Solutions Formula Sheet — NEET Chemistry

Q: What are the most important Solutions formulas for NEET?

This Solutions formula sheet covers all the high-yield Chemistry formulas, definitions and facts you need for NEET, across Methods of expressing concentration — molarity, molality, mole fraction, Vapour pressure and Raoult law, Ideal and non-ideal solutions, Colligative properties and determination of molecular mass, Van't Hoff factor and abnormal molar mass — each shown with the key result and, where useful, a worked example.

Q: Is this Solutions formula sheet free?

Yes — the full chapter formula sheet is free to read online, no login or payment required.

Q: How should I revise Solutions for NEET?

Blurt the Solutions key points from memory, then check against this sheet to find your gaps — and practise a few previous-year questions on the chapter to make sure you can apply them under time pressure.

Every key Solutions formula, definition and fact for NEET Chemistry in one place — with common examiner traps and worked examples. Free to read; blurt from memory, then check your gaps.

Syllabus — topics coveredNCERT · 5 sub-topics

Methods of expressing concentration — molarity, molality, mole fraction
Vapour pressure and Raoult law
Ideal and non-ideal solutions
Colligative properties and determination of molecular mass
Van't Hoff factor and abnormal molar mass

Expressing Concentration

A solution is a homogeneous mixture

▸ $Solvent$ = component in larger amount; $solute$ = in smaller amount.
▸ $Binary$ solution = two components; can be solid/liquid/gas in solid/liquid/gas (9 types).
▸Concentration = how much solute per fixed amount of solution or solvent.

Mass %

\frac{w _{solute}}{w _{solution}} \times 100

ppm

\frac{w _{solute}}{w _{solution}} \times 1 0^{6}

Mole fraction

x_{A} = \frac{n _{A}}{n _{A} + n _{B}}, \sum x = 1

Molarity

M = \frac{n _{solute}}{V _{solution} ( L )}

Molality

m = \frac{n _{solute}}{w _{solvent} ( kg )}

🎯 Exam · Temperature dependence

Molarity

changes with temperature (volume expands).

Molality

mole fraction

and

mass %

are temperature-

independent

(based on mass). Prefer molality for colligative-property work.

💡 Tip · Quick link

For dilute aqueous solutions, molarity

\approx

molality (since density

\approx 1

L^{- 1}

). 1 litre of water

= 55.5

mol.

🚫 Examiner Trap · Examiner traps

(1)

Molarity

is temperature-dependent (volume expands);

molality

, mole fraction, mass% are NOT — use molality for colligative work. (2) Molality uses kg of

solvent

, molarity uses L of

solution

. (3) Dilute aqueous:

M \approx m

(density

\approx 1

). (4) ppm

=

mass-ratio

\times 1 0^{6}

, for very dilute solutions.

Solubility & Henry's Law

Solubility

▸ $Solid in liquid$ : a saturated solution is in dynamic equilibrium with undissolved solute; mostly increases with temperature for endothermic dissolution.
▸ $Gas in liquid$ : solubility rises with $pressure$ and falls with $temperature$ (dissolution is exothermic — Le Chatelier).
▸Polar solutes dissolve in polar solvents ('like dissolves like').

Henry's law: At constant temperature, the partial pressure of a gas over a solution is proportional to its mole fraction in the solution:

p = K_{H} x

. A larger Henry's constant

K_{H}

means a less soluble gas.

Straight-line plot of gas mole fraction versus partial pressure with slope one over K_H, and the consequences that pressure raises and temperature lowers gas solubility — $p = K_{H} x$ : dissolved gas $\propto$ its partial pressure.

Henry's law in action

▸ $Soda water / soft drinks$ : bottled under high $C O_{2}$ pressure to force gas into solution.
▸ $Scuba diving (bends)$ : high pressure dissolves more $N_{2}$ in blood; rapid ascent releases bubbles — tanks use diluted air (He).
▸ $Anoxia$ : low $O_{2}$ partial pressure at high altitude lowers blood $O_{2}$ .

★ Remember · Raoult is a special case of Henry

Both say partial pressure

\propto

mole fraction. When the solute is the solvent itself,

K_{H} \to p^{0}

(pure VP), and Henry's law becomes Raoult's law.

✎ Example · Why warm soda goes flat

Explain using Henry's law.

Opening the bottle drops the $C O_{2}$ partial pressure above the liquid.
Lower $p \Rightarrow$ lower $x_{C O_{2}}$ , so gas escapes.
Warming further lowers gas solubility.

∴ Lower p and higher T both reduce dissolved

C O_{2}

🚫 Examiner Trap · Examiner traps

(1) Gas solubility

↑

with PRESSURE,

↓

with TEMPERATURE (dissolution exothermic). (2) Henry

p = K_{H} x

: a

larger

K_{H}

means a

less

soluble gas. (3) Raoult is the special case of Henry when

K_{H} \to p^{0}

. (4) Soda fizzes on opening (lower

p_{C O_{2}}

); divers get bends (more

N_{2}

dissolves at depth).

Raoult's Law & Vapour Pressure

Raoult's law: For a solution of volatile liquids, the partial vapour pressure of each component is proportional to its mole fraction:

p_{1} = p_{1}^{0} x_{1}

p_{2} = p_{2}^{0} x_{2}

, where

p^{0}

is the pure-component vapour pressure.

Total vapour pressure (Dalton + Raoult)

p_{t o t a l} = p_{1}^{0} x_{1} + p_{2}^{0} x_{2} = p_{1}^{0} + (p_{2}^{0} - p_{1}^{0}) x_{2}

Varies linearly with

x_{2}

between

p_{1}^{0}

and

p_{2}^{0}

Vapour pressure versus mole fraction for an ideal solution: two dashed partial-pressure lines and a solid total-pressure line varying linearly between the pure vapour pressures — Ideal solution: straight-line $p_{t o t a l}$ between $p_{1}^{0}$ and $p_{2}^{0}$ .

Vapour-phase composition

y_{i} = \frac{p _{i}}{p _{t o t a l}}

y_{i}

= mole fraction of component i in the vapour. The vapour is always

richer

in the more volatile component (higher

p^{0}

💡 Tip · Non-volatile solute

If the solute is non-volatile, only the solvent contributes to vapour pressure:

p_{solution} = p_{1}^{0} x_{1} < p_{1}^{0}

— the basis of all colligative properties.

✎ Example · Vapour-phase enrichment

CHC l_{3}

(

p^{0} = 200

) and

C H_{2} C l_{2}

(

p^{0} = 415

mm Hg), with

x_{C H_{2} C l_{2}} = 0.688

. Find

y_{C H_{2} C l_{2}}

$p_{C H_{2} C l_{2}} = 0.688 \times 415 = 285.5$ ; $p_{CHC l_{3}} = 0.312 \times 200 = 62.4$ .
$p_{t o t a l} = 347.9$ mm Hg.
$y = 285.5/347.9$ .

∴

y_{C H_{2} C l_{2}} = 0.82

(vapour richer in the more volatile component).

🚫 Examiner Trap · Examiner traps

(1) The vapour is always

richer in the more volatile

component (higher

p^{0}

). (2) Non-volatile solute: only solvent contributes,

p = p_{1}^{0} x_{1} < p_{1}^{0}

(basis of colligative properties). (3)

p_{t o t a l}

varies

linearly

with

x_{2}

only for an ideal solution. (4) Use mole fractions in the LIQUID for Raoult, in the VAPOUR (

y_{i}

) for composition.

Ideal & Non-ideal Solutions

Ideal solution: Obeys Raoult's law over the entire composition range, with

Δ_{mi x} H = 0

and

Δ_{mi x} V = 0

. Occurs when solute-solvent (A–B) interactions equal A–A and B–B (e.g. benzene + toluene, n-hexane + n-heptane).

Two vapour-pressure plots: positive deviation curving above the Raoult line, and negative deviation curving below it — Real solutions deviate when A–B interactions differ from A–A / B–B.

	Positive dev.	Negative dev.
A–B force	weaker	stronger
Vapour pressure	$>$ Raoult	$<$ Raoult
$Δ_{mi x} H$	$+$ (endo)	$-$ (exo)
$Δ_{mi x} V$	$+$	$-$
Azeotrope	min. boiling	max. boiling

Recognising the deviations

▸ $Positive$ : A–B attraction weaker $\Rightarrow$ molecules escape easily $\Rightarrow$ VP higher. Ex: ethanol + acetone, $C S_{2}$ + acetone.
▸ $Negative$ : A–B attraction stronger (often new H-bonds) $\Rightarrow$ VP lower. Ex: phenol + aniline, chloroform + acetone, $HN O_{3}$ + water.

Azeotrope: A binary mixture that boils at constant temperature with the same composition in liquid and vapour, so it

cannot

be separated by fractional distillation. Large positive deviation → minimum-boiling (e.g. 95% ethanol-water); large negative deviation → maximum-boiling (e.g. 68%

HN O_{3}

-water).

★ Remember · Sign logic

Positive deviation

\Rightarrow

weaker forces

\Rightarrow

Δ_{mi x} H > 0

Δ_{mi x} V > 0

. Negative deviation

\Rightarrow

stronger forces

\Rightarrow

Δ_{mi x} H < 0

Δ_{mi x} V < 0

. Ideal

\Rightarrow

both zero.

🚫 Examiner Trap · Examiner traps

(1)

Positive

deviation: A–B weaker

\Rightarrow

>

Raoult,

Δ_{mi x} H > 0

, min-boiling azeotrope (ethanol+water).

Negative

: A–B stronger, VP

<

Raoult,

Δ_{mi x} H < 0

, max-boiling azeotrope (HN

O_{3}

+water). (2) Ideal

\Rightarrow Δ_{mi x} H = Δ_{mi x} V = 0

. (3) An azeotrope CANNOT be separated by fractional distillation. (4) Don't confuse min-boiling (positive) with max-boiling (negative).

Colligative Properties — VP, Boiling & Freezing

Colligative property: A property that depends only on the number of solute particles, not their nature. The four are: relative lowering of vapour pressure, elevation of boiling point, depression of freezing point, and osmotic pressure.

Relative lowering of VP

\frac{p _{1}^{0} - p _{1}}{p _{1}^{0}} = x_{2}

Boiling-point elevation

Δ T_{b} = K_{b} m

Freezing-point depression

Δ T_{f} = K_{f} m

Two plots: solution vapour-pressure curve below the solvent curve, raising the boiling point at 1 atm and lowering the freezing point where it meets the solid curve — A non-volatile solute lowers VP, so $T_{b}$ rises and $T_{f}$ falls.

Molar mass from colligative data

M_{2} = \frac{1000 w _{2} K _{b}}{Δ T _{b} w _{1}} = \frac{1000 w _{2} K _{f}}{Δ T _{f} w _{1}}

w_{2}

= solute mass,

w_{1}

= solvent mass (g),

m = \frac{1000 w _{2}}{M _{2} w _{1}}

. Also

M_{2} = \frac{w _{2} M _{1}}{( Δ p / p _{1}^{0} ) w _{1}}

from VP lowering.

Solvent	$K_{b}$	$K_{f}$ (K kg mo $l^{- 1}$ )
Water	0.52	1.86
Benzene	2.53	5.12
Chloroform	3.63	4.79
Camphor	5.95	40.0

🎯 Exam · Why salt/glycol are used

K_{b} = \frac{R M _{1} T _{b}^{2}}{1000 Δ _{v a p} H}

K_{f} = \frac{R M _{1} T _{f}^{2}}{1000 Δ _{f u s} H}

are properties of the

solvent

only. Salt melts ice (fp depression); ethylene glycol is antifreeze + coolant (raises

T_{b}

, lowers

T_{f}

🚫 Examiner Trap · Examiner traps

(1) Colligative properties depend on the

number

of solute particles, NOT their nature. (2)

K_{b}, K_{f}

are properties of the

solvent

only (camphor has a huge

K_{f} = 40

). (3) Use

molality

Δ T = K m

(not molarity). (4) A non-volatile solute raises

T_{b}

AND lowers

T_{f}

— opposite directions.

Osmotic Pressure

Osmosis & osmotic pressure (

Π

): Across a semipermeable membrane, solvent flows from pure solvent (or dilute) to a more concentrated solution — osmosis. The excess pressure that must be applied on the solution to just stop this flow is the osmotic pressure

Π

U-tube with a semipermeable membrane: solvent flows into the solution raising its level by height h, with osmotic-pressure equations and tonicity definitions — $Π = h ρ g$ ; it is the best colligative property for large molecules.

van't Hoff equation for osmotic pressure

Π = CRT = \frac{n _{2}}{V} RT \Rightarrow M_{2} = \frac{w _{2} RT}{Π V}

C = molarity, T in K,

R = 0.0821

L atm

K^{- 1}

l^{- 1}

. Measured at room temperature with large values even for dilute solutions.

Tonicity (same

T

)

▸ $Isotonic$ : equal $Π$ — no net osmosis (0.9% $NaCl$ saline is isotonic with blood).
▸ $Hypertonic$ : higher $Π$ — cell loses water and $shrinks$ (crenation).
▸ $Hypotonic$ : lower $Π$ — cell takes in water and $swells$ (may burst).

🎯 Exam · Reverse osmosis & uses

Applying

p > Π

on the solution pushes pure solvent

out

—

reverse osmosis

, used for sea-water desalination. Osmotic pressure is preferred for

macromolecules/proteins

(large, easily measured, room-T).

✎ Example · Molar mass of a protein

200 cm^{3}

solution has

1.26

g protein,

Π = 2.57 \times 1 0^{- 3}

bar at 300 K.

$M_{2} = \frac{w _{2} RT}{Π V}$ with $R = 0.083$ L bar $K^{- 1}$ mo $l^{- 1}$ , $V = 0.200$ L.
$M_{2} = \frac{1.26 \times 0.083 \times 300}{2.57 \times 1 0 ^{- 3} \times 0.200}$ .

∴

M_{2} \approx 61000 g mol^{- 1}

🚫 Examiner Trap · Examiner traps

(1)

Π = CRT

uses

molarity

C and T in kelvin. (2) Osmotic pressure is the BEST colligative property for macromolecules/proteins (large, measurable at room T). (3) Hypertonic

\Rightarrow

cell shrinks; hypotonic

\Rightarrow

swells/bursts; isotonic

\Rightarrow

no net flow. (4) Reverse osmosis needs

p > Π

(desalination).

van't Hoff Factor & Abnormal Molar Mass

Abnormal molar mass: When a solute dissociates or associates in solution, the number of particles changes, so the measured colligative property — and hence the calculated molar mass — differs from the expected value.

KCl dissociating into two ions giving i greater than 1, versus two acetic acid molecules associating into a dimer giving i less than 1 — Dissociation raises the particle count ( $i > 1$ ); association lowers it ( $i < 1$ ).

van't Hoff factor

i = \frac{normal molar mass}{abnormal molar mass} = \frac{observed colligative property}{calculated value} = \frac{particles after}{particles before}

i > 1

: dissociation (e.g.

KCl \to 2

K_{2} S O_{4} \to 3

i < 1

: association (e.g. acetic acid dimer in benzene,

i \to 0.5

i = 1

: no change.

Relative lowering of VP

\frac{p _{1}^{0} - p _{1}}{p _{1}^{0}} = i \frac{n _{2}}{n _{1}}

Boiling-point elevation

Δ T_{b} = i K_{b} m

Freezing-point depression

Δ T_{f} = i K_{f} m

Osmotic pressure

Π = i CRT

★ Remember · Degree of dissociation / association

Dissociation into n ions:

i = 1 + (n - 1) α

. Association of n monomers:

i = 1 - (1 - \frac{1}{n}) α

, where

α

is the degree of dissociation/association.

🎯 Exam · Particle count drives everything

For the same molality,

Δ T_{f}

order:

K_{2} S O_{4} (i \approx 3) > NaCl (i \approx 2) > glucose (i = 1)

. Electrolytes depress freezing point more than non-electrolytes.

🚫 Examiner Trap · Examiner traps

(1)

i > 1

dissociation (KCl

\to 2

K_{2} S O_{4} \to 3

);

i < 1

association (acetic acid dimer in benzene

i \to 0.5

). (2) Dissociation:

i = 1 + (n - 1) α

; association:

i = 1 - (1 - \frac{1}{n}) α

. (3) Multiply EVERY colligative formula by i for electrolytes. (4) Same molality:

Δ T_{f}

order

K_{2} S O_{4} >

NaCl

>

glucose (particle count).

More NEET Chemistry formula sheets

Frequently Asked Questions

What are the most important Solutions formulas for NEET?

This Solutions formula sheet covers all the high-yield Chemistry formulas, definitions and facts you need for NEET, across Methods of expressing concentration — molarity, molality, mole fraction, Vapour pressure and Raoult law, Ideal and non-ideal solutions, Colligative properties and determination of molecular mass, Van't Hoff factor and abnormal molar mass — each shown with the key result and, where useful, a worked example.

Is this Solutions formula sheet free?

Yes — the full chapter formula sheet is free to read online, no login or payment required.

How should I revise Solutions for NEET?

Blurt the Solutions key points from memory, then check against this sheet to find your gaps — and practise a few previous-year questions on the chapter to make sure you can apply them under time pressure.

Also useful: all formula sheets · NEET previous-year papers · most important chapters.