Professional Courses
Industry-relevant training in Business, Technology, and Design to help professionals and graduates upskill for real-world careers.
Categories
Interactive Games
Fun, engaging games to boost memory, math fluency, typing speed, and English skillsβperfect for learners of all ages.
Typing
Memory
Math
English Adventures
Knowledge
Enroll to start learning
Youβve not yet enrolled in this course. Please enroll for free to listen to audio lessons, classroom podcasts and take practice test.
Interactive Audio Lesson
Listen to a student-teacher conversation explaining the topic in a relatable way.
Overview of Solubility and Salts
Unlock Audio Lesson
Signup and Enroll to the course for listening the Audio Lesson
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?
Isn't it how much of a substance can dissolve in a solvent?
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?
Barium sulfate comes to mind!
Great example! Barium sulfate is a classic sparingly soluble salt with a very low solubility in water.
Lattice Enthalpy vs. Solvation Enthalpy
Unlock Audio Lesson
Signup and Enroll to the course for listening the Audio Lesson
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?
I think itβs the energy required to break apart the ions in the solid salt.
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 donβt dissolve ionic salts?
Because the solvation enthalpy is not enough to overcome the lattice enthalpy!
Exactly! Good job.
Solubility Product Constant (Ksp)
Unlock Audio Lesson
Signup and Enroll to the course for listening the Audio Lesson
Now letβs talk about Ksp, the solubility product constant. If we take barium sulfate again, what would its Ksp expression look like?
It's Ksp = [Ba2+][SO4^2β] right?
Exactly right! Ksp helps us understand the saturation level of the salt in solution. What changes can indicate the degree of solubility?
If we add more of one of the ions, like sulfate, it could precipitate the salt?
Correct! Thatβs called the common ion effect! This is a crucial concept because it impacts how we control solubility in various applications.
Common Ion Effect
Unlock Audio Lesson
Signup and Enroll to the course for listening the Audio Lesson
Lastly, let's cover the common ion effect. How does adding a common ion around a sparingly soluble salt affect its solubility?
It would decrease the solubility, right? Because the equilibrium shifts to reduce the concentration of that common ion?
Perfect! With increased common ion concentration, the equilibrium moves to favor the undissolved solid.
So, if we have sodium chloride in water and then add HCl, a common ion, sodium chloride precipitates, right?
Absolutely!
Applications
Unlock Audio Lesson
Signup and Enroll to the course for listening the Audio Lesson
Can anyone think of real-life scenarios where understanding solubility equilibria is useful?
In environmental science, like water treatment!
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
Read a summary of the section's main ideas. Choose from Basic, Medium, or Detailed.
Quick Overview
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) β Ba2+(aq) + SO4^2β(aq)
The relevant equilibrium constant is known as the solubility product constant, denoted as Ksp, represented mathematically as:
Ksp = [Ba2+][SO4^2β]
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.
Youtube Videos
Audio Book
Dive deep into the subject with an immersive audiobook experience.
Solubility Variability
Unlock Audio Book
Signup and Enroll to the course for listening the Audio Book
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 doesnβt dissolve. While the sponge absorbs and holds water, the rock remains unchanged in the water.
Factors Influencing Solubility
Unlock Audio Book
Signup and Enroll to the course for listening the Audio Book
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
Unlock Audio Book
Signup and Enroll to the course for listening the Audio Book
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
Unlock Audio Book
Signup and Enroll to the course for listening the Audio Book
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)
Unlock Audio Book
Signup and Enroll to the course for listening the Audio Book
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) + SO42β(aq).
The equilibrium constant is given by the equation: K = {[Ba2+][SO42β]} / [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 (SO42β), 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
Unlock Audio Book
Signup and Enroll to the course for listening the Audio Book
It is expected from Le Chatelierβs 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 Chatelierβs principleβfavoring 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 clappingβthatβs 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 Γ 10β10.
-
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.
π΅ Rhymes Time
-
Sparingly soluble salts, keep the ions at bay, Lattice must break, less they say.
π 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.
π§ Other Memory Gems
-
Remember: Lattice - L, Solvation - S; L before S - Salts dissolve after breaking bonds.
π― Super Acronyms
KSP
- Keep Salt Present
- ensures equilibrium!
Flash Cards
Review key concepts with flashcards.
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.