Liquid Properties and Hydrogen Bonding

Intermolecular Forces

The Kinetic Molecular Theory explains that molecules in a liquid are in constant motion but are held together by intermolecular forces — attractive forces between neighbouring molecules that are much weaker than covalent bonds within a molecule. These forces determine whether a substance is a gas, liquid, or solid at a given temperature, and they control properties like boiling point, vapour pressure, and viscosity.
Four types of intermolecular forces exist: dipole-dipole forces, ion-dipole forces, dipole-induced dipole forces, and London dispersion forces (instantaneous dipole-induced dipole forces).
Strength comparison: covalent bonds >> hydrogen bonds > dipole-dipole > London dispersion forces.
Physical origin: Intermolecular forces arise from electrostatic attractions between partial charges on molecules — they have nothing to do with valence electrons, unlike covalent or ionic bonding.

Types of Intermolecular Forces (Weakest to Strongest)

1
London dispersion forces — present in all molecules
2
Dipole-induced dipole forces (Debye forces)
3
Dipole-dipole forces — polar molecules only
4
Ion-dipole forces — ions + polar molecules
5
Hydrogen bonding — strongest intermolecular force
In polar molecules like HCl, the difference in electronegativity between bonded atoms creates partial positive and partial negative charges. The positive end of one molecule attracts the negative end of another, producing dipole-dipole forces. Thermal energy prevents perfect alignment, but a net attraction persists. These forces are roughly one percent as strong as a covalent bond.
Strength depends on two factors: electronegativity difference between bonded atoms, and the distance between molecules.
In gases these forces are very weak because intermolecular distances are large; in liquids they are reasonably strong.
Greater dipole-dipole strength means higher melting points, boiling points, heats of vapourisation, and heats of sublimation.

Molecules Showing Dipole-Dipole Attractions

•
HCl — Cl(δ⁻) attracts H(δ⁺) of neighbouring molecule
•
CHCl₃ (chloroform) — three Cl(δ⁻) attract H(δ⁺) of adjacent molecule
When a polar molecule approaches a non-polar molecule, the partial charge of the polar molecule distorts the electron cloud of the non-polar one, inducing a temporary dipole. The resulting attraction is called dipole-induced dipole force (Debye force). These forces are weaker than dipole-dipole forces but important in mixtures of polar and non-polar substances.
Polar molecule acts as the source: its permanent partial charge creates an electric field that shifts the electron cloud of the nearby non-polar molecule.
Induced dipole is temporary: it exists only while the polar molecule is nearby.
Strength depends on polarizability of the non-polar molecule — larger electron clouds are more easily distorted.
London dispersion forces are the weakest intermolecular forces, yet they are present in ALL molecules — polar and non-polar alike. In a non-polar molecule like helium, the electrons are in constant motion. At any instant, the electron density may be uneven, creating an instantaneous dipole. This momentary dipole distorts the electron cloud of a neighbouring atom, inducing a second dipole. The fleeting attraction between them is the London force.
Very short-lived: dipoles vanish as quickly as they form, but new ones constantly appear in different orientations.
Dominant force for non-polar molecules like Cl₂, H₂, and noble gases (He, Ne, Ar).
Also present in polar molecules but usually overshadowed by dipole-dipole and hydrogen bonding.
Discovered by Fritz London in 1930 — sometimes called London forces or dispersion forces.
The strength of London dispersion forces depends on two key factors: the size of the electron cloud (and thus the polarizability) and the number of atoms in the molecule. Both factors explain why boiling points increase down groups in the periodic table and why larger hydrocarbons have much higher boiling points than smaller ones.
Polarizability increases with atomic/molecular size: larger electron clouds are more easily distorted, creating stronger London forces.
Noble gases (Group VIIIA): boiling points increase from He (−268.6 °C) to Rn (−61.8 °C) because outer electrons are farther from the nucleus and more easily polarized.
Halogens (Group VIIA): fluorine is a gas (B.P. −188.1 °C) while iodine is a solid (B.P. 184.4 °C) — the same trend from greater polarizability.
Number of atoms matters: C₂H₆ boils at −88.6 °C but C₆H₁₄ boils at 68.7 °C because longer chains provide more surface area for London forces to act along.

Boiling Points of Halogens and Noble Gases

•
He: −268.6 °C | F: −188.1 °C
•
Ne: −245.9 °C | Cl: −34.6 °C
•
Ar: −185.7 °C | Br: 58.8 °C
•
Kr: −152.3 °C | I: 184.4 °C
•
Xe: −107.1 °C | At: 337 °C
•
Rn: −61.8 °C

Evaporation

According to the Kinetic Molecular Theory, molecules in a liquid have a range of kinetic energies — some move slowly, others fast. If a high-energy molecule reaches the surface, it may overcome the intermolecular forces holding it and escape into the gas phase. This spontaneous conversion of liquid to vapour is called evaporation, and it occurs at all temperatures.
Evaporation is a surface phenomenon: only molecules at the surface can escape because molecules in the bulk are surrounded on all sides.
Cooling effect: when high-energy molecules leave, the average kinetic energy of the remaining liquid decreases — the liquid cools. Heat flows from the surroundings to the liquid, and the surroundings also cool.
Occurs at all temperatures: unlike boiling, evaporation does not require the vapour pressure to equal atmospheric pressure.
The rate of evaporation depends on several factors. Since evaporation occurs from the surface, increasing the surface area allows more molecules to escape per unit time. For liquids with the same surface area, the rate is controlled by temperature and the strength of intermolecular forces.
Surface area: larger surface → more molecules exposed → faster evaporation.
Temperature: higher temperature → more high-energy molecules → faster evaporation.
Intermolecular forces: weaker forces → molecules escape more easily → faster evaporation. Gasoline evaporates much faster than water because its London forces are weaker than water's hydrogen bonds.
Wind/airflow: removes vapour molecules from above the surface, shifting the equilibrium and accelerating evaporation.

Vapour Pressure

When a liquid evaporates in a closed container, vapour molecules accumulate above the surface. These molecules collide with the container walls (exerting pressure) and also with the liquid surface, where some are recaptured — a process called condensation. Eventually, the rate of evaporation equals the rate of condensation, and the system reaches dynamic equilibrium. At this point, the pressure exerted by the vapour is the equilibrium vapour pressure of the liquid.
At dynamic equilibrium, the number of molecules leaving the liquid surface equals the number returning — the vapour pressure remains constant
Independent of amount: vapour pressure depends only on temperature and the nature of the liquid — not on the amount of liquid, the volume of the container, or the surface area.
Surface area cancels out: a larger surface area increases both the rate of evaporation AND the rate of condensation equally, so equilibrium vapour pressure is unchanged.
Nature of the liquid matters: stronger intermolecular forces mean fewer molecules escape, so lower vapour pressure. Glycerol (0.00016 torr at 20 °C) has a far lower vapour pressure than isopentane (580 torr at 20 °C).

Vapour Pressure of Water at Various Temperatures

1
0 °C → 4.579 torr
2
10 °C → 9.209 torr
3
20 °C → 17.54 torr
4
37 °C → 47.07 torr
5
50 °C → 92.51 torr
6
70 °C → 233.7 torr
7
90 °C → 527.8 torr
8
100 °C → 760.0 torr
The most important factor controlling vapour pressure is temperature. As temperature rises, the kinetic energy of molecules increases, so more molecules have sufficient energy to escape the liquid surface. The vapour pressure increases nonlinearly — the increment grows larger as the boiling point is approached.
Nonlinear increase: from 0 °C to 10 °C, water's vapour pressure rises by 4.63 torr (4.579 → 9.209), but from 90 °C to 100 °C it rises by 232.2 torr (527.8 → 760.0).
Intermolecular forces set the starting point: at 0 °C, water starts at 4.8 torr while diethyl ether starts near 200 torr because ether's weaker forces allow far more molecules to escape.
Rapid rise near boiling point: vapour pressure curves are steep near the boiling point — a small temperature increase causes a large pressure increase.
Vapour pressure can be measured experimentally using the manometric method. The liquid is frozen, the air above it is evacuated with a vacuum pump, and the process is repeated until all air is removed. The liquid is then warmed to the desired temperature in a thermostat, and the difference in mercury levels in the two limbs of the manometer gives the vapour pressure.
The vapour pressure of the liquid equals atmospheric pressure plus the mercury height difference in the manometer
=Vapour pressure of the liquid(torr)
=Atmospheric pressure(torr)
=Difference in mercury column heights between the two manometer limbs(torr (mm Hg))
Freeze-evacuate-thaw cycle: repeated freezing and evacuation removes all dissolved air so that only the liquid's vapour remains above the surface.
Mercury column facing the liquid is depressed more because the total pressure on that side equals atmospheric pressure plus vapour pressure.
Thermostat maintains constant temperature during measurement — vapour pressure is highly temperature-dependent.

Boiling Point

When a liquid is heated, its vapour pressure increases. The boiling point is the temperature at which the vapour pressure of the liquid equals the external atmospheric pressure. At this temperature, bubbles of vapour form in the interior of the liquid — their internal pressure exceeds atmospheric pressure, so they rise to the surface and burst. A constant stream of bubbles characterises boiling.
Boiling occurs when the vapour pressure inside bubbles within the liquid matches the external pressure pushing down on the liquid surface
=Vapour pressure of the liquid(torr (or atm))
=External atmospheric pressure(torr (or atm))
→
Normal boiling point at sea level (1 atm)
Temperature stays constant during boiling: at the boiling point, added heat does not raise the temperature — it only breaks intermolecular forces and converts liquid to vapour.
Molar heat of vapourisation: the amount of heat required to vapourise one mole of a liquid at its boiling point. For water, this is 40.6 kJ mol⁻¹.
Stronger intermolecular forces → higher boiling point: water (100 °C, H-bonding) boils much higher than diethyl ether (34.6 °C, weaker forces).

Boiling Points of Common Liquids at 760 torr

•
Carbon disulphide: 46.30 °C
•
Acetone: 56.00 °C
•
Carbon tetrachloride: 76.50 °C
•
Ethanol: 78.26 °C
•
Benzene: 80.15 °C
•
Acetic acid: 118.50 °C
•
Phenol: 181.80 °C
•
Water: 100.00 °C
•
Naphthalene: 218.00 °C
Since boiling occurs when vapour pressure equals external pressure, changing the external pressure directly changes the boiling point. At higher external pressures, the liquid needs more heat to reach that pressure, so the boiling point rises. At lower pressures, the liquid boils at a lower temperature. This principle has practical applications in cooking and industrial distillation.
Higher pressure → higher boiling point: a pressure cooker traps steam, raising internal pressure and the boiling temperature of water, so food cooks faster.
Lower pressure → lower boiling point: water boils at 69 °C on Mount Everest (323 torr) and at 98 °C in Murree Hills (700 torr), compared to 100 °C at sea level (760 torr).
Vacuum distillation: by lowering the pressure, liquids boil at temperatures below their normal boiling points. Glycerin decomposes at 290 °C (its normal B.P.), but under vacuum at 50 torr it boils at just 210 °C and can be distilled without decomposition.

Energetics of Phase Changes

When matter changes state, energy is absorbed or released. At constant pressure, this heat change is called the enthalpy change (). Three key enthalpy changes accompany physical transitions: molar heat of fusion, molar heat of vapourisation, and molar heat of sublimation. All three are positive (endothermic) because energy must be supplied to overcome intermolecular forces.
The enthalpy change is always largest for sublimation, intermediate for vapourisation, and smallest for fusion — reflecting the increasing molecular separation in each process
=Molar heat of sublimation (solid → vapour)(kJ/mol)
=Molar heat of vapourisation (liquid → vapour)(kJ/mol)
=Molar heat of fusion (solid → liquid)(kJ/mol)
Molar heat of fusion ($\Delta H_f$): heat absorbed when one mole of solid melts to liquid at its melting point (1 atm).
Molar heat of vapourisation ($\Delta H_v$): heat absorbed when one mole of liquid vaporises at its boiling point (1 atm). Water: 40.6 kJ mol⁻¹.
Molar heat of sublimation ($\Delta H_s$): heat absorbed when one mole of solid sublimes directly to vapour.
because molecules separate much farther apart during vapourisation than during melting — the change in potential energy is greater.
$\Delta H_v$ values reflect intermolecular force strength: polar molecules like H₂O (40.6 kJ/mol) and NH₃ (21.7 kJ/mol) have higher than non-polar ones like CH₄ (8.6 kJ/mol).

Heats of Vapourisation of Selected Substances

•
H₂O: +40.6 kJ/mol
•
NH₃: +21.7 kJ/mol
•
SO₂: +24.3 kJ/mol
•
HCl: +15.6 kJ/mol
•
I₂: +22.0 kJ/mol
•
C₆H₁₄: +30.1 kJ/mol
•
C₃H₈: +16.9 kJ/mol
•
F₂: +5.9 kJ/mol
•
CH₄: +8.6 kJ/mol
Every change of state eventually reaches dynamic equilibrium — a condition where two opposing processes occur at equal rates. At 0 °C, ice and liquid water coexist in dynamic equilibrium. The concept of dynamic equilibrium extends to all reversible physical changes: solid-liquid, liquid-vapour, and solid-vapour transitions all reach this balanced state under the right conditions of temperature and pressure.
Phase changes are reversible — at equilibrium, forward and reverse processes occur at equal rates, so the amounts of each phase remain constant
Liquid-vapour equilibrium: at any temperature, a liquid in a closed container establishes dynamic equilibrium with its vapour.
Solid-liquid equilibrium: at the melting point, solid and liquid coexist in dynamic equilibrium (e.g., ice and water at 0 °C).
Dynamic equilibrium is the fate of all reversible changes — physical and chemical alike.

Hydrogen Bonding

Hydrogen bonding is the strongest type of intermolecular force. It occurs when a hydrogen atom covalently bonded to a highly electronegative atom (F, O, or N) is attracted to a lone pair on another electronegative atom in a neighbouring molecule. The small size of hydrogen allows it to approach closely, creating a strong electrostatic attraction that acts as a bridge between two electronegative atoms. Hydrogen bonds are about twenty times weaker than covalent bonds but significantly stronger than ordinary dipole-dipole forces.
Requires two conditions: a hydrogen atom bonded to F, O, or N (the hydrogen-bond donor), and a lone pair on F, O, or N on another molecule (the hydrogen-bond acceptor).
Water as the classic example: each H₂O molecule has two hydrogen atoms (donors) and two lone pairs on oxygen (acceptors), allowing it to form up to four hydrogen bonds.
Chlorine rarely forms hydrogen bonds, but in chloroform (CHCl₃), the three Cl atoms pull electron density away from C, giving the H atom sufficient partial positive charge to hydrogen-bond with the oxygen of acetone.
Strength order: H-F > H-O > H-N, reflecting electronegativity differences.

Compounds That Exhibit Hydrogen Bonding

•
H₂O — two H donors, two lone-pair acceptors on O
•
NH₃ — three H donors, one lone-pair acceptor on N
•
HF — one H donor, three lone-pair acceptors on F
•
C₂H₅OH (ethanol) — O-H group enables H-bonding
•
CH₃COOH (acetic acid) — O-H and C=O groups
•
CHCl₃ + CH₃COCH₃ (acetone) — intermolecular H-bond
Hydrogen bonding dramatically affects the boiling points of covalent hydrides. In groups VA, VIA, and VIIA, the first-row hydrides (NH₃, H₂O, HF) have anomalously high boiling points compared to the rest of their group. This anomaly is absent in Group IVA because those elements are not electronegative enough to form hydrogen bonds.
Group VIA trend: H₂O (100 °C) is liquid at room temperature, while H₂S (−61 °C), H₂Se (−42 °C), and H₂Te (−2 °C) are all gases — the gap is due entirely to hydrogen bonding in water.
Group VA trend: NH₃ (−33.4 °C) has a much higher boiling point than PH₃ (−88 °C), AsH₃ (−55 °C), and SbH₃ (−17 °C).
Group VIIA trend: HF (19.9 °C) boils far above HCl (−85 °C), HBr (−67 °C), and HI (−38 °C).
Why HF does not beat H₂O: although F is more electronegative than O, each HF molecule can form only one hydrogen bond (one H, multiple acceptors but limited by geometry). Each H₂O molecule can form two hydrogen bonds (two H atoms + two lone pairs = four attachment points), giving water a more extensive H-bond network.
NH₃ forms fewer H-bonds per molecule: it has three H atoms but only one lone pair, limiting it to roughly one effective H-bond per molecule on average.
HCl is borderline: its slightly higher boiling point than expected for pure London/dipole-dipole forces suggests very weak H-bonding, but it is not a strong hydrogen-bonded system.

Boiling Points of Hydrides

•
CH₄: −164 °C | NH₃: −33.4 °C | H₂O: 100 °C | HF: 19.9 °C
•
SiH₄: −112 °C | PH₃: −88 °C | H₂S: −61 °C | HCl: −85 °C
•
GeH₄: −88 °C | AsH₃: −55 °C | H₂Se: −42 °C | HBr: −67 °C
•
SnH₄: −52 °C | SbH₃: −17 °C | H₂Te: −2 °C | HI: −38 °C
Hydrogen bonding is the key to understanding why some substances dissolve in water and others do not. Water itself is an extensive H-bonded network. Substances that can also form hydrogen bonds with water — like ethyl alcohol and small carboxylic acids — dissolve readily. Non-polar substances like hydrocarbons cannot form H-bonds with water and are insoluble.
Like dissolves like (via H-bonding): ethanol (C₂H₅OH) is miscible with water because the −OH group forms hydrogen bonds with water molecules.
Small carboxylic acids dissolve: acetic acid (CH₃COOH) is soluble because its −OH and C=O groups can H-bond with water.
Hydrocarbons are insoluble: they are non-polar and lack −OH, −NH, or −FH groups, so no hydrogen bonding is possible with water.
Hydrogen bonding has wide-ranging applications across chemistry and biology. It explains the cleansing action of soaps, the structure and stability of proteins and DNA, the stickiness of glue and honey, the adhesive properties of paints and dyes, and the tensile strength of fibres like cotton and silk.
Soaps and detergents: the polar (hydrophilic) part of the molecule is water-soluble due to hydrogen bonding, while the non-polar (hydrophobic) alkyl/benzyl portion dissolves grease — together they emulsify dirt.
Proteins (alpha-helix): long chains of amino acids coil into helices. In the alpha-helix, >NH and >C=O groups on adjacent turns are linked by hydrogen bonds, with about 27 amino acid residues per turn of the helix.
DNA (double helix): two spiral strands coil around a common axis (18–20 Å diameter). H-bonds between complementary base pairs (A–T, G–C) link the two strands together.
Carbohydrates: glucose, fructose, and sucrose all contain −OH groups responsible for intermolecular hydrogen bonding.
Textiles: hydrogen bonding gives cotton, silk, and synthetic fibres their rigidity and tensile strength.
Paints, dyes, glue, and honey: their adhesive and sticky properties arise from hydrogen bonding between molecules.

Anomalous Behaviour of Water

Water exhibits a unique property among liquids: its density is maximum at 4 °C, not at its freezing point. When water freezes to ice, the molecules arrange into a regular tetrahedral lattice held together by an extensive hydrogen bond network. This ordered structure contains large empty spaces, causing ice to occupy about 9% more volume than liquid water. As a result, ice is less dense than liquid water and floats.
Tetrahedral structure of ice: each water molecule sits at the centre of a tetrahedron, connected to four neighbours via hydrogen bonds (analogous to the carbon lattice in diamond).
Liquid water is denser: in the liquid state, H-bonds constantly break and reform, allowing molecules to pack more closely — hence liquid water is denser than ice.
Density maximum at 4 °C: as liquid water cools from room temperature to 4 °C, it contracts and becomes denser. Below 4 °C, the expanding effect of forming more ordered H-bond clusters outweighs thermal contraction, and density decreases.
Ice floats on water: the open hexagonal structure of ice has about 9% more volume than the same mass of liquid water.
The anomalous density behaviour of water has profound ecological consequences. In winter, as atmospheric temperature drops, the surface layer of a lake or pond cools to 4 °C and sinks because it is the densest water. Below 4 °C, the cooled water becomes less dense and stays on top, eventually freezing. The layer of ice acts as an insulating blanket, preventing the water below from freezing and allowing aquatic life to survive throughout winter.
Freezing from the top down: the surface freezes first because cold water (below 4 °C) is less dense and stays on top of warmer water beneath.
Ice insulates: the layer of ice has low thermal conductivity and prevents further heat loss from the water underneath.
Survival of aquatic life: fish and plants survive under the ice blanket because liquid water persists below the frozen surface, maintaining temperatures around 4 °C.
Without H-bonding in water, ponds would freeze from the bottom up, killing all aquatic life — the pattern of life on Earth would be fundamentally different.