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Common-ion effect
Common-ion effect
Common-ion effect
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In chemistry, the common-ion effect refers to the decrease in solubility of an ionic precipitate by the addition to the solution of a soluble compound with an ion in common with the precipitate.[1] This behaviour is a consequence of Le Chatelier's principle for the equilibrium reaction of the ionic association/dissociation. The effect is commonly seen as an effect on the solubility of salts and other weak electrolytes. Adding an additional amount of one of the ions of the salt generally leads to increased precipitation of the salt, which reduces the concentration of both ions of the salt until the solubility equilibrium is reached. The effect is based on the fact that both the original salt and the other added chemical have one ion in common with each other.

Examples of the common-ion effect

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Dissociation of hydrogen sulfide in presence of hydrochloric acid

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Hydrogen sulfide (H2S) is a weak electrolyte. It is partially ionized when in aqueous solution, therefore there exists an equilibrium between un-ionized molecules and constituent ions in an aqueous medium as follows:

H2S ⇌ H+ + HS−

By applying the law of mass action, we have

Hydrochloric acid (HCl) is a strong electrolyte, which nearly completely ionizes as

HCl → H+ + Cl−

If HCl is added to the H2S solution, H+ is a common ion and creates a common ion effect. Due to the increase in concentration of H+ ions from the added HCl, the equilibrium of the dissociation of H2S shifts to the left and keeps the value of Ka constant. Thus the dissociation of H2S decreases, the concentration of un-ionized H2S increases, and as a result, the concentration of sulfide ions decreases.

Solubility of barium iodate in presence of barium nitrate

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Barium iodate, Ba(IO3)2, has a solubility product Ksp = [Ba2+][IO3−]2 = 1.57 x 10−9. Its solubility in pure water is 7.32 x 10−4 M. However in a solution that is 0.0200 M in barium nitrate, Ba(NO3)2, the increase in the common ion barium leads to a decrease in iodate ion concentration. The solubility is therefore reduced to 1.40 x 10−4 M, about five times smaller.[1]

Solubility effects

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Main article: solubility equilibrium

A practical example used very widely in areas drawing drinking water from chalk or limestone aquifers is the addition of sodium carbonate to the raw water to reduce the hardness of the water. In the water treatment process, highly soluble sodium carbonate salt is added to precipitate out sparingly soluble calcium carbonate. The very pure and finely divided precipitate of calcium carbonate that is generated is a valuable by-product used in the manufacture of toothpaste.

The salting-out process used in the manufacture of soaps benefits from the common-ion effect. Soaps are sodium salts of fatty acids. Addition of sodium chloride reduces the solubility of the soap salts. The soaps precipitate due to a combination of common-ion effect and increased ionic strength.

Sea, brackish and other waters that contain appreciable amount of sodium ions (Na+) interfere with the normal behavior of soap because of common-ion effect. In the presence of excess Na+, the solubility of soap salts is reduced, making the soap less effective.

Buffering effect

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Main article: buffer solution

A buffer solution contains an acid and its conjugate base or a base and its conjugate acid.[2] Addition of the conjugate ion will result in a change of pH of the buffer solution. For example, if both sodium acetate and acetic acid are dissolved in the same solution they both dissociate and ionize to produce acetate ions. Sodium acetate is a strong electrolyte, so it dissociates completely in solution. Acetic acid is a weak acid, so it only ionizes slightly. According to Le Chatelier's principle, the addition of acetate ions from sodium acetate will suppress the ionization of acetic acid and shift its equilibrium to the left. Thus the percent dissociation of the acetic acid will decrease, and the pH of the solution will increase. The ionization of an acid or a base is limited by the presence of its conjugate base or acid.

NaCH3CO2(s) → Na+(aq) + CH3CO2−(aq)
CH3CO2H(aq) ⇌ H+(aq) + CH3CO2−(aq)

This will decrease the hydronium concentration, and thus the common-ion solution will be less acidic than a solution containing only acetic acid.

Exceptions

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Many transition-metal compounds violate this rule due to the formation of complex ions, a scenario not part of the equilibria that are involved in simple precipitation of salts from ionic solution. For example, copper(I) chloride is insoluble in water, but it dissolves when chloride ions are added, such as when hydrochloric acid is added. This is due to the formation of soluble CuCl2− complex ions.

Uncommon-ion effect

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Sometimes adding an ion other than the ones that are part of the precipitated salt itself can increase the solubility of the salt. This "salting in" is called the "uncommon-ion effect" (also "salt effect" or the "diverse-ion effect"). It occurs because as the total ion concentration increases, inter-ion attraction within the solution can become an important factor.[3] This alternate equilibrium makes the ions less available for the precipitation reaction. This is also called odd ion effect.

References

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Common-ion effect

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The common-ion effect is a phenomenon in chemical equilibria where the addition of an ion that is already present in the equilibrium system shifts the position of the equilibrium, typically suppressing the dissociation or solubility of the involved species according to . This effect occurs when a soluble compound introduces a "common "—an shared between the added compound and the equilibrium —leading to a reduction in the concentration of that from the equilibrium reaction. In the context of solubility equilibria, the common-ion effect significantly decreases the of sparingly soluble ionic salts. For instance, the of (AgCl), governed by the equilibrium AgCl(s) ⇌ Ag⁺(aq) + Cl⁻(aq) with a solubility product constant (Ksp) of 1.8 × 10−10, is approximately 1.3 × 10−5 M in pure but drops to about 1.8 × 10−9 M in a 0.10 M NaCl solution due to the added Cl⁻ ions shifting the equilibrium leftward. This principle can also be applied to prevent by maintaining ion concentrations below saturation levels through controlled addition of common ions. The common-ion effect extends to acid-base equilibria, particularly for weak acids and bases, where it influences and extent. For a weak acid like acetic acid (CH3COOH ⇌ H⁺ + CH3COO⁻), adding (NaCH3COO) introduces excess CH3COO⁻, driving the equilibrium left and decreasing [H⁺], which raises the solution's . A similar suppression occurs in weak base systems; for example, adding (NH4Cl) to (NH3 + H2O ⇌ NH4⁺ + OH⁻) provides NH4⁺, reducing [OH⁻] and lowering basicity, as demonstrated by the fading of indicator color from pink to colorless. This aspect is crucial in buffer solutions, where common ions from conjugate pairs stabilize against changes.

Definition and Mechanism

Definition

The common-ion effect refers to the reduction in the degree of ionization of a weak electrolyte or the solubility of a sparingly soluble ionic compound when another ionic compound sharing a common ion is introduced into the solution, causing a shift in the chemical equilibrium. This phenomenon occurs because the added common ion increases the concentration of one of the products in the dissociation equilibrium, suppressing further dissociation to maintain the equilibrium constant. The effect applies broadly to ionic solutions involving weak electrolytes, such as acids and bases that partially dissociate, and to sparingly soluble salts that establish low-concentration equilibria in water. Unlike the general influence of , which alters activities through electrostatic interactions across all species in solution, the common-ion effect specifically arises from the mass-action response to the elevated concentration of the shared . This process presupposes the concept of ionic dissociation, where a separates into its constituent ions in solution, as exemplified by the equilibrium for a weak acid:
\ceHA⇌H++A−\ce{HA ⇌ H+ + A-}
The introduction of additional \ceA−\ce{A-} ions from an external source shifts this equilibrium to the left, reducing the concentration of \ceH+\ce{H+}. The underlying driver is , which predicts that the system will counteract the change by favoring the reverse reaction.

Mechanism

The common-ion effect arises from the principles of , where the addition of an common to an existing equilibrium shifts the position of the equilibrium in response to the disturbance. According to , an increase in the concentration of a common —such as through the addition of a soluble salt—prompts the system to counteract this change by favoring the reverse reaction, thereby reducing the extent of or dissolution. This qualitative shift maintains the dynamic equilibrium in the solution, where forward and reverse processes occur continuously at equal rates until perturbed. In the context of weak electrolytes, such as a weak in , the common-ion effect manifests as a suppression of the electrolyte's dissociation. The presence of the common , often introduced via its conjugate base from a salt, increases the concentration of that in solution, driving the equilibrium toward the undissociated form of the weak electrolyte. This dominance of the reverse reaction reduces the percent dissociation, as the system adjusts to minimize the excess concentration per . The result is a lower concentration of from the weak electrolyte itself, preserving the balance of activities in the dynamic equilibrium. For sparingly soluble salts, the mechanism similarly involves a shift in the . Adding a source of the common elevates its concentration, prompting the equilibrium to favor of the solid to reduce the overall levels in solution. This qualitative reduction in occurs because the external common contributes to the ion product, pushing the system leftward to reattain equilibrium without relying solely on the salt's own dissolution. In this dynamic process, the rates of dissolution and adjust until the activities stabilize, underscoring the role of equilibrium in ion-involved solutions.

Quantitative Aspects

Equilibrium Constants and Derivations

The , KaK_a, quantifies the extent of dissociation for a weak acid HA in according to the equilibrium HA⇌H++A−\text{HA} \rightleftharpoons \text{H}^+ + \text{A}^- where Ka=[H+][A−][HA]K_a = \frac{[\text{H}^+][\text{A}^-]}{[\text{HA}]}. When a common , such as A⁻ from a salt like NaA, is introduced, it increases the total [A⁻] in solution, shifting the equilibrium to the left and suppressing [H⁺]. The derivation for the suppressed [H⁺] starts from the KaK_a expression, assuming the added common concentration dominates such that total [A⁻] ≈ [A⁻]₀ (initial added) + [A⁻] from dissociation, but for small dissociation, [H⁺] ≈ Ka[HA][A−]totalK_a \frac{[\text{HA}]}{[\text{A}^-]_{\text{total}}}, where [A⁻]ₙotal includes the common contribution. For weak bases, the base dissociation constant KbK_b describes the equilibrium B+H2O⇌BH++OH−\text{B} + \text{H}_2\text{O} \rightleftharpoons \text{BH}^+ + \text{OH}^- with Kb=[BH+][OH−][B]K_b = \frac{[\text{BH}^+][\text{OH}^-]}{[\text{B}]}, as exemplified by (NH₃). Adding a common like BH⁺ (e.g., from NH₄Cl) elevates total [BH⁺], suppressing further dissociation and reducing [OH⁻]. The suppressed [OH⁻] is derived analogously from the KbK_b expression: [OH⁻] ≈ Kb[B][BH+]totalK_b \frac{[\text{B}]}{[\text{BH}^+]_{\text{total}}}, incorporating the added common into the denominator. In solubility equilibria of sparingly soluble salts, the solubility product constant KspK_{sp} governs the dissolution, such as for silver chloride: AgCl(s)⇌Ag++Cl−\text{AgCl}(s) \rightleftharpoons \text{Ag}^+ + \text{Cl}^- where Ksp=[Ag+][Cl−]K_{sp} = [\text{Ag}^+][\text{Cl}^-]. Introducing a common ion, such as Cl⁻ from NaCl, increases [Cl⁻] total, driving the equilibrium toward the solid phase and reducing [Ag⁺] to maintain KspK_{sp}. The general form yields the suppressed ion concentration as [Ag⁺] = Ksp[Cl−]total\frac{K_{sp}}{[\text{Cl}^-]_{\text{total}}}, with [Cl⁻]ₙotal accounting for the common ion. The common-ion effect's mathematical foundation across these equilibria follows a unified derivation pattern: begin with the expression, incorporate the total concentration of the common into the relevant term, and solve for the suppressed species (e.g., [H⁺], [OH⁻], or solubility-related ), assuming negligible change from the weak process relative to the added . This approach highlights how the presence of the common directly inversely affects the dissociated concentration while keeping the invariant.

Calculations for Solubility

The common-ion effect can be quantitatively analyzed through solubility calculations for sparingly soluble salts, where the presence of an added common suppresses dissolution according to . Consider the of iodate, Ba(IO₃)₂, in a solution containing , Ba(NO₃)₂, as a representative example. The dissolution equilibrium is Ba(IO₃)₂(s) ⇌ Ba²⁺(aq) + 2 IO₃⁻(aq), with the product constant K_{sp} = [Ba²⁺][IO₃⁻]² = 4.0 × 10^{-9} at 25°C. Let s denote the molar of Ba(IO₃)₂ (in mol/L), and let C denote the initial concentration of the common Ba²⁺ from Ba(NO₃)₂ (assuming complete dissociation and no initial IO₃⁻). At equilibrium, [Ba²⁺] = C + s and [IO₃⁻] = 2s, so the K_{sp} expression becomes K_{sp} = (C + s)(2s)² = 4s²(C + s). To solve for s, rearrange the equation into the cubic form 4s³ + 4Cs² - K_{sp} = 0, which generally requires numerical methods or successive approximations for exact solutions. However, when C ≫ s (typically valid for low-solubility salts and moderate C values, such as C > 0.01 M), the approximation [Ba²⁺] ≈ C simplifies the equation to K_{sp} ≈ 4s²C, yielding s ≈ √(K_{sp} / (4C)). For instance, in a 0.020 M Ba(NO₃)₂ solution, s ≈ √(4.0 × 10^{-9} / (4 × 0.020)) = √(5.0 × 10^{-8}) ≈ 2.2 × 10^{-4} M, compared to the pure solubility of 1.0 × 10^{-3} M (from solving 4s³ = K_{sp}). This demonstrates a solubility reduction by a factor of about 4.5. The approximation's validity depends on the condition C ≫ s holding true, which fails for highly soluble salts or very low C (e.g., C ≈ 10^{-4} M), where the contribution of s to [Ba²⁺] becomes significant, leading to errors exceeding 20-25% if ignored. In such cases, successive approximations—iteratively substituting estimated s back into the full equation—or solving the cubic exactly is necessary; for example, with C = 10^{-4} M, initial approximation yields s ≈ 3.2 × 10^{-3} M (overestimating significantly as C ≪ s), but refinement gives s ≈ 9.5 × 10^{-4} M. Error analysis shows that for salts with K_{sp} > 10^{-6}, approximations often break down even at low C due to higher inherent . All concentrations are expressed in mol/L (M), assuming ideal behavior at low ionic strengths (μ < 0.01 M), where activity coefficients γ ≈ 1; at higher μ, non-ideal effects require correction via γ_{Ba^{2+}} and γ_{IO_3^-} in the thermodynamic K_{sp} = a_{Ba^{2+}} a_{IO_3^-}^2 = γ_{Ba^{2+}}[Ba²⁺] · (γ_{IO_3^-}[IO₃⁻])² [Ba²⁺][IO₃⁻]², potentially introducing 20-50% errors if neglected, but these are typically omitted in introductory calculations.

Illustrative Examples

Weak Electrolyte Dissociation

The common-ion effect significantly suppresses the dissociation of weak electrolytes when a soluble salt providing the common ion is added to the solution. A classic example is the ionization of acetic acid (\ceCH3COOH\ce{CH3COOH}), a weak acid with an acid dissociation constant Ka=1.8×10−5K_a = 1.8 \times 10^{-5} at 25°C. The equilibrium is: \ceCH3COOH⇌H++CH3COO−\ce{CH3COOH ⇌ H+ + CH3COO-} In a pure 0.10 M acetic acid solution, the percent dissociation is approximately 1.3%, corresponding to a hydrogen ion concentration [\ceH+]≈1.3×10−3[\ce{H+}] \approx 1.3 \times 10^{-3} M, calculated using the approximation for weak acids where [\ceH+]≈Ka⋅C[\ce{H+}] \approx \sqrt{K_a \cdot C}
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