Grade 12
  1. Solid state  1.1. Classification of solids  1.2. Crystal Lattice and Unit Cell  1.3. Packing Efficiency  1.4. Density of unit cell  1.5. Imperfections in solids (point and line defects)  1.6. Electrical Properties of Solids (Conductors, Insulators and Semiconductors)  1.7. Magnetic Properties of Solids (Diamagnetic, Paramagnetic and Ferromagnetic Materials)
  2. Solution  2.1. Types of Solutions (Homogeneous and Heterogeneous)  2.2. Expressing concentration of solutions (mass %, volume %, molarity, molality, normality, mole fraction)  2.3. Solubility of solids and gases in liquids (Henry's law)  2.4. Raoult's Law and Ideal and Non-Ideal Solutions  2.5. Syndrome properties  2.6. Abnormal molar mass and Van't Hoff factor
  3. Electrochemistry  3.1. Electrochemical cells (galvanic and electrolytic cells)  3.2. Galvanic cell and EMF of the cell  3.3. Standard electrode potential and electrochemical series  3.4. Nernst equation and its applications  3.5. Conductivity and electrolytic cells (specific, molar and equivalent conductivity)  3.6. Kohlrausch's law and its applications  3.7. Batteries (primary and secondary cells)  3.8. Corrosion and its prevention
  4. Chemical kinetics  4.1. Rate of chemical reaction  4.2. Factors affecting the reaction rate (concentration, temperature, catalyst, surface area)  4.3. Order and molecularity of the reaction  4.4. Integrated Rate Equations (Zero, First and Second Order)  4.5. Half-life of a reaction  4.6. Collision theory and activation energy  4.7. Arrhenius equation and its applications
  5. Surface chemistry  5.1. Adsorption and its types (Physicosorption and Chemisorption)  5.2. Catalysis and its types (homogeneous and heterogeneous)  5.3. Colloids and their properties (Tyndall effect, Brownian motion, coagulation)  5.4. Emulsions and gels  5.5. Applications of Colloids
  6. General principles and processes of separation of elements  6.1. Presence of metals and minerals  6.2. Concentration of ores (gravity separation, froth flotation, magnetic separation)  6.3. Extraction of raw metals (roasting, calcination, reduction)  6.4. Thermodynamic and electrochemical principles of metallurgy  6.5. Refining of metals
  7. P-block elements  7.1. Group 15 elements (nitrogen and phosphorus)  7.2. Group 16 Elements (Oxygen, Sulphur and their Compounds)  7.3. Group 17 Elements (Halogens and their Compounds)  7.4. Group 18 Elements  7.5. Properties and Trends in p-Block Elements  7.6. Important compounds of p-block elements (ammonia, phosphine, sulphuric acid, hydrochloric acid)  7.7. Interhalogen compounds and their applications
  8. d-block and f-block elements  8.1. General Properties of Transition Metals  8.2. Properties of Inner Transition Metals (Lanthanoids and Actinoids)  8.3. Lanthanoid and actinoid contractions  8.4. Coordination Chemistry of d-Block Elements
  9. Coordination compounds  9.1. Werner's theory and modern concepts  9.2. Coordination number and type of ligand  9.3. Nomenclature of Coordination Compounds  9.4. Isomerism (geometrical and optical) in coordination compounds  9.5. Bonding in Coordination Compounds (Valence Bond Theory and Crystal Field Theory)  9.6. Stability and applications of coordination compounds in medicine and industry
  10. Haloalkanes and haloarenes  10.1. Classification and nomenclature of haloalkanes and haloarenes  10.2. Physical and Chemical Properties of Haloalkanes and Haloarenes  10.3. Mechanism of Nucleophilic Substitution Reactions (SN1 and SN2)  10.4. Uses and environmental effects (ozone depletion, CFCs)
  11. Alcohol, phenol and ether  11.1. Structure and classification of alcohols, phenols and ethers  11.2. Preparation and properties of alcohols  11.3. Preparation and properties of phenol  11.4. Preparation and properties of ethers  11.5. Chemical Reactions and Uses
  12. Aldehydes, ketones and carboxylic acids  12.1. Structure and nomenclature in aldehydes, ketones and carboxylic acids  12.2. Preparation and properties of aldehydes and ketones  12.3. Chemical reactions in aldehydes, ketones and carboxylic acids (nucleophilic addition, oxidation and reduction)  12.4. Preparation and Properties of Carboxylic Acids  12.5. Chemical Reactions and Uses of Carboxylic Acids
  13. Nitrogen-containing organic compounds  13.1. Amines (classification, structure, basicity)  13.2. Preparation and properties of amines  13.3. Reactions of Amines (Diazotization, Carbylamine Reaction)  13.4. Diazonium salts and their applications in paint industry
  14.1. Carbohydrates (classification and properties)  14.2. Proteins (Structure and Function)  14.3. Enzymes (Mechanism and Function)  14.4. Nucleic acids (DNA and RNA)  14.5. Vitamins and hormones
  15. Polymer  15.1. Classification of Polymers (Addition, Condensation, Copolymers)  15.2. Types of polymerization (free radical, ionic, condensation)  15.3. Important Polymers and Their Uses  15.4. Biodegradable and non-biodegradable polymers
  16. Chemistry in Everyday Life  16.1. Drugs and their classification (analgesics, antipyretics, antibiotics, antacids)  16.2. Chemicals in food and cosmetics (preservatives, artificial sweeteners, antioxidants)  16.3. Cleaning agents and detergents (soaps, synthetic detergents, surfactants)  16.4. Environmental Chemistry (Pollution, Green Chemistry, Sustainable Chemistry)

Grade 12 → Solution ↓


Solubility of solids and gases in liquids (Henry's law)


In chemistry, solubility is a fundamental property that describes how substances dissolve in a solvent to form a solution. In particular, the solubility of solids and gases in liquids is important in a variety of scientific and industrial processes. One of the guiding principles for understanding the solubility of gases in liquids is Henry's Law. Let's explore solubility with simple explanations, examples, and visual aids to develop a clear understanding of how it works.

What is solubility?

Solubility is defined as the maximum amount of solute (solid, liquid, or gas) that can dissolve in a solvent (usually a liquid) at a specified temperature and pressure to form a homogeneous solution. Solubility is often expressed in terms of concentration, such as grams of solute per 100 milliliters of solvent, or molarity, which is moles of solute per liter of solution.

Factors affecting solubility

Temperature

The effect of temperature on solubility depends on the nature of the solute and the solvent:

  • Solids: Most solid solutes become more soluble in liquid solvents as the temperature increases. A common example is sugar dissolving more quickly in hot water than in cold water.
  • Gases: The solubility of gases in liquids generally decreases as the temperature increases. For example, cold carbonated beverages retain their fizz better than beverages at room temperature because carbon dioxide is more soluble in cold liquids.

Pressure

Pressure has a more significant effect on the solubility of gases than on solids:

  • Increasing pressure increases the solubility of gases in liquids. This principle is famously illustrated by Henry's Law, which we will explore shortly.
  • Changes in pressure have a negligible effect on the solubility of solid solutes in liquid solvents.

Nature of solute and solvent

The chemical nature and structure of both the solute and the solvent play an important role in solubility:

  • Solutes that share the same chemical polarity with their solvents are generally more soluble. This is often summarized by the rule of thumb "like dissolves like." For example, table salt (NaCl) dissolves well in water, which is a polar solvent, but not in a nonpolar solvent such as oil.

Henry's law explained

Henry's law provides a quantitative relationship between the solubility of a gas in a liquid and the pressure of that gas above the liquid. It states that the amount of gas dissolved in a liquid is directly proportional to the partial pressure of that gas above the liquid.

c = kH * P

Where:

  • c is the concentration of the dissolved gas (e.g., in mol/L).
  • kH is the Henry's law constant, which varies depending on the gas, solvent, and temperature.
  • P is the partial pressure of the gas above the liquid.

Visual example of Henry's law

Consider a simple system where the gas is above the liquid, such as carbon dioxide above water in a closed bottle.

gas dissolved in water gas above liquid

Increasing the pressure of carbon dioxide increases its solubility in water, as Henry's Law shows. This relationship forms the basis of many industrial processes, such as the carbonation of beverages.

Textual example of Henry's law

Let us consider a scenario where the partial pressure of oxygen in water is doubled. According to Henry's law, the solubility of oxygen in water should also double. This concept explains processes such as the increase in oxygen absorption in the aquatic environment when atmospheric pressure increases.

Applications of Henry's law

Henry's law is not just theoretical; it has many practical applications:

Carbonation of beverages

Carbonated beverages such as soda and sparkling water are made by dissolving carbon dioxide gas in water under high pressure. When the pressure is reduced by opening the bottle or can, the gas escapes, producing the fizz.

Diving and decompression

When divers are underwater, the increase in pressure causes more nitrogen to dissolve in their blood. If they ascend too quickly, the rapid decrease in pressure causes the nitrogen to quickly come out of solution, forming bubbles that can cause decompression sickness, or "the bends."

Oxygen therapy in medicine

Henry's law is important for understanding how oxygen dissolves in the blood. In medical scenarios, increasing the partial pressure of oxygen can increase its solubility, helping patients who require supplemental oxygen therapy.

Understanding the solubility of solids

While Henry's law primarily addresses gases, the solubility of solid solutes in liquid solvents is equally important. Although pressure has little effect, temperature is an important factor for solids.

Visual example of solubility of solids

Let's look at the process of dissolving a solid substance like salt in water:

Salt crystals ions dissolved in water

A salt crystal dissociates into ions when dissolved in water. This process is facilitated by the polar nature of water molecules that interact with the ionic solid.

Textual example of solubility of solids

Consider adding a teaspoon of sugar to a cup of hot tea. You will notice that the sugar dissolves more quickly than if you add it to a cup of cold tea. This shows how temperature affects the solubility of solid solutes in liquids.

Conclusion

Understanding the solubility of solids and gases in liquids is essential in chemistry, which is primarily driven by principles such as Henry's law for gases. By controlling factors such as temperature and pressure, and considering the nature of the solute and solvent, we can predict and manipulate solubility to serve a variety of practical applications, from food and beverage production to medical treatment and environmental science.


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Solubility of solids and gases in liquids (Henry's law)

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