CBSE Class 12 Chemistry Revision Notes Chapter 1 Solutions

Solutions are homogeneous mixtures of two or more components, usually studied through solute, solvent, concentration and physical properties. In CBSE Class 12 Chemistry, this chapter explains liquid solutions, vapour pressure, colligative properties and abnormal molar mass.

Solutions explains how substances mix to form uniform mixtures. In daily life, pure substances are rare. Air, alloys, medicines, salt water and blood plasma are all linked with mixtures or solutions. The chapter mainly studies liquid solutions and their important properties.

Use these CBSE Class 12 Chemistry Revision Notes Chapter 1 to revise definitions, formulas, laws and examples from the 2026–27 chapter. Start with types of solutions and concentration units. Then revise solubility, Henry’s Law, Raoult’s Law, ideal solutions, non-ideal solutions and colligative properties.

Key Takeaways

• Solution: A homogeneous mixture of two or more components with uniform composition.
• Solvent: The component present in the largest amount and deciding the physical state.
• Henry’s Law: Gas solubility increases with pressure at constant temperature.
• Colligative Properties: Depend on the number of solute particles, not their chemical nature.

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Access 30 Minutes Class 12 Chemistry Chapter 1 Solutions Notes

This chapter can be revised quickly if you divide it into formulas, laws and comparison tables. Spend the first few minutes on concentration units because many numericals begin from them.

Then revise Henry’s Law, Raoult’s Law and colligative properties. These topics connect theory with numerical questions.

What are Solutions in Class 12 Chemistry Chapter 1?

A solution is a homogeneous mixture of two or more than two components. Its composition and properties remain uniform throughout the mixture.

Most examples in this chapter are binary solutions. A binary solution has two components.

Solute, Solvent and Binary Solution

The solvent decides the physical state of the solution. If water is the solvent, the solution is usually liquid.

Types of Solutions in Chemistry Chapter 1

Solutions may be gaseous, liquid or solid. The type depends on the physical state of the solvent.

Liquid solutions are most important in this chapter. These include gases, liquids or solids dissolved in a liquid solvent.

Concentration of Solutions: Important Formula Notes

Concentration tells how much solute is present in a given amount of solution or solvent. It can be expressed in different units.

Some units depend on mass, while some depend on volume. This is why temperature affects molarity but not molality.

Mass Percentage

Mass percentage gives the mass of a component in 100 parts by mass of solution.

Mass percentage = Mass of component / Total mass of solution × 100

Example: 10% glucose by mass means 10 g glucose is present in 100 g solution.

Volume Percentage

Volume percentage gives the volume of a component in 100 parts by volume of solution.

Volume percentage = Volume of component / Total volume of solution × 100

Example: 10% ethanol solution means 10 mL ethanol is present in 100 mL solution.

Mass by Volume Percentage

Mass by volume percentage is the mass of solute dissolved in 100 mL of solution.

It is commonly used in medicine and pharmacy.

Parts Per Million

Parts per million is used when the solute is present in very small quantity.

Parts per million = Number of parts of component / Total number of parts of all components × 10⁶

It is useful for expressing pollutants in water or air.

Mole Fraction

Mole fraction is the ratio of moles of one component to total moles of all components.

For component A:

xA = nA / (nA + nB)

For a solution, the sum of mole fractions is always 1.

x1 + x2 = 1

Mole fraction is very useful in vapour pressure and gas mixture calculations.

Molarity

Molarity is the number of moles of solute dissolved in one litre of solution.

Molarity = Moles of solute / Volume of solution in litre

Unit: mol L⁻¹ or M

Molarity changes with temperature because volume changes with temperature.

Molality

Molality is the number of moles of solute dissolved in one kilogram of solvent.

Molality = Moles of solute / Mass of solvent in kg

Unit: mol kg⁻¹ or m

Molality does not change with temperature because it depends on mass.

Quick Comparison of Molarity and Molality

Solubility in Solutions Class 12 Chemistry Notes

Solubility is the maximum amount of solute that can dissolve in a specified amount of solvent at a specified temperature.

It depends on the nature of solute and solvent, temperature and pressure.

Solubility of a Solid in a Liquid

A solid dissolves in a liquid when solute-solvent interactions are similar. This is often explained as like dissolves like.

Polar solutes dissolve in polar solvents. Non-polar solutes dissolve in non-polar solvents.

Examples:

When solute dissolves, dissolution takes place. When solute particles separate from solution, crystallisation takes place.

At equilibrium:

Solute + Solvent ⇌ Solution

A saturated solution contains the maximum amount of dissolved solute under given conditions.

Effect of Temperature on Solid Solubility

If dissolution is endothermic, solubility increases with rise in temperature.

If dissolution is exothermic, solubility decreases with rise in temperature.

Effect of Pressure on Solid Solubility

Pressure has no significant effect on the solubility of solids in liquids.

This is because solids and liquids are almost incompressible.

Solubility of a Gas in a Liquid

Gas solubility increases with pressure. More pressure means more gas particles strike the liquid surface and enter the solution.

Gas solubility usually decreases with rise in temperature. This is why aquatic life is more comfortable in cold water than warm water.

Henry’s Law in Class 12 Chemistry Chapter 1

Henry’s Law gives the relation between gas pressure and gas solubility in a liquid.

At constant temperature, the solubility of a gas in a liquid is directly proportional to the partial pressure of the gas above the liquid.

p = KHx

Here:

Higher KH means lower solubility of the gas at the same pressure.

Applications of Henry’s Law

Scuba divers use air diluted with helium to reduce the harmful effect of nitrogen under high pressure.

Vapour Pressure and Raoult’s Law Notes

Vapour pressure is the pressure exerted by vapour over a liquid at equilibrium.

In solutions, vapour pressure depends on the nature of components and their mole fractions.

Raoult’s Law for Volatile Liquid Solutions

For a solution of volatile liquids, the partial vapour pressure of each component is directly proportional to its mole fraction.

For component 1:

p1 = p1°x1

For component 2:

p2 = p2°x2

Total vapour pressure:

ptotal = p1 + p2

ptotal = x1p1° + x2p2°

This relation is important for numericals based on vapour pressure and mole fraction.

Raoult’s Law for Non-Volatile Solute

When a non-volatile solute is added to a volatile solvent, vapour pressure decreases.

This happens because solute particles occupy part of the surface. Fewer solvent molecules escape into vapour phase.

For solvent:

p1 = x1p1°

The lowering of vapour pressure depends on the amount of non-volatile solute, not its identity.

Raoult’s Law as a Special Case of Henry’s Law

Henry’s Law is written as:

p = KHx

Raoult’s Law is written as:

p = p°x

Both laws show that pressure is proportional to mole fraction. Raoult’s Law becomes a special case of Henry’s Law when KH becomes equal to p°.

Ideal and Non-Ideal Solutions

Liquid-liquid solutions are classified as ideal and non-ideal solutions on the basis of Raoult’s Law.

An ideal solution obeys Raoult’s Law over the entire range of concentration.

For ideal solutions:

ΔmixH = 0

ΔmixV = 0

This means no heat is absorbed or evolved during mixing. The volume of solution is equal to the sum of volumes of components.

Examples of nearly ideal solutions:

Non-Ideal Solutions

A non-ideal solution does not obey Raoult’s Law over the entire range of concentration.

It shows either positive deviation or negative deviation.

Positive Deviation from Raoult’s Law

Positive deviation occurs when solute-solvent interactions are weaker than interactions in pure components.

Molecules escape more easily. So vapour pressure becomes higher than expected.

Examples:

• Ethanol and acetone
• Carbon disulphide and acetone

Negative Deviation from Raoult’s Law

Negative deviation occurs when solute-solvent interactions are stronger than interactions in pure components.

Molecules do not escape easily. So vapour pressure becomes lower than expected.

Examples:

• Phenol and aniline
• Chloroform and acetone

Azeotropes

Azeotropes are binary mixtures having the same composition in liquid and vapour phase. They boil at a constant temperature.

They cannot be separated by fractional distillation at azeotropic composition.

Ethanol-water mixture forms an azeotrope containing about 95% ethanol by volume.

Nitric acid-water azeotrope contains about 68% nitric acid and 32% water by mass.

Colligative Properties and Molar Mass

Colligative properties depend on the number of solute particles in solution. They do not depend on the nature of solute particles.

There are four colligative properties in this chapter.

Relative Lowering of Vapour Pressure

When a non-volatile solute is added to a solvent, vapour pressure decreases.

Relative lowering of vapour pressure is equal to mole fraction of solute.

Δp1 / p1° = x2

For dilute solutions:

Δp1 / p1° = n2 / n1

Here, n2 is moles of solute and n1 is moles of solvent.

Elevation of Boiling Point

Boiling point increases when a non-volatile solute is added to a solvent.

Elevation of boiling point is directly proportional to molality.

ΔTb = Kb × m

Here:

This property helps in determining molar mass of solute.

Depression of Freezing Point

Freezing point decreases when a non-volatile solute is added to a solvent.

Depression of freezing point is directly proportional to molality.

ΔTf = Kf × m

Here:

Antifreeze used in car engines works on this principle.

Osmotic Pressure

Osmosis is the flow of solvent through a semipermeable membrane from pure solvent to solution.

Osmotic pressure is the pressure required to stop osmosis.

π = CRT

Here:

Osmotic pressure is useful for determining molar mass of macromolecules.

Abnormal Molar Mass and van’t Hoff Factor

Some solutes show abnormal molar mass because they associate or dissociate in solution.

Association means two or more particles combine. Dissociation means one particle splits into more particles.

The van’t Hoff factor explains this change.

i = Observed colligative property / Calculated colligative property

It can also be written as:

i = Normal molar mass / Abnormal molar mass

Example: NaCl dissociates into Na⁺ and Cl⁻, so the number of particles increases.

Quick Formula Table for Class 12 Chemistry Chapter 1 Solutions

Important Terms in Solutions Chapter

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