1.3416e-5
M
1.8000e-9
M
1.8000e-9
M
1.0000e-1
M
7,453.56
×
99.99
%
1.3416e-5
M
1.8000e-9
M
1.8000e-9
M
1.0000e-1
M
7,453.56
×
99.99
%
The Common Ion Effect Calculator determines how the presence of a common ion reduces the solubility of a sparingly soluble salt. When a solution already contains one of the ions produced by a dissolving salt, Le Chatelier's principle predicts that the equilibrium shifts to the left, decreasing solubility. This calculator compares solubility in pure water versus solubility in a solution containing a specified concentration of the common ion. The common ion effect is fundamentally important in qualitative analysis, selective precipitation, water treatment, and pharmaceutical formulation. It explains why adding NaCl to a solution of AgCl dramatically reduces silver chloride's solubility.
Consider salt MmAn dissolving in a solution already containing the anion A at concentration C:
$$K_{sp} = [M]^m \cdot ([A]_{added} + ns)^n$$
For a 1:1 salt (m = n = 1) with common anion at concentration C:
$$K_{sp} = s \cdot (C + s) \approx s \cdot C$$
Since s << C in most cases, the approximation gives:
$$s = \frac{K_{sp}}{C}$$
Compare this with solubility in pure water:
$$s_{pure} = \sqrt{K_{sp}} \quad \text{(for 1:1 salt)}$$
The reduction factor = spure / scommon shows how many times less soluble the salt becomes. The percent reduction quantifies the decrease.
A large reduction factor (e.g., >100×) means the common ion dramatically suppresses dissolution. Higher common ion concentrations produce greater reductions. The percent reduction close to 100% indicates nearly complete suppression of solubility. This is why the common ion effect is used in gravimetric analysis to ensure complete precipitation, and in water softening to remove dissolved minerals.
Inputs
Results
In pure water: s = √(1.8 × 10⁻¹⁰) = 1.34 × 10⁻⁵ M. With 0.1 M Cl⁻: s = 1.8 × 10⁻¹⁰/0.1 = 1.8 × 10⁻⁹ M. Solubility reduced by 7454×, or 99.99%.
Inputs
Results
PbCl₂ → Pb²⁺ + 2Cl⁻. In pure water: Ksp = 4s³, s = (1.7 × 10⁻⁵/4)^(1/3) = 0.0161 M. With 0.2 M Cl⁻: Ksp = s × (0.2)² = 0.04s, s = 4.25 × 10⁻⁴ M. 97% reduction.
The common ion effect is the decrease in solubility of a sparingly soluble salt when a soluble salt containing a common ion is added to the solution. It is a direct consequence of Le Chatelier's principle.
No. Ksp is a thermodynamic constant that depends only on temperature. The common ion effect changes the molar solubility but not Ksp itself.
Either the cation or anion can be the common ion, depending on what is already in solution. For example, adding NaCl provides Cl⁻ as common ion for AgCl, while adding AgNO₃ provides Ag⁺.
The decrease depends on the common ion concentration and salt stoichiometry. For a 1:1 salt, adding 0.1 M common ion typically reduces solubility by several orders of magnitude.
Yes. It is used in water treatment (precipitation of unwanted ions), toothpaste (fluoride and calcium equilibria), and drug formulation (controlling drug solubility).
The common ion effect involves an ion that participates in the solubility equilibrium. Salting out involves adding a different salt that decreases solubility through ionic strength effects, not equilibrium shifts.
No, the common ion effect always decreases solubility. However, adding certain ions that form complexes with the dissolved species can increase solubility (complex ion formation effect).
At high ionic strengths, activity coefficients deviate from 1, and actual solubility may be higher than predicted by the simple Ksp calculation. The Debye-Hückel equation can correct for this.
Yes, the common ion effect applies to acid-base equilibria too. Adding a common ion (e.g., acetate to acetic acid) suppresses dissociation and changes the pH. This is the basis of buffer solutions.
The calculator assumes the common ion concentration is much larger than the contribution from dissolution (C >> ns), allowing simplification of the equilibrium expression. This is valid when the common ion concentration is at least 100× larger than the pure-water solubility.
Roboculator Team
The Roboculator Team explains calculations, planning tools, and practical formulas in clear language for real-life situations.
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