What is the fundamental definition of the common-ion effect in chemistry?

In chemistry, the common-ion effect describes the reduction in solubility of an ionic precipitate when a soluble compound, which shares an ion with the precipitate, is added to the solution. This means that if you have a slightly soluble salt, and you add another salt that provides one of the same ions, the original salt will become even less soluble.

Which scientific principle governs the behavior observed in the common-ion effect?

The behavior of the common-ion effect is a direct consequence of Le Chatelier's principle. This principle states that if a dynamic equilibrium is disturbed by changing the conditions, the position of equilibrium shifts to counteract the change.

How does the common-ion effect specifically influence the solubility of ionic precipitates?

The common-ion effect leads to a decrease in the solubility of an ionic precipitate. When a common ion is added, the equilibrium of the ionic association/dissociation reaction shifts, causing more of the precipitate to form and thus reducing the amount of dissolved ions.

Beyond salts, what other types of chemical compounds commonly demonstrate the common-ion effect?

The common-ion effect is commonly observed as an effect on the solubility of salts and other weak electrolytes. Weak electrolytes are substances that only partially ionize or dissociate in a solution.

What happens to the concentration of both ions of a salt when an additional amount of one of its common ions is introduced into a solution?

Adding an additional amount of one of the ions of the salt generally leads to increased precipitation of the salt. This increased precipitation reduces the concentration of both ions of the salt in the solution until the solubility equilibrium is re-established.

Describe the ionization behavior of hydrogen sulfide (H₂S) when dissolved in an aqueous solution.

Hydrogen sulfide (H₂S) is classified as a weak electrolyte, meaning it only partially ionizes when dissolved in an aqueous solution. This partial ionization establishes an equilibrium between the un-ionized H₂S molecules and their constituent ions, H⁺ and HS⁻.

Write the equilibrium reaction that represents the partial ionization of hydrogen sulfide in water.

The equilibrium reaction for the partial ionization of hydrogen sulfide in an aqueous medium is represented as: H₂S ⇌ H⁺ + HS⁻. This shows that hydrogen sulfide can reversibly dissociate into a hydrogen ion and a hydrosulfide ion.

How is the acid dissociation constant (K_a) for hydrogen sulfide mathematically expressed?

According to the law of mass action, the acid dissociation constant (K_a) for hydrogen sulfide is expressed as: K_a = [H⁺,HS⁻] / [H₂S]. This expression relates the concentrations of the products (ions) to the concentration of the reactant (un-ionized acid) at equilibrium.

What is the nature of hydrochloric acid (HCl) as an electrolyte?

Hydrochloric acid (HCl) is a strong electrolyte, which means it ionizes almost completely when dissolved in a solution. Its dissociation reaction is represented as: HCl → H⁺ + Cl⁻.

When hydrochloric acid is added to an H₂S solution, which ion acts as the common ion?

When hydrochloric acid (HCl) is added to a hydrogen sulfide (H₂S) solution, the hydrogen ion (H⁺) acts as the common ion. Both H₂S and HCl produce H⁺ ions in solution.

According to Le Chatelier's principle, how does the addition of H⁺ ions from HCl affect the equilibrium of H₂S dissociation?

Due to the increase in the concentration of H⁺ ions from the added HCl, Le Chatelier's principle dictates that the equilibrium of the dissociation of H₂S will shift to the left. This shift occurs to counteract the stress of the added H⁺, consuming some of the excess H⁺ and HS⁻ to form more un-ionized H₂S, while keeping the K_a value constant.

What is the resulting impact on the concentration of un-ionized H₂S molecules and sulfide ions (HS⁻) when HCl is added?

When HCl is added, the dissociation of H₂S decreases, leading to an increase in the concentration of un-ionized H₂S molecules. As a result, the concentration of sulfide ions (HS⁻) in the solution also decreases.

What is the chemical formula for barium iodate, and what is its solubility product constant (K_sp)?

The chemical formula for barium iodate is Ba(IO₃)₂. Its solubility product constant (K_sp) is given as 1.57 x 10⁻⁹. The K_sp is a measure of the solubility of a sparingly soluble ionic compound.

What is the solubility of barium iodate in pure water?

The solubility of barium iodate in pure water is 7.32 x 10⁻⁴ M. This value represents the maximum concentration of barium iodate that can dissolve in pure water at a given temperature.

How does the presence of barium nitrate, Ba(NO₃)₂, affect the solubility of barium iodate?

In a solution that is 0.0200 M in barium nitrate, Ba(NO₃)₂, the solubility of barium iodate is significantly reduced. This occurs because barium nitrate introduces the common ion, barium (Ba²⁺), into the solution, shifting the solubility equilibrium of barium iodate.

Quantitatively, how much is the solubility of barium iodate reduced when it is dissolved in a 0.0200 M barium nitrate solution compared to pure water?

The solubility of barium iodate is reduced to 1.40 x 10⁻⁴ M when dissolved in a 0.0200 M barium nitrate solution. This is approximately five times smaller than its solubility in pure water, which is 7.32 x 10⁻⁴ M.

In the context of water treatment, how is the common-ion effect utilized to reduce water hardness?

In water treatment, the common-ion effect is widely used to reduce water hardness, particularly when drawing drinking water from chalk or limestone aquifers. Highly soluble sodium carbonate salt is added to the raw water to precipitate out sparingly soluble calcium carbonate, thereby removing hardness-causing calcium ions.

Which specific chemical is added to raw water to facilitate the precipitation of calcium carbonate?

Sodium carbonate is the specific chemical added to raw water to facilitate the precipitation of calcium carbonate. Sodium carbonate is a highly soluble salt that provides carbonate ions, which are common with calcium carbonate.

What is the commercial value of the calcium carbonate precipitate obtained from water treatment?

The very pure and finely divided precipitate of calcium carbonate generated during the water treatment process is a valuable by-product. It is used in the manufacture of products such as toothpaste.

Explain how the common-ion effect plays a role in the salting-out process during soap manufacturing.

The salting-out process, used in the manufacture of soaps, benefits from the common-ion effect. Soaps are sodium salts of fatty acids, and the addition of sodium chloride introduces a common ion (sodium, Na⁺) which reduces the solubility of the soap salts, causing them to precipitate.

What type of chemical compounds are soaps, and how does the addition of sodium chloride affect their solubility?

Soaps are chemical compounds that are sodium salts of fatty acids. The addition of sodium chloride reduces the solubility of these soap salts, causing them to precipitate out of the solution due to a combination of the common-ion effect and increased ionic strength.

Why do waters containing significant amounts of sodium ions, such as sea or brackish water, diminish the effectiveness of soap?

Sea, brackish, and other waters containing appreciable amounts of sodium ions (Na⁺) interfere with the normal behavior of soap due to the common-ion effect. The excess Na⁺ reduces the solubility of the soap salts, making the soap less effective at cleaning.

What are the two main components that characterize a buffer solution?

A buffer solution is characterized by containing an acid and its conjugate base, or alternatively, a base and its conjugate acid. These components work together to resist changes in pH.

How does the common-ion effect manifest when a conjugate ion is added to a buffer solution, specifically regarding its pH?

When a conjugate ion is added to a buffer solution, the common-ion effect causes a change in the pH of the solution. For example, adding the conjugate base to a weak acid buffer will increase the pH.

Describe the dissociation behavior of sodium acetate when dissolved in a solution.

Sodium acetate (NaCH₃CO₂) is a strong electrolyte, meaning it dissociates completely when dissolved in a solution. It breaks down into sodium ions (Na⁺) and acetate ions (CH₃CO₂⁻), as shown by the reaction: NaCH₃CO₂(s) → Na⁺(aq) + CH₃CO₂⁻(aq).

How does acetic acid behave as an electrolyte in solution?

Acetic acid (CH₃CO₂H) is a weak acid, which means it only ionizes slightly when dissolved in a solution. It establishes an equilibrium where only a small fraction of its molecules dissociate into hydrogen ions (H⁺) and acetate ions (CH₃CO₂⁻).

When both sodium acetate and acetic acid are present in the same solution, how do the acetate ions from sodium acetate influence the ionization of acetic acid?

When both sodium acetate and acetic acid are dissolved in the same solution, the acetate ions (CH₃CO₂⁻) from the completely dissociated sodium acetate act as a common ion. According to Le Chatelier's principle, these added acetate ions suppress the ionization of the weak acetic acid and shift its equilibrium to the left.

What is the overall effect on the percent dissociation of acetic acid and the pH of the solution due to the common-ion effect in an acetate buffer?

Due to the common-ion effect, the percent dissociation of acetic acid will decrease, meaning less of it ionizes. Consequently, the pH of the solution will increase, making it less acidic than a solution containing only acetic acid.

How does the hydronium concentration in a common-ion solution containing acetic acid and sodium acetate compare to a solution containing only acetic acid?

In a common-ion solution containing both acetic acid and sodium acetate, the hydronium (H⁺) concentration will be decreased compared to a solution containing only acetic acid. This reduction in hydronium ions makes the common-ion solution less acidic.

Under what circumstances might the common-ion effect not apply, particularly concerning certain metal compounds?

The common-ion effect may not apply to many transition-metal compounds, particularly when the formation of complex ions is involved. This scenario introduces different equilibria that are not part of the simple precipitation of salts from an ionic solution.

Provide a specific example of a transition-metal compound that defies the typical common-ion effect.

Copper(I) chloride (CuCl) is an example of a transition-metal compound that defies the typical common-ion effect. Although it is insoluble in water, it dissolves when chloride ions are added, such as from hydrochloric acid, even though chloride is a common ion.

Why does copper(I) chloride dissolve when chloride ions are added, despite chloride being a common ion?

Copper(I) chloride dissolves when chloride ions are added because it forms soluble complex ions, specifically CuCl₂⁻. The formation of these complex ions shifts the equilibrium away from the insoluble CuCl, allowing it to dissolve.

What type of chemical species is formed that causes copper(I) chloride to dissolve in the presence of additional chloride ions?

The formation of soluble coordination complex ions, such as CuCl₂⁻, is what causes copper(I) chloride to dissolve in the presence of additional chloride ions. These complex ions are stable and keep the copper in solution.

Define the uncommon-ion effect and list its alternative names.

The uncommon-ion effect, also known as the salt effect or the diverse-ion effect, describes a phenomenon where adding an ion that is not part of the precipitated salt itself can actually increase the solubility of the salt. This is sometimes referred to as salting in.

In contrast to the common-ion effect, how does the uncommon-ion effect influence the solubility of a salt?

In contrast to the common-ion effect, which decreases solubility, the uncommon-ion effect causes an increase in the solubility of a salt. This means more of the salt can dissolve in the presence of a non-common ion.

What is the primary mechanism or factor that causes the solubility of a salt to increase in the presence of an uncommon ion?

The primary mechanism for the uncommon-ion effect is that as the total ion concentration in the solution increases, inter-ion attraction within the solution becomes a significant factor. These attractions can stabilize the dissolved ions, making them less likely to precipitate.

How does the increased total ion concentration in the uncommon-ion effect impact the availability of ions for precipitation?

In the uncommon-ion effect, the increased total ion concentration leads to greater inter-ion attraction. This alternate equilibrium makes the ions of the sparingly soluble salt less available for the precipitation reaction, effectively increasing the salt's solubility.

Common-ion Effect Wiki: Ionic Equilibria Unveiled: The Common-Ion Effect in Chemical Systems

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What is the Common-Ion Effect?

Decreased Solubility Explained

In the realm of chemistry, the common-ion effect describes a phenomenon where the solubility of an ionic precipitate significantly decreases when a soluble compound containing an ion common to the precipitate is introduced into the solution.^[1] This principle is a direct manifestation of Le Chatelier's principle, which dictates that a system at equilibrium will adjust itself to counteract any applied stress.

Equilibrium Shifts

When an additional amount of one of the constituent ions of a sparingly soluble salt is added to a solution, the system's solubility equilibrium is perturbed. To re-establish equilibrium, the system shifts towards the formation of more solid precipitate, thereby reducing the concentrations of both ions in the solution until the solubility product constant (K_sp) is satisfied. This effect is particularly pronounced for weak electrolytes and sparingly soluble salts.

Illustrative Examples

Hydrogen Sulfide Dissociation

Consider hydrogen sulfide (H₂S), a weak electrolyte that partially ionizes in aqueous solution, establishing an equilibrium:

H₂S(aq) ⇌ H⁺(aq) + HS⁻(aq)

The acid dissociation constant (K_a) for this equilibrium is expressed as:

K_a = [H⁺][HS⁻] / [H₂S]

If a strong electrolyte like hydrochloric acid (HCl) is introduced, it dissociates almost completely:

HCl(aq) → H⁺(aq) + Cl⁻(aq)

The influx of H⁺ ions, common to both equilibria, stresses the H₂S dissociation. According to Le Chatelier's principle, the H₂S equilibrium shifts to the left, decreasing the dissociation of H₂S and consequently reducing the concentration of sulfide ions (HS⁻).

Barium Iodate Solubility

Barium iodate, Ba(IO₃)₂, is a sparingly soluble salt with a solubility product constant (K_sp) of 1.57 x 10⁻⁹. Its solubility in pure water is 7.32 x 10⁻⁴ M. The dissolution equilibrium is:

Ba(IO₃)₂(s) ⇌ Ba²⁺(aq) + 2IO₃⁻(aq)

However, if barium nitrate, Ba(NO₃)₂, a soluble salt, is added to the solution to achieve a 0.0200 M concentration of Ba²⁺ ions, the common-ion effect becomes evident. The increased concentration of Ba²⁺ ions shifts the equilibrium to the left, causing more Ba(IO₃)₂ to precipitate. This reduces the solubility of barium iodate to approximately 1.40 x 10⁻⁴ M, a reduction by a factor of about five.^[1]

Practical Solubility Effects

Water Treatment Applications

A significant practical application of the common-ion effect is observed in water treatment processes, particularly when dealing with hard water sourced from chalk or limestone aquifers. Sodium carbonate (Na₂CO₃), a highly soluble salt, is intentionally added to raw water. This addition facilitates the precipitation of sparingly soluble calcium carbonate (CaCO₃), effectively reducing the water's hardness. The finely divided calcium carbonate precipitate generated is often a valuable byproduct, utilized in industries such as toothpaste manufacturing.

Soap Manufacturing: Salting-Out

The "salting-out" process, a critical step in the manufacture of soaps, also leverages the common-ion effect. Soaps are essentially sodium salts of fatty acids. By adding sodium chloride (NaCl) to the soap solution, the concentration of Na⁺ ions (a common ion with the sodium fatty acid salts) increases. This increase reduces the solubility of the soap salts, causing them to precipitate out of the solution. This process is further enhanced by the general increase in ionic strength, which also contributes to the reduced solubility.

Soap Ineffectiveness in Hard Water

Conversely, the common-ion effect explains why soaps are less effective in sea, brackish, or other hard waters that contain appreciable amounts of sodium (Na⁺) ions. The presence of these excess Na⁺ ions reduces the solubility of the soap salts, preventing them from forming stable lather and effectively cleaning. This interference highlights the importance of water quality in the efficacy of cleaning agents.

The Buffering Effect

Stabilizing pH

The common-ion effect is fundamental to the operation of buffer solutions. A buffer solution is typically composed of a weak acid and its conjugate base, or a weak base and its conjugate acid.^[2] The presence of both species in significant concentrations allows the solution to resist drastic changes in pH upon the addition of small amounts of acid or base.

Acetic Acid and Acetate Example

Consider a solution containing both sodium acetate (NaCH₃CO₂) and acetic acid (CH₃CO₂H). Sodium acetate, being a strong electrolyte, dissociates completely:

NaCH₃CO₂(s) → Na⁺(aq) + CH₃CO₂⁻(aq)

Acetic acid, a weak acid, only partially ionizes, establishing an equilibrium:

CH₃CO₂H(aq) ⇌ H⁺(aq) + CH₃CO₂⁻(aq)

The addition of acetate ions (CH₃CO₂⁻) from sodium acetate acts as a common ion. According to Le Chatelier's principle, this increased concentration of acetate ions suppresses the ionization of acetic acid, shifting its equilibrium to the left. This results in a decrease in the hydronium (H⁺) concentration, making the common-ion solution less acidic than a solution containing only acetic acid, and thus increasing the pH.

Notable Exceptions

Complex Ion Formation

While the common-ion effect generally predicts a decrease in solubility, certain transition-metal compounds present an exception due to their ability to form stable complex ions. This phenomenon is not governed by the simple precipitation equilibria of ionic salts. For instance, copper(I) chloride (CuCl) is largely insoluble in water. However, upon the addition of chloride ions (e.g., from hydrochloric acid), its solubility surprisingly increases. This is because the added chloride ions react with CuCl to form soluble complex ions, such as CuCl₂⁻, effectively removing Cu⁺ ions from the simple dissolution equilibrium and driving the reaction towards dissolution.

The Uncommon-Ion Effect

Salting In and Diverse Ions

In contrast to the common-ion effect, the "uncommon-ion effect," also known as the "salt effect" or "diverse-ion effect," describes situations where the addition of an ion *not* common to the precipitated salt can actually *increase* the salt's solubility. This seemingly counterintuitive behavior arises because as the total ionic concentration in the solution increases, the inter-ionic attractions within the solution become more significant.^[3] These increased attractions effectively "shield" the ions of the sparingly soluble salt, making them less available for precipitation and thus enhancing their solubility. This alternate equilibrium pathway leads to a phenomenon often termed "salting in."

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Ionic Equilibria Unveiled

💡 Dive in with Flashcard Learning! 💡

What is the Common-Ion Effect? ⚛️

📉 Decreased Solubility Explained

⚖️ Equilibrium Shifts

Illustrative Examples 💡

💨 Hydrogen Sulfide Dissociation

💧 Barium Iodate Solubility

Practical Solubility Effects 🧪

🚰 Water Treatment Applications

🧼 Soap Manufacturing: Salting-Out

🌊 Soap Ineffectiveness in Hard Water

The Buffering Effect ⚗️

🛡️ Stabilizing pH

🍎 Acetic Acid and Acetate Example

Notable Exceptions 🚫

🔗 Complex Ion Formation

The Uncommon-Ion Effect ➕

⬆️ Salting In and Diverse Ions

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📜 Important Notice

Dive in with Flashcard Learning!

What is the Common-Ion Effect?

Decreased Solubility Explained

Equilibrium Shifts

Illustrative Examples

Hydrogen Sulfide Dissociation

Barium Iodate Solubility

Practical Solubility Effects

Water Treatment Applications

Soap Manufacturing: Salting-Out

Soap Ineffectiveness in Hard Water

The Buffering Effect

Stabilizing pH

Acetic Acid and Acetate Example

Notable Exceptions

Complex Ion Formation

The Uncommon-Ion Effect

Salting In and Diverse Ions

Teacher's Corner

True or False?

Test Your Knowledge!

Gamer's Corner

Discover other topics to study!

References

Feedback & Support

Disclaimer

Important Notice