common ion effect. CaCl2 dissociates into Ca2+ and 2Cl. The presence of extra Cl decreases the solubility of AgCl. So, S2 will be less than S1.
04

- Solubility of AgCl in 0.01 M NaCl (S3)

Again, a common ion effect is present. NaCl dissociates into Na+ and Cl, adding more Cl to the solution. Thus, S3 will be less than S1. The solubility of AgCl will be similarly affected as with CaCl2, hence S3S2.
05

- Solubility of AgCl in 0.05 M AgNO3 (S4)

Add a common ion effect of Ag+. AgNO3 dissociates into Ag+ and NO3. The presence of extra Ag+ decreases the solubility of AgCl significantly. Therefore, S4 will be much less than S1.
06

- Conclusion

From these observations, we can see that S1 will be greater than S2S3 which will be greater than S4.

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Key Concepts

These are the key concepts you need to understand to accurately answer the question.

Solubility Product Constant (Ksp)
The Solubility Product Constant, commonly abbreviated as Ksp, is a key concept in understanding the solubility of sparingly soluble salts. For a salt like AgCl, it dissociates in water as:
AgClAg++Cl The Ksp expression for AgCl is given by:
Ksp=[Ag+][Cl] Here, [Ag+] and [Cl] represent the molar concentrations of the respective ions in a saturated solution at equilibrium. The Ksp value is fixed at a given temperature and indicates how much of the salt can dissolve to form a saturated solution.
Common Ion Effect
The Common Ion Effect refers to the decrease in solubility of an ionic compound when a common ion is added to the solution. For example:
  • When CaCl2 or NaCl is added to a solution containing AgCl, both dissociate to release Cl ions.
  • This increase in Cl ion concentration shifts the equilibrium to the left, reducing the solubility of AgCl.
Therefore, the solubilities S2 and S3 are less than S1, the solubility of AgCl in pure water.
Chemical Equilibrium
Chemical Equilibrium occurs when the rates of the forward and reverse reactions are equal, resulting in constant concentrations of reactants and products.
In the context of solubility, the equilibrium involving AgCl is described by the dissociation reaction:
AgClAg++Cl The equilibrium expression derived from this reaction reflects the product of the ion concentrations raised to the power of their coefficients in the balanced chemical equation.
Dissociation Reactions
Dissociation Reactions involve the separation of an ionic compound into its individual ions in solution. For instance:
  • AgCl dissociates as AgClAg++Cl
  • Similarly, CaCl2 dissociates into Ca2+ and Cl ions as CaCl2Ca2++2Cl
The introduction of additional ions from these dissociations can affect the equilibrium and solubility of other compounds, demonstrating the impact of the common ion effect.
Molar Solubility
Molar Solubility is defined as the number of moles of a solute that can dissolve per liter of solution to form a saturated solution. It directly relates to the Ksp value for a given compound.
If the molar solubility of AgCl in pure water is S1, then:
Ksp=S12 When common ions are added, such as Cl ions from CaCl2 or NaCl, the molar solubility decreases. This effect is due to the shift in equilibrium, aligning with Le Chatelier’s principle, thereby reducing the amount of AgCl that can dissolve in the presence of common ions.

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