Common ion effect

Introduction

The Common Ion Effect (CIE) is founded on Le Chatelier's Principle, which governs the chemical equilibrium of salts as well as other weak electrolytes including associated ions in solutions. This phenomenon is frequently employed to modify the solubilities of salts with weak electrolytes, as well as to separate salts from liquids. The common-ion effect (CIE) explains the dampening impact on the ionization of one electrolyte when some other electrolyte with a common ion is introduced. The CIE is the phenomenon wherein the solubility of a dispersed electrolyte decreases when some other electrolyte is introduced to the liquid with one ion that would be the same as the dispersed electrolyte.

What is the common ion effect?

This is typically focused on Le Chatelier's Principle, which states that if the proportion of any of the reactants is raised, the equilibrium moves in the path of the changes made to the equilibrium. Whenever a various salt, as well as an electrolyte containing one common ion, is followed by the addition containing an already disintegrated salt as well as an electrolyte, the composition of the common ion increases, causing the equilibrium of the dissolved salt or electrolyte to change to the initial salt or electrolyte side, allowing its solubility to decrease. Imagine the hypothetical salt 𝐴𝐵 plus electrolyte 𝐶𝐵. In this case, 𝐶𝐵 is a weaker electrolyte than 𝐴𝐵. Because when salt 𝐴𝐵 is introduced to a liquid that just contains 𝐶𝐵, because of the CIE, 𝐶𝐵 ionization decreases, as does its solubility, as well as the equilibrium moves to the left.

Effect on solubility

The existence of common ions influences the solubility of a salt in a liquid. In everyday life, this influence on solubility is exploited for several reasons.

This phenomenon is used to get water supply from aquifers comprising limestone or chalk. Aquifers are subsurface water layers that have been combined with permeable rocks and other elements. Sodium carbonate salt ($\mathrm{NA_{2}CO_{3}}$) is introduced to this combination to extract access to clean water. Because of the additional precipitation, the hardness of the water is reduced.

The addition of $\mathrm{NA_{2}CO_{3}}$ in water purification causes calcium carbonate ($\mathrm{CACO_{3}}$), which is slightly soluble in water, to precipitate away. This pure ppt is later employed in the production of toothpaste as a beneficial by-product of the phenomena of CIE.

This action is also very important in the manufacture of soap. They are fatty acid sodium salts. When $\mathrm{NaCl}$ is added to soap during the manufacturing process, the solubility of the soap salts is reduced. As a result of the CIE as well as an increase in ion strength, the soaps settle. This concept is also used in the filtration of salt. The introduction of a common ion to an already saturated salt solution causes additional salt to precipitate. This is also soaked with salt, refining the previously existent salt.

pH and the common ion effect

Whenever a conjugate ion is introduced to a buffer solution, the pH level changes due to the CIE.

This action may be seen, for example, when sodium acetate ($\mathrm{CH_{3}COOHNa}$) plus acetic acid ($\mathrm{CH_{3}COOH}$) is dissolved in a liquid to produce acetate ions. The $\mathrm{CH_{3}COOHNa}$ in the liquid entirely disintegrates, whereas the $\mathrm{CH_{3}COOH}$ only partially ionizes. This is because $\mathrm{CH_{3}COOH}$ is a weaker acid while sodium acetate is a stronger electrolyte.

The additional acetate ions produced by sodium acetate prevent $\mathrm{CH_{3}COOH}$ ionization, shifting the balance to the left thus lowering the pH of the solution. As a result, the CIE containing sodium acetate ($\mathrm{CH_{3}COOHNa}$) plus acetic acid ($\mathrm{CH_{3}COOH}$) will ultimately have a higher pH by being less acidic than an acetic acid liquid.

Examples of common ion effect

Observe what occurs when lead(𝐼𝐼) chloride ($\mathrm{PbCl_{2}}$) is dissolved in water but then sodium chloride ($\mathrm{NaCl}$) is added to the saturated liquid.

Because lead (𝐼𝐼)chloride ($\mathrm{PbCl_{2}}$) is very weakly soluble in water, the corresponding equilibrium exists:

$$\mathrm{Pb^{2+}(aq)\:+\:2Cl^{-}\:\rightarrow\:PbCl_{2}}$$

The resultant solution has double the amount of chloride ions plus lead ions. When sodium chloride is added to this solution, both lead(𝐼𝐼) chloride ($\mathrm{PbCl_{2}}$) but also sodium chloride bearing the chlorine anion is formed. Sodium chloride ($\mathrm{NaCl}$) ionizes to produce sodium as well as chloride ions −

$$\mathrm{Na^{2+}(aq)\:+\:Cl^{-}\:\rightarrow\:NaCl}$$

The extra chlorine anion produced by this reaction reduces the solubility of lead(II) chloride ($\mathrm{PbCl_{2}}$), changing the lead chloride reaction equilibrium to compensate for the 𝐶𝑙 addition. As a result, part of the chloride is removed being converted into lead(𝐼𝐼) chloride ($\mathrm{PbCl_{2}}$).

The CIE arises if a chemical is only sparingly soluble. In any liquid containing a common ion, the chemical will be less soluble. While the case with lead chloride ($\mathrm{PbCl_{2}}$) used a common anion, the same concepts apply to a common cation.

Factors involved in common ion effect

• Nature of electrolyte

• Temperature

• Volume/Dilution

• Mixing of ions

Importance of common effect

Circumstances are generated in QA to extract particular cations while leaving other cations in liquid. This is accomplished by adding an adequate quantity of the strong electrolyte with a common ion. A weak electrolyte's dissociation is inhibited to the required level.

The CIE was also used in the gravimetric analysis. Gravimetric precipitates are very slightly soluble in the precipitating liquid. Using the CIE, the precipitate's solubility may be reduced. This approach cannot be utilized for gravimetric estimation if the solubility is not poor. This is accomplished mostly by introducing a tiny excess of the precipitating agent with a common ion.

Conclusion

The precipitation of an ionic compound is determined by the presence of 1 or more common ions in a liquid. The precipitation of the mixture is governed by Le Chatiliers' Principle. When the amount of either of the ionic species rises, the reaction reverses to reduce the tension on the product side. We studied the common ion effect, its solubility, and its importance. Furthermore, we also gained knowledge about pH and common ion effects as well as some examples for better understanding in this article.

FAQs

1. What exactly is the common-ion effect (CIE)?

It is the result of Le Chatlier's principle. It is the phenomenon of suppressing one electrolyte's solubility by introducing a second electrolyte to the same liquid when the new electrolyte shares 1 ion with the electrolyte previously existing in the liquid.

2. Which of the following chemical pairings will exhibit the common-ion effect(CIE)?

• Sodium hydroxide ($\mathrm{NaOH}$) with sodium carbonate (𝑵𝒂𝑪𝑶𝟑)

• $\mathrm{Na_{2}SO_{4}}$with $\mathrm{Mg(OH)_{2}}$

• $\mathrm{CH_{3}COONa}$ with $\mathrm{CH_{3}COOH}$

Since they contain the common ions 𝑁𝑎 as well as $\mathrm{CH_{3}COO}$, the pairings "a" as well as "c" exhibit the common-ion effect.

3. Is the common-ion effect (CIE) the same as Le Chatelier's principle?

No, the Le Chatelier principle and CIE are not the same thing. Though the CIE is based mostly on Le Chatelier's principle, it differs other than that the CIE only applies to the concentration of equilibrium parts, including only ions of a molecule. Whereas Le Chatelier's concept may be used for quantities of many substances in equilibrium, including not just ions but also the influence of pressure as well as other factors, it can also be limited to various physical phenomena.

4. Clarify the common ion effect (CIE) of $\mathrm{H_{3}O^{+}}$ in acetic acid ionization?

When a strong acid produces the common ion, $\mathrm{H_{3}O^{+}}$ , the equilibrium shifts to produce additional $\mathrm{CH_{3}COOH}$.

$$\mathrm{H_{3}O^{+}\:+\:C_{2}H_{3}O_{2}^{-}\:\rightarrow\:CH_{3}COOH\:+\:H_{2}O}$$

5. Why does the common ion effect (CIE) of $\mathrm{NH_{4}OH}$ and $\mathrm{NH_{4}Cl}$ occur?

The $\mathrm{NH_{4}OH}$ is a weak base that does not entirely ionize. As a result of the existence of the common ion $\mathrm{NH_{4}}$ + in ammonium chloride $\mathrm{NH_{4}Cl}$, it inhibits the ionization of the weak base ammonium hydroxide $\mathrm{NH_{4}OH}$, lowering the 𝑂𝐻− concentration, therefore, preventing increasing group cations from precipitating. As a result, the pair $\mathrm{NH_{4}OH}$ plus $\mathrm{NH_{4}Cl}$ exhibits a CIE.