Liquid And Solid State

Q: What are the types of intermolecular forces?

The attractive intermolecular forces are London or dispersion forces between all molecules, dipole-dipole forces between polar molecules, dipole-induced dipole forces between polar and non-polar molecules, and hydrogen bonds. The first three are van der Waals forces. Ion-dipole forces also exist but are not counted as van der Waals forces.

Q: What are London dispersion forces and what affects their strength?

London forces arise when a momentary shift of electrons creates a temporary dipole in one atom that induces a dipole in its neighbour, so the two attract. They act between all molecules, fall off as one over r to the sixth, and grow with polarisability, that is with more electrons, larger size and larger contact area.

Q: What is a hydrogen bond?

A hydrogen bond is a strong dipole-dipole attraction between a hydrogen atom bonded to nitrogen, oxygen or fluorine and a lone pair on nitrogen, oxygen or fluorine of another molecule. Its energy is 10 to 100 kJ per mole, which explains the high boiling point of water and the shapes of proteins and DNA.

Q: What is vapour pressure and how is it related to boiling point?

Vapour pressure is the pressure of a vapour in equilibrium with its liquid at a given temperature. It rises with temperature. A liquid boils when its vapour pressure equals the external pressure; at 1 atm this is the normal boiling point, and at 1 bar the slightly lower standard boiling point.

Q: Why does water boil at a lower temperature on mountains?

At high altitude the atmospheric pressure is lower, so the vapour pressure of water reaches it at a lower temperature. At 0.70 atm water boils near 90 °C, so food cooks slowly. A pressure cooker traps steam, raises the pressure and raises the boiling point above 100 °C.

Q: Why do surface tension and viscosity decrease with temperature?

Heating gives molecules more kinetic energy, which weakens the effect of intermolecular attraction. Surface molecules are pulled inward less effectively, so surface tension falls, and layers slide past one another more easily, so viscosity falls. For water, viscosity drops more than six-fold between 0 and 100 °C.

Q: Are intermolecular forces and liquids part of JEE Advanced 2026?

Yes. The JEE Advanced syllabus under States of Matter: Gases and Liquids lists intermolecular interactions, their types, distance dependence and effect on properties, and liquids with vapour pressure, surface tension and viscosity. The chapter was removed from NCERT Class 11, JEE Main and NEET, so it is examined in JEE Advanced.

Q: What questions on liquids and intermolecular forces come in JEE Advanced?

JEE Advanced asks you to rank boiling points or viscosities from the type and strength of intermolecular forces, apply the distance dependence of different forces, predict effects of temperature on vapour pressure, surface tension and viscosity, and use the Clausius-Clapeyron equation for boiling points at other pressures.

ChemistryStates Of MatterFor JEE aspirants

The liquid and solid state exist because intermolecular forces pull molecules together strongly enough to beat their thermal energy. These forces are London (dispersion), dipole-dipole and dipole-induced dipole forces (together called van der Waals forces), plus hydrogen bonds. Their strength and how fast they fall off with distance explain the vapour pressure, boiling point, surface tension and viscosity of liquids, and the rigidity of solids. Intermolecular forces and the liquid and solid state are JEE Advanced syllabus topics.

On this page1Forces vs thermal energy2Types of forces3London4Dipole forces5Hydrogen bond6Distance dependence7Vapour pressure8Surface tension9Viscosity10Solid state11Examples

Key Formulas - Quick Reference

★ Must learn Distance dependence of interaction energy: London, rotating dipole-dipole and dipole-induced dipole $\propto \frac{1}{r ^{6}}$ ; dipole-dipole in solids (fixed) $\propto \frac{1}{r ^{3}}$ ; ion-ion $\propto \frac{1}{r}$ .
Hydrogen bond energy: 10-100 kJ mol $^{- 1}$ (a covalent bond is roughly 150-570 kJ mol $^{- 1}$ ).
★ Must learn Boiling: vapour pressure = external pressure. Normal b.p. at 1 atm (water $100^{\circ} C$ ); standard b.p. at 1 bar (water $99.6^{\circ} C$ ).
Clausius-Clapeyron: $ln \frac{p _{2}}{p _{1}} = \frac{Δ _{vap} H}{R} (\frac{1}{T _{1}} - \frac{1}{T _{2}})$ .
★ Must learn Surface tension $γ$ : force per unit length, N m $^{- 1}$ (= J m $^{- 2}$ , surface energy); dimensions kg s $^{- 2}$ . Work to enlarge a surface $W = γ Δ A$ .
★ Must learn Viscosity: $F = η A \frac{d u}{d z}$ ; $η$ in Pa s (N s m $^{- 2}$ ); $1 poise = 1 g cm^{- 1} s^{- 1} = 0.1 Pa s$ .
★ Must learn Stronger intermolecular forces: lower vapour pressure, higher b.p., higher $γ$ , higher $η$ , higher $T_{c}$ .
Raising $T$ : vapour pressure rises; surface tension and liquid viscosity fall.

1. What Decides the State of Matter

Single molecules do not boil, wet or flow; these are bulk properties of many molecules together. Two opposing influences decide which state a substance is in:

Intermolecular forces pull molecules together.
Thermal energy, the energy of molecular motion, is directly proportional to temperature and drives molecules apart. The motion itself is called thermal motion.

Figure 1: The state of a substance is a tug of war. Intermolecular forces pull molecules together; thermal energy (

\propto T

) drives them apart. Cooling hands the win to the forces (gas to liquid to solid); heating hands it to thermal motion.

Property	Solid	Liquid	Gas
Shape	definite	takes the container's shape	takes the container's shape
Volume	definite	definite	fills the container
Spacing of particles	closest, fixed	close, random	far apart
Compressibility	negligible	very low	high
Which wins	forces, strongly	forces, just	thermal energy

When the forces are weak, a gas does not liquefy under pressure alone; its thermal energy must first be reduced by cooling. Once cooled enough, it liquefies easily. The chemical composition does not change with state: ice, water and steam are all $H_{2} O$ . Reaction rates, however, do depend on the physical state.

Key idea

State = forces versus thermal energy. Cool or compress and the forces win; heat and thermal motion wins.

2. Types of Intermolecular Forces

Intermolecular forces are the attractive and repulsive forces between neighbouring molecules (or atoms). They do not include ionic bonds between oppositely charged ions or covalent bonds inside a molecule. The attractive ones are called van der Waals forces, after Johannes van der Waals, who used them to explain why real gases are not ideal.

Figure 2: The family of intermolecular forces. All van der Waals forces fall off steeply with distance (

1/ r^{3}

1/ r^{6}

), which is why they act only between near neighbours. Ionic and covalent bonds are not intermolecular forces.

2.1 London (dispersion) forces

Atoms and non-polar molecules have no permanent dipole, yet they still attract. At any instant the electron cloud of one atom may be lopsided, creating a short-lived instantaneous dipole. That dipole distorts the cloud of a neighbour and induces a dipole in it, and the two attract. Fritz London explained this force, so it bears his name.

Figure 3: London (dispersion) forces. A momentary shift of the electron cloud in A creates a temporary dipole that induces one in B; the dipoles keep flipping but always attract. Energy

\propto \frac{1}{r ^{6}}

, effective only within about 500 pm, larger for bigger, more polarisable molecules.

Always attractive; present between all molecules, polar or not.
Energy $\propto \frac{1}{r ^{6}}$ , so it matters only at short range (about 500 pm).
Grows with polarisability: more electrons, larger size and a larger contact surface give stronger London forces. Boiling points rise He $<$ Ne $<$ Ar $<$ Kr $<$ Xe, and n-pentane (309 K) boils above its compact isomer neopentane (283 K).

2.2 Dipole-dipole forces

Polar molecules such as HCl carry permanent partial charges ( $δ +$ and $δ -$ ), each less than the electronic charge $1.6 \times 1 0^{- 19}$ C. Neighbouring molecules line up so that opposite ends attract. This is stronger than London forces between molecules of similar size but weaker than ion-ion attraction, because only partial charges are involved. The energy is $\propto \frac{1}{r ^{3}}$ between stationary dipoles (in solids) and $\propto \frac{1}{r ^{6}}$ between rotating dipoles (in gases and liquids). Polar molecules feel London forces as well, so their total attraction adds up.

2.3 Dipole-induced dipole forces

A polar molecule can induce a dipole in a non-polar neighbour by distorting its electron cloud. The energy is $\propto \frac{1}{r ^{6}}$ and increases with the dipole moment of the polar molecule and the polarisability of the non-polar one; larger molecules are more easily polarised. London forces act here too.

Figure 4: (a) Permanent dipoles line up head to tail; the partial charges are fractions of

1.6 \times 1 0^{- 19}

C. (b) A permanent dipole induces a dipole in a non-polar neighbour. In both cases London forces also act, and the effects add up.

Exam Trick Size can beat polarity. HCl is more polar than HI, yet HI boils higher (238 K vs 188 K), because its much larger electron cloud gives far stronger London forces. Compare polarity only between molecules of similar size.

2.4 Hydrogen bond

A hydrogen bond is a particularly strong dipole-dipole attraction between H bonded to a small, very electronegative atom (N, O or F) and a lone pair on N, O or F of another molecule. Chlorine forms only weak ones. Because only a few elements take part, it is treated as a separate category. Its energy, 10-100 kJ mol $^{- 1}$ , is large enough to control the structure of proteins and nucleic acids and to give water its unusually high boiling point.

Figure 5: A hydrogen bond is a strong dipole-dipole attraction between H on N, O or F and a lone pair on N, O or F of a neighbour. At 10 to 100 kJ mol

^{- 1}

it is much weaker than a covalent bond (about 400-570 kJ mol

^{- 1}

for O-H and H-F) but strong enough to give water its high boiling point and to hold proteins and DNA in shape.

Ion-dipole forces (for example

N a^{+}

surrounded by water molecules) are attractions between an ion and a polar molecule. They are not van der Waals forces.

2.5 Repulsive forces and distance dependence

When molecules are pushed into contact, their electron clouds and their nuclei repel. This repulsion rises very steeply as the distance shrinks, which is why liquids and solids, whose molecules already touch, are so hard to compress. The balance of attraction and repulsion fixes the most stable spacing:

Figure 6: (a) Attraction pulls molecules together until the short-range repulsion takes over; the balance point is the most stable spacing, and the steep wall is why liquids and solids hardly compress. (b) Doubling the distance cuts an ion-ion force to

\frac{1}{2}

but a London force to

\frac{1}{64}

Interaction	Acts between	Energy varies as
Ion-ion (not intermolecular)	ions	$1/ r$
Ion-dipole	ion and polar molecule	$1/ r^{2}$ (fixed dipole)
Dipole-dipole	polar molecules	$1/ r^{3}$ fixed; $1/ r^{6}$ rotating
Dipole-induced dipole	polar and non-polar	$1/ r^{6}$
London (dispersion)	all molecules	$1/ r^{6}$

Key idea

All molecules feel London forces; polar ones add dipole forces; H on N, O, F adds hydrogen bonds. The steeper the

1/ r^{n}

law, the shorter the range.

Quick Recall: tap to check

Which intermolecular forces act between

I_{2}

molecules in solid iodine?

Only London forces (non-polar), but they are strong because

I_{2}

has many electrons, so iodine is a solid at room temperature.

If the distance between two fixed dipoles in a solid doubles, what happens to their interaction energy?

It falls to

\frac{1}{8}

, since energy

\propto 1/ r^{3}

Is the attraction between

N a^{+}

and

H_{2} O

a van der Waals force?

No, it is an ion-dipole force.

3. The Liquid State

In a liquid, intermolecular forces are much stronger than in a gas and there is very little empty space. Molecules stay together, so a liquid has a definite volume, but they slide past one another, so it flows and takes the shape of its container. Liquids are denser than gases. Three properties follow directly from the forces: vapour pressure, surface tension and viscosity.

3.1 Vapour pressure

Partly fill an evacuated, closed container with a liquid. Molecules escape from the surface and the pressure of the vapour rises; soon vapour molecules return to the liquid as fast as others leave. The pressure then stops changing.

Figure 7: In a closed vessel the vapour pressure rises until the rate of condensation equals the rate of evaporation. The constant value is the equilibrium or saturated vapour pressure. It depends on the liquid and the temperature, not on the amount of liquid or the size of the vessel.

Equilibrium (saturated) vapour pressure: the pressure of the vapour in equilibrium with its liquid at a given temperature. Because evaporation depends on temperature, the temperature must always be stated.

In an open vessel, once the vapour pressure equals the external pressure, vapour can form throughout the bulk of the liquid and escape freely: the liquid boils. The boiling temperature at 1 atm is the normal boiling point; at 1 bar it is the standard boiling point, slightly lower because 1 bar is slightly less than 1 atm (water: $100^{\circ} C$ and $99.6^{\circ} C$ ).

Figure 8: Vapour pressure rises steeply with

T

(curves from measured Antoine constants). A liquid boils where its curve meets the external pressure; at 760 mm Hg these are the normal boiling points. At 293 K ether's vapour pressure is 440 mm Hg but water's only 17.5 mm Hg: weaker intermolecular forces, higher vapour pressure, lower boiling point.

At high altitude the air pressure is lower, so water boils below $100^{\circ} C$ and food cooks slowly; a pressure cooker raises the pressure and the boiling point.
Hospital autoclaves sterilise instruments in water boiling above $100^{\circ} C$ , under a weighted vent.
In a closed vessel a liquid never boils. As it is heated the vapour gets denser and the liquid less dense; when the two densities become equal the boundary between them vanishes. That temperature is the critical temperature (see Deviation From Ideal Gas Behaviour).

JEE Advanced Clausius-Clapeyron equation. Vapour pressure rises roughly exponentially with temperature:

ln p = - \frac{Δ _{vap} H}{R T} + C

, so between two temperatures

ln \frac{p _{2}}{p _{1}} = \frac{Δ _{vap} H}{R} (\frac{1}{T _{1}} - \frac{1}{T _{2}})

A plot of

ln p

against

1/ T

is a straight line of slope

- \frac{Δ _{vap} H}{R}

. For water (

Δ_{vap} H = 40.7

kJ mol

^{- 1}

p = 760

mm Hg at 373.15 K), the vapour pressure at 363.15 K comes out as

760 e^{- 0.362} = 529

mm Hg, close to the measured 526 mm Hg.

3.2 Surface tension

A molecule inside a liquid is pulled equally in all directions. A molecule at the surface has neighbours only beside and below it, so it feels a net inward pull and has more energy. A liquid therefore keeps as few molecules as possible at its surface: it minimises its surface area.

Figure 9: Surface molecules have no neighbours above them, so they feel a net inward pull and carry extra energy. The liquid shrinks its surface: drops are spherical, and in a thin tube a wetting liquid climbs (

h = \frac{2 γ cos θ}{r ρ g}

) while mercury is pushed down.

Surface energy: the energy needed to increase the surface area of a liquid by one unit (J m

^{- 2}

). Surface tension ( $γ$ ): the force acting per unit length perpendicular to a line drawn on the surface (N m

^{- 1}

, dimensions kg s

^{- 2}

). The two are numerically equal.

Drops of mercury (and of any liquid in free fall or in space) are spherical, the shape with the least surface for a given volume; on a table, gravity flattens them slightly.
Fire polishing: heating a sharp glass edge lets the softened glass pull into a rounded, smooth shape.
Liquids that wet a surface spread over it as a film; moist soil grains cling together because the water film pulls them together to reduce its area.
A liquid rises (water in glass) or falls (mercury in glass) in a thin capillary because of surface tension.
Stronger attractions give higher surface tension: water (72.8 mN m $^{- 1}$ at $20^{\circ} C$ ) versus ethanol (22.3 mN m $^{- 1}$ ). Heating raises molecular kinetic energy and lowers $γ$ .

3.3 Viscosity

Viscosity is a liquid's resistance to flow, caused by internal friction between layers sliding past one another. When a liquid flows over a fixed surface, the layer touching the surface is at rest and each layer above moves faster. This regular gradation of velocity is laminar flow.

Figure 10: In laminar flow each layer drags on the one above and slows it. The force needed to keep the layers sliding is

F = η A \frac{d u}{d z}

; the coefficient of viscosity

η

is the force per unit area when the velocity gradient is unity.

The force needed to keep the layers moving is proportional to the area of contact $A$ and to the velocity gradient $\frac{d u}{d z}$ :

F = η A \frac{d u}{d z}

The constant $η$ is the coefficient of viscosity: the force per unit area when the velocity gradient is unity. Its SI unit is N s m $^{- 2}$ = Pa s = kg m $^{- 1}$ s $^{- 1}$ ; the cgs unit is the poise (after Jean Poiseuille), with 1 poise = 1 g cm $^{- 1}$ s $^{- 1}$ = 0.1 kg m $^{- 1}$ s $^{- 1}$ . The greater the viscosity, the more slowly the liquid flows. Hydrogen bonds and strong van der Waals forces give high viscosity: glycerol, with three OH groups, is about 1400 times as viscous as water. Glass is so viscous that it behaves as a solid; NCERT calls it an extremely viscous liquid, and the Solid State unit classes it as an amorphous (pseudo) solid.

Figure 11: Heating weakens the effect of intermolecular attraction, so both properties fall (measured data for water). Surface tension drops by about 22% from 0 to

100^{\circ} C

; viscosity drops by a factor of more than 6, because molecules with more kinetic energy slip past one another easily.

Surface tension

A surface property: force per unit length, N m $^{- 1}$ . Caused by the unbalanced inward pull on surface molecules. Falls with temperature.

Viscosity

A bulk property: resistance to flow, Pa s. Caused by attraction between layers sliding past each other. Falls with temperature for liquids.

Gases behave the opposite way: their viscosity rises with temperature, because it comes from molecules carrying momentum between layers, and faster molecules carry more.

Exam Trick One rule covers every liquid property: stronger forces = lower vapour pressure, higher b.p., higher $γ$ , higher $η$ . And heating does the reverse for all of them except vapour pressure, which rises.

Key idea

Vapour pressure, surface tension and viscosity are three views of the same intermolecular attraction.

Quick Recall: tap to check

Does the vapour pressure of water in a sealed flask change if you add more water at the same temperature?

No. Saturated vapour pressure depends only on the liquid and the temperature.

Why do small drops of mercury form spheres?

Surface tension pulls the liquid into the shape with the least surface area for its volume, a sphere.

Arrange by viscosity: ethanol, water, glycerol.

Ethanol

<

water

<

glycerol: more hydrogen bonds per molecule, more resistance to flow.

4. The Solid State

In a solid, intermolecular (or interionic) forces dominate so completely that particles stay in fixed positions and can only vibrate. Solids therefore have a definite shape and volume, are rigid, and are practically incompressible. Heating supplies energy that first raises the temperature and then, at the melting point, breaks the rigid arrangement without changing the temperature:

Figure 12: Heat supplied during melting and boiling breaks intermolecular attractions instead of raising

T

, so the curve is flat. Boiling needs about seven times the heat of melting (40.7 vs 6.01 kJ mol

^{- 1}

), because vaporisation must separate the molecules completely.

Crystalline solids

Long-range regular order; sharp melting point; anisotropic (properties differ with direction). Examples: NaCl, quartz, ice, metals.

Amorphous solids

Only short-range order; soften over a range of temperature; isotropic. Examples: glass, rubber, plastics.

Unit cells, packing, density of crystals, defects and electrical and magnetic properties of solids are covered in detail on the Solid State pages (Class 12 NCERT, also in the JEE Advanced syllabus).

5. Comparing Substances

Figure 13: One judgement, five predictions. Size can beat polarity: HI (238 K) boils above HCl (188 K) because its larger electron cloud gives much stronger London forces.

5.1 The whole concept at a glance

Figure 14: The whole concept on one page. Every liquid property on the right is a consequence of the forces on the left.

6. Solved Examples

Solved Example 1

Which member of each pair has the higher boiling point, and why? (a) n-pentane or neopentane (b)

H_{2} O

H_{2} S

Solution:

(a) n-pentane (309 K vs 283 K): the long chain has more surface contact, so London forces are stronger.

(b) $H_{2} O$ (373 K vs 213 K): water forms hydrogen bonds; $H_{2} S$ does not.

(c) HI (238 K vs 188 K): its larger, more polarisable electron cloud gives stronger London forces, which outweigh the greater polarity of HCl.

Solved Example 2

Name the strongest intermolecular force in (a) liquid argon (b)

C H_{3} OH

(c)

NaCl

dissolved in water (d) HCl gas (e)

I_{2}

dissolved in

CC l_{4}

Solution:

(a) London forces. (b) Hydrogen bonding (O-H). (c) Ion-dipole ( $N a^{+}$ and $C l^{-}$ with water). (d) Dipole-dipole. (e) London forces (both non-polar).

Solved Example 3

At a hill station the air pressure is 0.70 atm. Estimate the boiling point of water there (

Δ_{vap} H = 40.7

kJ mol

^{- 1}

Solution:

Water boils when its vapour pressure is 0.70 atm. With $p_{1} = 1$ atm at $T_{1} = 373.15$ K:

\frac{1}{T _{2}} = \frac{1}{T _{1}} + \frac{R}{Δ _{vap} H} ln \frac{p _{1}}{p _{2}} = \frac{1}{373.15} + \frac{8.314}{40700} ln \frac{1}{0.70} = 2.7527 \times 1 0^{- 3} K^{- 1}

$T_{2} = 363$ K, about $90^{\circ} C$ . Food cooks slowly, hence the pressure cooker.

Solved Example 4

Water (

η = 1.0 \times 1 0^{- 3}

Pa s) flows between two plates, with a contact area of 0.10 m

^{2}

and a velocity gradient of 50 s

^{- 1}

. Find the force needed to maintain the flow.

Solution:

$F = η A \frac{d u}{d z} = (1.0 \times 1 0^{- 3}) (0.10) (50) = 5.0 \times 1 0^{- 3}$ N.

Solved Example 5

Which property of a liquid INCREASES when its temperature is raised?
(A) viscosity
(B) surface tension
(C) vapour pressure
(D) density

Solution:

Answer: (C). More molecules have enough energy to escape, so vapour pressure rises. Viscosity, surface tension and (usually) density fall.

Solved Example 6

The distance between two polar molecules held in fixed positions in a solid is doubled. Their dipole-dipole interaction energy becomes:
(A) 1/2
(B) 1/4
(C) 1/8
(D) 1/64 of the original

Solution:

Answer: (C). For stationary dipoles the energy $\propto \frac{1}{r ^{3}}$ , so doubling $r$ gives $\frac{1}{2 ^{3}} = \frac{1}{8}$ . Option (D) would apply to rotating dipoles or London forces, $\propto \frac{1}{r ^{6}}$ .

Practice Questions

Arrange $F_{2}$ , $C l_{2}$ , $B r_{2}$ , $I_{2}$ in order of increasing boiling point and give the reason.Answer: $F_{2}$ $<$ $C l_{2}$ $<$ $B r_{2}$ $<$ $I_{2}$ (85, 239, 332, 457 K): London forces grow with the number of electrons.
Why is glycerol far more viscous than ethanol?Answer: Glycerol has three OH groups per molecule and forms an extensive hydrogen-bond network; ethanol has one.
Why is the standard boiling point of water ( $99.6^{\circ} C$ ) lower than its normal boiling point?Answer: 1 bar is slightly less than 1 atm, so the vapour pressure reaches the external pressure at a lower temperature.
Convert a viscosity of 1.5 poise into SI units.Answer: 0.15 Pa s.
At $20^{\circ} C$ , which has the higher vapour pressure, diethyl ether or water? Why?Answer: Diethyl ether (about 440 vs 17.5 mm Hg): its molecules cannot hydrogen-bond to each other, so they escape more easily.
How much work is needed to increase the surface area of water by 2.0 m $^{2}$ at $20^{\circ} C$ ( $γ = 72.8$ mN m $^{- 1}$ )?Answer: $W = γ Δ A = 0.0728 \times 2.0 = 0.146$ J.
Why does a liquid heated in a sealed container never boil?Answer: The vapour cannot escape, so the pressure above the liquid rises with its vapour pressure; at the critical temperature the liquid-vapour boundary simply disappears.

Common Mistakes to Avoid

Watch out

Calling ion-dipole forces van der Waals forces. They are a separate, stronger category.
Assuming a polar molecule always has stronger intermolecular forces than a non-polar one. Size matters: $I_{2}$ is a solid, HCl a gas.
Treating a hydrogen bond as a covalent bond. It is an electrostatic attraction of 10-100 kJ mol $^{- 1}$ .
Using $1/ r^{6}$ for dipole-dipole forces in solids. Fixed dipoles follow $1/ r^{3}$ .
Thinking vapour pressure depends on the amount of liquid or the size of the container. It depends only on the liquid and the temperature.
Quoting a boiling point without a pressure. The normal (1 atm) and standard (1 bar) boiling points differ.
Writing surface tension in N m $^{- 2}$ . It is N m $^{- 1}$ (numerically equal to surface energy in J m $^{- 2}$ ).
Saying all viscosities rise with temperature. Liquid viscosity falls; gas viscosity rises.

Frequently Asked Questions

What are the types of intermolecular forces?

The attractive intermolecular forces are London or dispersion forces between all molecules, dipole-dipole forces between polar molecules, dipole-induced dipole forces between polar and non-polar molecules, and hydrogen bonds. The first three are van der Waals forces. Ion-dipole forces also exist but are not counted as van der Waals forces.

What are London dispersion forces and what affects their strength?

London forces arise when a momentary shift of electrons creates a temporary dipole in one atom that induces a dipole in its neighbour, so the two attract. They act between all molecules, fall off as one over r to the sixth, and grow with polarisability, that is with more electrons, larger size and larger contact area.

What is a hydrogen bond?

A hydrogen bond is a strong dipole-dipole attraction between a hydrogen atom bonded to nitrogen, oxygen or fluorine and a lone pair on nitrogen, oxygen or fluorine of another molecule. Its energy is 10 to 100 kJ per mole, which explains the high boiling point of water and the shapes of proteins and DNA.

What is vapour pressure and how is it related to boiling point?

Vapour pressure is the pressure of a vapour in equilibrium with its liquid at a given temperature. It rises with temperature. A liquid boils when its vapour pressure equals the external pressure; at 1 atm this is the normal boiling point, and at 1 bar the slightly lower standard boiling point.

Why does water boil at a lower temperature on mountains?

At high altitude the atmospheric pressure is lower, so the vapour pressure of water reaches it at a lower temperature. At 0.70 atm water boils near 90 °C, so food cooks slowly. A pressure cooker traps steam, raises the pressure and raises the boiling point above 100 °C.

Why do surface tension and viscosity decrease with temperature?

Heating gives molecules more kinetic energy, which weakens the effect of intermolecular attraction. Surface molecules are pulled inward less effectively, so surface tension falls, and layers slide past one another more easily, so viscosity falls. For water, viscosity drops more than six-fold between 0 and 100 °C.

Are intermolecular forces and liquids part of JEE Advanced 2026?

Yes. The JEE Advanced syllabus under States of Matter: Gases and Liquids lists intermolecular interactions, their types, distance dependence and effect on properties, and liquids with vapour pressure, surface tension and viscosity. The chapter was removed from NCERT Class 11, JEE Main and NEET, so it is examined in JEE Advanced.

What questions on liquids and intermolecular forces come in JEE Advanced?

JEE Advanced asks you to rank boiling points or viscosities from the type and strength of intermolecular forces, apply the distance dependence of different forces, predict effects of temperature on vapour pressure, surface tension and viscosity, and use the Clausius-Clapeyron equation for boiling points at other pressures.

Previous year questions on Liquid And Solid State

1 question from past papers, each with a step-by-step solution.

JEE Main 2025 Apr 3 Shift 2, Chemistry Q16

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