Solutions: Complete Chapter Notes for Class 12 Physical Chemistry By PW

Solutions can become challenging in NEET preparation when several concentration terms, laws and numerical concepts appear together. A question may require you to distinguish between molarity and molality, apply Raoult’s law, calculate a colligative property or account for abnormal molar mass, making it important to recognise the concept before choosing the formula.

The PW Solutions Complete Chapter One-Shot Revision Video for Class 12 NEET brings the important concepts and numerical applications together for chapter revision. Use these notes alongside the video to revise key formulas, concepts and problem-solving approaches before attempting NEET questions.

Meaning Of A Solution

A solution has two important features:

• It is a homogeneous mixture.

• Its components do not undergo a chemical reaction.

For example, dissolving sugar in water is a physical process. Sugar and water retain their chemical identities. Tea can also be considered a solution because its ingredients are distributed uniformly.

(*Memory Tip: Identify a solution through two points—homogeneous mixture and no chemical reaction.)*

Solute And Solvent

• Solvent: The component that determines the physical state of the solution. If both components have the same physical state, the component present in greater amount is the solvent.

• Solute: The component dissolved in the solvent.

The amount-based rule is used only when physical state cannot distinguish the components.

Types Of Solutions

The physical state of a solution is determined by its solvent.

Questions can be asked directly from the combinations of solute and solvent, so each category should be revised carefully.

Concentration Of A Solution

Concentration expresses the amount of solute present in a given amount of solution or solvent. Important concentration terms include:

• Molarity

• Molality

• Mole fraction

• Percentage concentration

• Parts per million

• Normality

• Formality

• Strength

Concentration units should be understood through their definitions rather than memorised only as formulas. This is important for conversion-based problems.

Molarity

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

Molarity = moles of solute/volume of solution in litres

A 2 M NaCl solution contains 2 moles of NaCl in 1 litre of solution.

Molarity depends on volume. Since volume generally increases with temperature, molarity decreases when temperature increases.

Molality

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

Molality = moles of solute/mass of solvent in kilograms

Unlike molarity, molality does not involve volume. Therefore, it is independent of temperature.

Mole Fraction

The mole fraction of a component is:

Mole fraction = moles of the component / total moles of all components

For a binary solution:

X_A + X_B = 1

(*Memory Tip: Mole fractions represent parts of one complete whole, so the sum of all mole fractions is always 1.)*

Percentage Concentration

For mass percentage, a 20% solution contains 20 g of solute in 100 g of solution.

Parts Per Million

ppm = mass of solute/mass of solution × 10⁶

Similarly:

• ppb uses 10⁹.

• ppt uses 10¹².

These units are used for very dilute solutions.

Mixing And Dilution Of Solutions

Molarity is an intensive property, so molarities cannot be added directly. Instead, calculate the total amount of solute and divide it by the total volume.

For mixing two solutions:

Final molarity = (M₁V₁ + M₂V₂) / (V₁ + V₂)

Volumes must be expressed in the same unit.

A useful form is:

Millimoles = molarity × volume in millilitres

Dilution

Dilution means adding solvent without changing the amount of solute. The volume increases and concentration decreases.

M₁V₁ = M₂V₂

Here, the final volume is V₂. The volume of solvent added is:

Volume added = final volume − initial volume

(*Memory Tip: Do not confuse the final volume with the volume of water added.)*

Henry’s Law And Gas Solubility

The solubility of a gas in a liquid depends on:

• Nature of the gas

• Nature of the solvent

• Temperature

• Pressure

Henry’s law is:

p = K_H × x

where:

• p is the partial pressure of the gas.

• K_H is Henry’s constant.

• x is the mole fraction of the dissolved gas.

At constant pressure:

• Higher K_H means lower gas solubility.

• Lower K_H means higher gas solubility.

Effect Of Pressure And Temperature

• Increasing pressure increases the solubility of a gas.

• Increasing temperature generally decreases gas solubility.

(*Memory Tip: Pressure pushes gas into the liquid, while heat drives gas out of the liquid.)

These principles explain carbonation in cold drinks, decompression sickness in divers, and low oxygen availability at high altitudes.

Vapour Pressure And Raoult’s Law

Vapour pressure is the pressure exerted by vapour above a liquid. It depends on the temperature and nature of the liquid, but not on its volume or number of moles.

Vapour pressure increases with temperature. Liquids with stronger intermolecular forces generally have lower vapour pressure and higher boiling points.

Raoult’s Law For A Binary Solution

For two volatile components A and B:

• Partial pressure of A = P_A° × X_A

• Partial pressure of B = P_B° × X_B

The total vapour pressure is:

P_total = P_A°X_A + P_B°X_B

Here, the mole fractions refer to the liquid phase.

Dalton’s law gives:

P_total = P_A + P_B

For vapour-phase mole fractions:

• P_A = P_total × Y_A

• P_B = P_total × Y_B

Non-Volatile Solute

A non-volatile solute does not contribute to vapour pressure. Therefore:

P_solution = P_solvent° × X_solvent

Since the solvent mole fraction is less than 1, the solution vapour pressure is lower than that of the pure solvent.

Ideal And Non-Ideal Solutions

An ideal solution obeys Raoult’s law at every concentration.

For an ideal solution:

• Solute–solute, solvent–solvent, and solute–solvent interactions are nearly equal.

• Change in enthalpy of mixing is zero.

• Change in volume of mixing is zero.

• Change in entropy of mixing is positive.

• Change in Gibbs free energy of mixing is negative.

Examples include benzene–toluene and n-hexane–n-heptane.

Positive Deviation

Positive deviation occurs when solute–solvent interactions are weaker than the original interactions.

Consequences:

• More molecules escape into the vapour phase.

• Vapour pressure is higher than expected.

• Enthalpy and volume of mixing are positive.

Examples include:

• Ethanol–acetone

• Alcohol–water

• Carbon tetrachloride–toluene

Negative Deviation

Negative deviation occurs when solute–solvent interactions are stronger.

Consequences:

• Fewer molecules escape.

• Vapour pressure is lower than expected.

• Enthalpy and volume of mixing are negative.

Examples include:

• Acetone–chloroform

• Phenol–aniline

• Electrolyte–water systems

(*Memory Tip: Weak A–B forces give positive deviation; strong A–B forces give negative deviation.)

Azeotropes

An azeotrope is a non-ideal solution that boils at a constant composition and cannot be separated by fractional distillation.

• Positive deviation forms a minimum-boiling azeotrope.

• Negative deviation forms a maximum-boiling azeotrope.

Examples:

• Ethanol–water: minimum-boiling azeotrope

• Nitric acid–water: maximum-boiling azeotrope

(*Memory Tip: Positive means minimum boiling; negative means maximum boiling.)

Colligative Properties

Colligative properties depend on the number of solute particles, not their chemical identity. They are:

1. Relative lowering of vapour pressure

2. Elevation in boiling point

3. Depression in freezing point

4. Osmotic pressure

Relative Lowering Of Vapour Pressure

For a non-volatile solute:

Relative lowering of vapour pressure = mole fraction of solute

Elevation In Boiling Point

ΔT_b = K_bm

The solution boils at a higher temperature because its vapour pressure is lowered.

Depression In Freezing Point

ΔT_f = K_fm

The solution freezes at a lower temperature than the pure solvent.

For water:

• K_b = 0.52 K kg mol⁻¹

• K_f = 1.86 K kg mol⁻¹

Osmotic Pressure

Osmosis is the movement of solvent through a semipermeable membrane from a pure solvent or dilute solution toward a more concentrated solution.

Osmotic pressure is the pressure required to stop osmosis.

π = CRT

where C is molarity, R is the gas constant, and T is temperature in kelvin.

• Isotonic solutions: Equal osmotic pressures

• Hypertonic solution: Higher osmotic pressure

• Hypotonic solution: Lower osmotic pressure

In reverse osmosis, external pressure forces solvent in the opposite direction. This process is widely used for water purification.

Van’t Hoff Factor

The van’t Hoff factor corrects colligative-property equations when solute particles associate or dissociate.

Corrected equations include:

• ΔT_b = iK_bm

• ΔT_f = iK_fm

• π = iCRT

Values Of The van’t Hoff Factor

For dissociation:

i = 1 + (n − 1)α

For association:

i = 1 − α + α/n

where n is the number of particles involved, and α is the degree of association or dissociation.

(*Memory Tip: Dissociation increases particles and i; association decreases particles and i.)

Revising Solutions effectively requires more than memorising formulas. Focus on identifying the concentration term, understanding the relevant law and applying the correct relationship to numerical questions. The PW Solutions Complete Chapter One-Shot Revision Video For Class 12 NEET can help you revise the chapter's important concepts and strengthen your preparation before practising questions. 

NEET Resources  You Might Like: