Chemistry Solutions For NEET 2026: Concepts, Formulas and Tricks

Solutions is an important Physical Chemistry chapter for NEET 2026. It is short and scoring, and it connects closely with Mole Concept, Thermodynamics, Equilibrium, and Electrochemistry. Many questions are direct and formula-based, so a quick revision of the core ideas can save valuable time in the exam hall.

Questions in this chapter often test definitions, concentration terms, temperature dependence, stoichiometry, and solubility. They also cover vapour pressure, boiling-point elevation, freezing-point depression, osmotic pressure, and abnormal molar mass. A clear understanding of solvent-solute notation and particle changes helps you solve both numerical and conceptual questions quickly and accurately.

What Is A Solution And What Are Its Components?

A solution is a homogeneous mixture of two or more components.

• Solute: Component present in the smaller quantity.

• Solvent: Component present in the larger quantity. It generally determines the physical state of the solution.

If the solvent is solid, the solution is a solid solution. If the solvent is liquid, it is a liquid solution.

For calculations, the solvent is treated as component 1 and the solute as component 2. Therefore:

• Moles of solvent = n1

• Moles of solute = n2

Types Of Solutions Based On Physical State

Solutions are classified by the physical states of the solute and the solvent. Since the solvent determines the physical state of the solution, identifying it first makes the type easy to name. The table below lists common combinations with examples.

An amalgam is a solution involving mercury and another metal. Solute and solvent must be identified using quantity and the physical-state criterion, not merely by the presence of mercury.

Mole Fraction

The mole fraction of the solute is:

x2 = n2 / (n1 + n2)

The mole fraction of the solvent is:

x1 = n1 / (n1 + n2)

Their sum is:

x1 + x2 = 1

For an aqueous NaOH solution containing 44 g solution and 8 g NaOH:

• Mass of water = 44 − 8 = 36 g

• Moles of NaOH = 8 / 40 = 0.2 mol

• Moles of water = 36 / 18 = 2 mol

Therefore:

xNaOH = 0.2 / (0.2 + 2) = 1 / 11

Concentration Terms

Concentration terms describe how much solute is present in a given amount of solution or solvent. They are grouped by whether they depend on volume, which makes them temperature-dependent, or on mass, which makes them temperature-independent.

Molality

Molality is represented by a small m. It is the number of moles of solute per kilogram of solvent:

m = moles of solute/mass of solvent in kg

Using masses in grams:

m = (w2 × 1000) / (M2 × w1)

Unit: mol kg−1 or molal.

Molality is based on solvent mass, so it is temperature-independent.

Molarity

Molarity is represented by a capital M. It is the number of moles of solute per litre of solution:

M = moles of solute/volume of solution in litres

If volume is in millilitres:

M = (w2 / M2) × (1000 / V)

Number of moles:

n2 = M × volume in litres

Number of millimoles:

millimoles = M × volume in millilitres

Strength

Strength is the mass of solute present per volume of solution:

Strength = mass of solute/volume of solution

For mass in grams and volume in litres:

Strength = w2 / V

Unit: g L−1.

Normality

Normality is represented by a capital N. It is the number of gram equivalents of solute per litre of solution:

N = gram equivalents/volume in litres

The important relation is:

N = M × n-factor

Equivalent mass:

Equivalent mass = molar mass / n-factor

Number of equivalents:

Equivalents = moles × n-factor

Formality

For ionic solids, formality is used in place of molarity in the discussed context.

For:

AlCl3 → Al3+ + 3Cl−

A formal concentration of 0.1 gives:

• Al3+ concentration = 0.1

• Cl− concentration = 0.3

Useful Relations

Molarity from mass percentage and density:

M = (10 × percentage × density) / molar mass of solute

Relation between molarity and molality:

m = (1000M) / (1000d − M M2)

Dilution And Mixing Of Solutions

During dilution or concentration, the number of solute moles remains constant:

M1V1 = M2V2

A molarity–volume graph is a rectangular hyperbola because M × V is constant.

When two solutions are mixed:

Mmixture = (M1V1 + M2V2) / total volume

For concentration terms:

n-Factor and Acid–Base Mixing

The n-factor is always positive.

• For an atom: n-factor = valency

• For an ion: n-factor = absolute value of charge

• For an acid: n-factor = number of H+ ions donated

• For oxidation: n-factor = electrons lost or increase in oxidation number

• For reduction: n-factor = electrons gained or decrease in oxidation number

For mixing acid and base solutions:

Acid equivalents = NaVa

Base equivalents = NbVb

Equal Equivalents

If NaVa = NbVb, complete neutralisation occurs. For a strong acid and strong base:

• The solution is neutral.

• pH = 7

• No salt hydrolysis occurs.

Acid In Excess

If NaVa > NbVb:

[H+] = (NaVa − NbVb) / total volume

The mixture is acidic and pH is less than 7.

Base In Excess

If NbVb > NaVa:

[OH−] = (NbVb − NaVa) / total volume

The mixture is basic.

Solubility And Henry’s Law

Solubility is the maximum amount of solute that dissolves in a solvent. It depends on:

• Nature of solute and solvent

• Temperature

• Pressure

For solids in liquids, pressure has little effect. The principle is:

Like dissolves like

Polar solutes dissolve better in polar solvents, while nonpolar solutes dissolve better in nonpolar solvents.

For gases in liquids:

• Gas solubility decreases as temperature increases.

• Gas solubility increases as pressure increases.

Henry’s law is:

P = KH × X

where P is gas pressure, KH is Henry’s law constant, and X is the mole fraction of dissolved gas.

Solubility is inversely related to KH:

• Higher KH means lower solubility.

• Lower KH means higher solubility.

In a very dilute solution:

Xgas is approximately equal to moles of gas / moles of liquid.

Applications include carbonated drinks, diving equipment, and high-altitude effects. Helium is used in diving mixtures to reduce nitrogen-related problems. Reduced oxygen pressure at high altitudes can cause anoxia.

Raoult’s Law And Vapour Pressure

For a volatile component:

PA = PA° × XA

For two volatile liquids:

Total pressure = PA°XA + PB°XB

A liquid boils when its vapour pressure becomes equal to external pressure.

Adding a non-volatile solute lowers the vapour pressure of the solvent. This causes:

• Elevation in boiling point

• Depression in freezing point

For vapour composition:

YA = PA / total pressure

Ideal And Non-Ideal Solutions

An ideal solution obeys Raoult’s law throughout the composition range.

Properties of ideal solutions:

• Volume change on mixing is zero.

• Heat change on mixing is zero.

• Forces between unlike molecules are similar to forces between like molecules.

Examples include:

• Hexane and heptane

• Benzene and toluene

• Chloroethane and bromoethane

Deviations From Raoult’s Law

Non-ideal solutions do not follow Raoult’s law over the full composition range. They deviate in one of two ways, depending on whether the forces between unlike molecules are weaker or stronger than those between like molecules. The table below compares the two cases.

Positive deviation forms a minimum-boiling azeotrope. Negative deviation forms a maximum-boiling azeotrope.

An azeotrope has the same liquid and vapour composition and cannot be separated by ordinary distillation.

Colligative Properties

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

The four properties are:

1. Relative lowering of vapour pressure

2. Elevation in boiling point

3. Depression in freezing point

4. Osmotic pressure

Elevation In Boiling Point

ΔTb = Kb × m

For electrolytes:

ΔTb = iKb × m

Depression In Freezing Point

ΔTf = Kf × m

For electrolytes:

ΔTf = iKf × m

The solution freezing point is:

Freezing point of solution = freezing point of pure solvent − ΔTf

Osmotic Pressure

Osmotic pressure is the pressure required to stop osmosis:

π = CRT

For electrolytes:

π = iCRT

Osmotic pressure is preferred for finding the molar mass of biomolecules because it can be measured at room temperature.

Van’t Hoff Factor

The Van’t Hoff factor accounts for dissociation or association.

• Non-electrolyte: i = 1

• Dissociation: i > 1

• Association: i < 1

For dissociation into n particles:

i = 1 + (n − 1)α

For association of n particles into one:

i = 1 + (1/n − 1)α

It can also be calculated as:

i = normal molar mass / abnormal molar mass

For benzoic acid:

• In water, it dissociates and i > 1.

• In benzene, it forms dimers and i < 1.

Common Mistakes To Avoid In Solutions Numericals

Most errors in this chapter come from small slips in units, signs, and trends rather than from the concepts themselves. Keep an eye on these points while solving numericals:

• Molarity vs molality: Molarity uses litres of solution, while molality uses kilograms of solvent. Check the units before substituting.

• Formality and n-factor: A 0.1 formal AlCl₃ solution gives 0.1 Al³⁺ , but 0.3 Cl⁻, and the n-factor is always positive.

• Deviations from Raoult’s law: Positive deviation gives a lower boiling point, and negative deviation gives a higher one.

• Boiling and freezing points: Add ΔTb to the pure solvent’s boiling point, but subtract ΔTf from its freezing point.

• Van’t Hoff factor: Include i for electrolytes, and enter α as a fraction, so 30% becomes 0.3.

Solutions combines a small set of concentration formulas with clear rules on solubility, vapour pressure, and colligative properties. Revise these regularly, practise PYQs, and you can solve most numerical and conceptual questions in this chapter quickly.

Related Sources

• NEET Exam 

• NEET Admit Card

• NEET Application Form

• NEET Results

• NEET Syllabus