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6.13 - SOLUBILITY EQUILIBRIA OF SPARINGLY SOLUBLE SALTS

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Overview of Solubility and Salts

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Teacher
Teacher

Today, we'll discuss solubility, particularly focusing on sparingly soluble salts and their equilibria. Can anyone tell me what solubility means in a chemical context?

Student 1
Student 1

Isn't it how much of a substance can dissolve in a solvent?

Teacher
Teacher

Exactly! Solubility refers to the maximum amount of solute that can dissolve in a solvent at equilibrium. Now, when we talk about sparingly soluble salts, can anyone think of an example?

Student 3
Student 3

Barium sulfate comes to mind!

Teacher
Teacher

Great example! Barium sulfate is a classic sparingly soluble salt with a very low solubility in water.

Lattice Enthalpy vs. Solvation Enthalpy

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Teacher
Teacher

Let's dive deeper into what affects the solubility of salts. We have two key concepts: lattice enthalpy and solvation enthalpy. Can someone explain what lattice enthalpy entails?

Student 2
Student 2

I think it’s the energy required to break apart the ions in the solid salt.

Teacher
Teacher

Correct! And solvation enthalpy is the energy released when ions interact with solvent molecules and get surrounded by them. It must exceed the lattice enthalpy for dissolution to occur. Can anyone tell me why non-polar solvents don’t dissolve ionic salts?

Student 4
Student 4

Because the solvation enthalpy is not enough to overcome the lattice enthalpy!

Teacher
Teacher

Exactly! Good job.

Solubility Product Constant (Ksp)

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Teacher
Teacher

Now let’s talk about Ksp, the solubility product constant. If we take barium sulfate again, what would its Ksp expression look like?

Student 1
Student 1

It's Ksp = [Ba2+][SO4^2–] right?

Teacher
Teacher

Exactly right! Ksp helps us understand the saturation level of the salt in solution. What changes can indicate the degree of solubility?

Student 3
Student 3

If we add more of one of the ions, like sulfate, it could precipitate the salt?

Teacher
Teacher

Correct! That’s called the common ion effect! This is a crucial concept because it impacts how we control solubility in various applications.

Common Ion Effect

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Teacher
Teacher

Lastly, let's cover the common ion effect. How does adding a common ion around a sparingly soluble salt affect its solubility?

Student 2
Student 2

It would decrease the solubility, right? Because the equilibrium shifts to reduce the concentration of that common ion?

Teacher
Teacher

Perfect! With increased common ion concentration, the equilibrium moves to favor the undissolved solid.

Student 4
Student 4

So, if we have sodium chloride in water and then add HCl, a common ion, sodium chloride precipitates, right?

Teacher
Teacher

Absolutely!

Applications

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Teacher
Teacher

Can anyone think of real-life scenarios where understanding solubility equilibria is useful?

Student 1
Student 1

In environmental science, like water treatment!

Teacher
Teacher

Exactly! Controlling solubility is essential in environmental management and industrial processes. Understanding Ksp is key to making predictions about contamination and purification techniques.

Introduction & Overview

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Quick Overview

This section discusses the solubility equilibria of sparingly soluble salts, emphasizing factors like lattice enthalpy and solvation enthalpy.

Standard

The chapter provides an overview of how the solubility of sparingly soluble salts in water is determined by the balance between lattice enthalpy and solvation enthalpy, defining solubility product constant (Ksp) and exploring the common ion effect on solubility.

Detailed

Detailed Summary

In this section, we delve into the solubility equilibria of sparingly soluble salts, exploring the factors involved in their dissolution in aqueous solutions. The primary concepts include lattice enthalpy, which is the energy required to separate ions in a structured solid, and solvation enthalpy, the energy released when ions interact with a solvent.

Factors affecting solubility are categorized into three categories:
- Category I: Soluble salts with a solubility greater than 0.1 M
- Category II: Slightly soluble salts with solubility between 0.01 M and 0.1 M
- Category III: Sparingly soluble salts with solubility less than 0.01 M

For example, the equilibrium involving barium sulfate can be expressed as:

BaSO4(s) ⇌ Ba2+(aq) + SO4^2–(aq)

The relevant equilibrium constant is known as the solubility product constant, denoted as Ksp, represented mathematically as:

Ksp = [Ba2+][SO4^2–]

The section highlights how Ksp is constant at a specific temperature and indicates the saturation level of the salt in solution. Different salts exhibit diverse solubility behaviors, which can be influenced by the presence of common ions. An increase in the concentration of a common ion typically decreases the solubility of the salt due to the common ion effect, shifting the equilibrium according to Le Chatelier's principle.

By understanding these concepts, one can analyze and predict the behavior of salts in various aqueous environments.

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Audio Book

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Solubility Variability

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We have already known that the solubility of ionic solids in water varies a great deal. Some of these (like calcium chloride) are so soluble that they are hygroscopic in nature and even absorb water vapour from atmosphere. Others (such as lithium fluoride) have so little solubility that they are commonly termed as insoluble.

Detailed Explanation

Solubility refers to the ability of a substance to dissolve in a solvent, particularly in water for ionic solids. The solubility of ionic compounds can differ significantly. For instance, calcium chloride dissolves readily in water, which means it has high solubility and can attract moisture from the air, making it hygroscopic. In contrast, lithium fluoride dissolves poorly in water, classifying it as insoluble due to its low solubility.

Examples & Analogies

Think of a sponge (highly soluble salt) that easily soaks up water, compared to a rock (insoluble salt) under water that doesn’t dissolve. While the sponge absorbs and holds water, the rock remains unchanged in the water.

Factors Influencing Solubility

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The solubility depends on a number of factors important amongst which are the lattice enthalpy of the salt and the solvation enthalpy of the ions in a solution.

Detailed Explanation

Several factors influence the solubility of salts in water. Lattice enthalpy is the energy required to separate ions in a solid crystal lattice. A higher lattice enthalpy means stronger attractions between ions, making the salt less soluble. Solvation enthalpy is the energy released when ions interact with the solvent. For dissolution to occur, the energy released via solvation must overcome the energy required to separate the ions, signifying that effective solvation can increase solubility.

Examples & Analogies

Imagine trying to break apart a tightly held group of friends (lattice enthalpy). If someone comes along (water) and offers them something exciting to attract their attention (solvation), they are more likely to separate and engage with that offer. Without sufficient attraction, they won't separate easily.

Classification of Salts by Solubility

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Each salt has its characteristic solubility which depends on temperature. We classify salts on the basis of their solubility in the following three categories:
Category I: Soluble Solubility > 0.1M
Category II: Slightly Soluble 0.01M<Solubility< 0.1M
Category III: Sparingly Soluble Solubility < 0.01M

Detailed Explanation

Salts can be categorized based on their solubility in water. Category I includes salts that dissolve greatly in water (greater than 0.1M), Category II includes slightly soluble salts (between 0.01M and 0.1M), and Category III includes sparingly soluble salts (less than 0.01M). This classification helps in understanding how much of a specific salt can be dissolved in a solution under given conditions.

Examples & Analogies

Think of it like different kinds of drinks. A soda (soluble) dissolves easily in water, whereas a syrup (slightly soluble) mixes but remains thick. Some compounds, like a pinch of salt in a glass of water (sparingly soluble), dissolve very little. This helps us understand how well or poorly components can mix.

Equilibrium of Sparingly Soluble Ionic Salts

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We shall now consider the equilibrium between the sparingly soluble ionic salt and its saturated aqueous solution.

Detailed Explanation

Sparingly soluble salts reach a state of equilibrium in a saturated solution where the rate of dissolution equals the rate of precipitation. In a saturated solution of a salt like barium sulfate, it exists in a dynamic balance where some solid salt keeps dissolving while an equal amount precipitates back. This creates a stable concentration of ions in the solution, allowing the salt to remain in equilibrium with its solid form.

Examples & Analogies

Imagine a busy coffee shop where customers continually enter and exit. People come in (dissolve) to buy coffee while others leave (precipitate) after enjoying it. Eventually, the number of customers remains stable while the processes of entering and leaving continue simultaneously.

Solubility Product Constant (Ksp)

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Let us now have a solid like barium sulphate in contact with its saturated aqueous solution. The equilibrium between the undisolved solid and the ions in a saturated solution can be represented by the equation:
BaSO4(s) Ba2+(aq) + SO42–(aq).
The equilibrium constant is given by the equation: K = {[Ba2+][SO42–]} / [BaSO4].

Detailed Explanation

The solubility product constant, Ksp, quantifies the extent of solubility of sparingly soluble salts. For barium sulfate, Ksp is expressed in terms of the concentrations of its dissociation products, barium ions and sulfate ions, while the concentration of the solid is omitted, as it remains constant. Therefore, Ksp equals the product of the concentrations of the ions at equilibrium.

Examples & Analogies

Think of Ksp as a recipe for a perfect balance. To make a cake, you need specific amounts of flour (Ba2+) and sugar (SO42–), but no matter how much cake you have, it doesn't change the amount of flour itself in the recipe.

Common Ion Effect on Solubility

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It is expected from Le Chatelier’s principle that if we increase the concentration of any one of the ions, it should combine with the ion of its opposite charge and some of the salt will be precipitated.

Detailed Explanation

The common ion effect describes how the solubility of a salt decreases when a common ion is added to the solution. By increasing the concentration of one of the ions involved in the solubility equilibrium, the system shifts equilibrium according to Le Chatelier’s principle—favoring the formation of the solid phase over the dissolved ions, thus leading to precipitation.

Examples & Analogies

Imagine a concert where everyone is clapping and cheering (ions in solution), but if someone starts a loud chant (adding a common ion), it shifts focus and people start standing at attention instead of clapping—that’s similar to how precipitation occurs in solution as equilibrium shifts.

Definitions & Key Concepts

Learn essential terms and foundational ideas that form the basis of the topic.

Key Concepts

  • Lattice Enthalpy: Energy needed to separate ions in a solid.

  • Solvation Enthalpy: Energy released during the interaction of ions with water.

  • Ksp: The equilibrium constant for the solubility of a sparingly soluble salt.

  • Common Ion Effect: The reduction in solubility of a salt due to the presence of a common ion.

Examples & Real-Life Applications

See how the concepts apply in real-world scenarios to understand their practical implications.

Examples

  • Barium sulfate has very low solubility in water, represented by a Ksp of 1.1 × 10–10.

  • The common ion effect can be observed when sodium chloride is added to a barium sulfate solution, decreasing its solubility.

Memory Aids

Use mnemonics, acronyms, or visual cues to help remember key information more easily.

🎵 Rhymes Time

  • Sparingly soluble salts, keep the ions at bay, Lattice must break, less they say.

📖 Fascinating Stories

  • Imagine a salty feast where each grain of salt relates to its watery cousins. Each salt has a story, where solvation helps it come together with friends in water.

🧠 Other Memory Gems

  • Remember: Lattice - L, Solvation - S; L before S - Salts dissolve after breaking bonds.

🎯 Super Acronyms

KSP

  • Keep Salt Present
  • ensures equilibrium!

Flash Cards

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Glossary of Terms

Review the Definitions for terms.

  • Term: Lattice Enthalpy

    Definition:

    The energy required to separate ions in a solid ionic compound.

  • Term: Solvation Enthalpy

    Definition:

    The energy released when ions are solvated or surrounded by solvent molecules.

  • Term: Solubility Product Constant (Ksp)

    Definition:

    An equilibrium constant for a sparingly soluble ionic compound in a saturated solution.

  • Term: Common Ion Effect

    Definition:

    The decrease in solubility of a sparingly soluble salt when a common ion is added to the solution.