Questions based on concentration calculations, vapor pressure, osmotic pressure, and colligative properties often require careful formula application and unit conversion in NEET Chemistry. Many students face issues while connecting concepts like mole fraction, Raoult’s Law, Henry’s Law, and Van ’t Hoff factor within numerical problems. 

The Solutions chapter in Chemistry explains how solutes behave in different solvents and how dissolved particles affect the physical properties of solutions. Regular revision and numerical practice help improve calculation accuracy and conceptual understanding. 

Definition and Types of Solutions

A solution is a homogeneous mixture in which one or more substances are completely dissolved in another substance. A solution contains two components:

• Solvent: The component present in larger quantity that determines the physical state of the solution.

• Solute: The component present in smaller quantity that gets dissolved in the solvent.

Solutions can exist in solid, liquid, or gaseous states depending on the physical state of the solute and solvent.

Major Types of Solutions

The properties of a solution depend on the nature of the solute, solvent, temperature, and concentration.

Concentration Terms

Concentration terms describe the amount of solute present in a given amount of solvent or solution. Some concentration terms depend on temperature, while others remain independent of temperature.

Temperature-Dependent Concentration Terms

Molarity (M)

Molarity is defined as the number of moles of solute present in one litre of solution.

Formula:

M = Number of moles of solute / Volume of solution in litres

Alternative Formula:

M = (W × 1000) / (Molar Mass × Volume in mL)

Where:

• W = mass of solute in grams

• Volume = volume of solution in mL

Normality (N)

Normality is defined as the number of gram equivalents of solute present in one litre of solution.

Formula:

N = Gram equivalents of solute / Volume of solution in litres

Relation between Normality and Molarity:

N = M × n-factor

Mass by Volume Percentage (% w/V)

Formula:

% w/V = (Mass of solute in grams / Volume of solution in mL) × 100

Temperature-Independent Concentration 

Molality (m)

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

Formula:

m = Number of moles of solute / Mass of solvent in Kg

Alternative Formula:

m = (W × 1000) / (Molar Mass × Mass of solvent in grams)

Mole Fraction (X)

Mole fraction is the ratio of the number of moles of one component to the total number of moles present in the solution.

For a binary solution:

Xₐ = nₐ / (nₐ + nᵦ)

Xᵦ = nᵦ / (nₐ + nᵦ)

Also:

Xₐ + Xᵦ = 1

Mass Percentage (% w/w)

Formula:

% w/w = (Mass of component / Total mass of solution) × 100

Parts Per Million (ppm)

Used for expressing very small concentrations.

Formula:

ppm = (Mass of component / Total mass of solution) × 10⁶

Solubility and Henry’s Law

Solubility refers to the maximum amount of solute that can dissolve in a given amount of solvent at a specific temperature.

For gases dissolved in liquids, solubility depends strongly on pressure and temperature.

Henry’s Law

Henry’s Law states that at constant temperature, the solubility of a gas in a liquid is directly proportional to the partial pressure of the gas above the liquid.

Formula:

P = Kₕ × X

Where:

• P = partial pressure of gas

• Kₕ = Henry’s Law constant

• X = mole fraction of gas

A higher value of Kₕ indicates lower solubility of gas.

Applications of Henry’s Law include:

• Carbonated beverages

• Deep-sea diving

• High-altitude breathing problems

Vapor Pressure and Raoult’s Law

Vapor pressure is the pressure exerted by vapours in equilibrium with the liquid in a closed container.

Raoult’s Law explains how vapor pressure changes when solutions are formed.

Raoult’s Law for Volatile Solutions

For volatile liquids:

Pₐ = Pₐ° × Xₐ

Pᵦ = Pᵦ° × Xᵦ

Total vapor pressure:

P(total) = Pₐ + Pᵦ

P(total) = Pₐ°Xₐ + Pᵦ°Xᵦ

Where:

• Pₐ° and Pᵦ° are vapor pressures of pure components

• Xₐ and Xᵦ are mole fractions

Raoult’s Law for Non-Volatile Solutes

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

Formula:

P(solution) = P° × X(solvent)

The lowering of vapor pressure depends on the mole fraction of solvent.

Ideal and Non-Ideal Solutions

Ideal Solutions

Ideal solutions obey Raoult’s Law over the entire concentration range.

Conditions for ideal solutions:

ΔH(mixing) = 0

ΔV(mixing) = 0

In ideal solutions, intermolecular forces between unlike molecules are nearly equal to those between like molecules.

Examples:

• Benzene + Toluene

• n-Hexane + n-Heptane

Non-Ideal Solutions

Non-ideal solutions do not obey Raoult’s Law and show deviations.

Positive Deviation

Occurs when:

P(total) > Pₐ°Xₐ + Pᵦ°Xᵦ

Conditions:

• ΔH(mixing) > 0

• ΔV(mixing) > 0

Intermolecular forces between unlike molecules are weaker.

Example:

• Ethanol + Acetone

Negative Deviation

Occurs when:

P(total) < Pₐ°Xₐ + Pᵦ°Xᵦ

Conditions:

• ΔH(mixing) < 0

• ΔV(mixing) < 0

Intermolecular forces between unlike molecules are stronger.

Example:

• Chloroform + Acetone

Azeotropes

Azeotropes are liquid mixtures that boil at constant temperature and have the same composition in liquid and vapour phases.

Types of Azeotropes

Minimum Boiling Azeotropes

Formed by solutions showing positive deviation from Raoult’s Law.

Example:

• Ethanol and Water mixture

Maximum Boiling Azeotropes

Formed by solutions showing negative deviation from Raoult’s Law.

Example:

• Nitric Acid and Water mixture

Colligative Properties

Colligative properties depend only on the number of solute particles present in the solution and not on the nature of the solute.

Relative Lowering of Vapor Pressure

Formula:

(P° − P) / P° = X(solute)

Where:

• P° = vapor pressure of pure solvent

• P = vapor pressure of solution

Elevation of Boiling Point

The boiling point of a solution becomes higher than that of the pure solvent.

Formula:

ΔTb = Kb × m

Where:

• ΔTb = elevation in boiling point

• Kb = ebullioscopic constant

• m = molality

Depression of Freezing Point

The freezing point of a solution becomes lower than that of the pure solvent.

Formula:

ΔTf = Kf × m

Where:

• ΔTf = depression in freezing point

• Kf = cryoscopic constant

• m = molality

Osmotic Pressure

Osmotic pressure is the pressure required to stop osmosis through a semipermeable membrane.

Formula:

π = C × R × T

Where:

• π = osmotic pressure

• C = molarity

• R = gas constant

• T = temperature in Kelvin

Solutions having equal osmotic pressure are called isotonic solutions.

Van ’t Hoff Factor (i)

Van ’t Hoff factor corrects colligative property calculations when solutes dissociate or associate in solution.

Formula:

i = Observed Colligative Property / Calculated Colligative Property

Also:

i = Calculated Molar Mass / Observed Molar Mass

For Dissociation

For electrolytes like NaCl and MgCl₂:

i > 1

Formula:

i = 1 + (n − 1)α

Where:

• α = degree of dissociation

• n = number of ions formed

For Association

For association reactions like acetic acid dimerisation:

i < 1

Formula:

i = 1 + [(1/n) − 1]β

Where:

• β = degree of association

• n = number of molecules combining

For Non-Electrolytes

For glucose and urea:

i = 1

Modified Colligative Property Equations

Relative lowering of vapor pressure:

(P° − P) / P° = i × X(solute)

Elevation of boiling point:

ΔTb = i × Kb × m

Depression of freezing point:

ΔTf = i × Kf × m

Osmotic pressure:

π = i × C × R × T

Solutions: Complete Study Resources By PW

Physics Wallah provides multiple study and revision resources for chapter-wise NEET preparation. These resources help improve conceptual understanding, formula revision, and numerical-solving skills.