05/14 Intermolecular Forces and Phases of Matter

A progressive guide to intermolecular attractions, particle behavior, liquid and solid properties, phase changes, heating curves, vapor pressure, and phase diagrams, including key quantitative relationships and examples.

Particle-Level View and

The physical state of a substance reflects a competition between particle kinetic energy and attractions between particles. Stronger attractions favor condensed phases, while greater kinetic energy favors separation into the gas phase.

In a solid, particles are closely packed and mainly vibrate about fixed positions. The substance has a definite shape and volume.
In a liquid, particles remain close together but can move past one another. The substance has a definite volume but takes the shape of its container.
In a gas, particles are widely separated and move independently. The substance has neither a definite shape nor a definite volume.

Increasing temperature increases the average kinetic energy of the particles. Phase changes generally alter attractions between particles without breaking the covalent bonds within the molecules.

Comparing the main attractions

The most important attractions to identify are , , , and .

result from temporary dipoles and occur in every atom and molecule. They become stronger with more electrons, greater polarizability, and greater molecular contact area.
occur between polar molecules with permanent partial charges.
occurs when hydrogen bonded to nitrogen, oxygen, or fluorine is attracted to a lone pair on nitrogen, oxygen, or fluorine in a neighboring particle.
occur between ions and polar molecules and help explain the dissolution of ionic substances in polar solvents.

When comparing substances, identify the strongest important attraction first, then consider molecular size, shape, surface area, and polarizability. Less-branched molecules often have greater contact area than compact, highly branched isomers. Larger halogen molecules generally have stronger dispersion forces and higher boiling points.

Stronger effective attractions usually correspond to higher boiling points, higher melting points, greater viscosities, and greater surface tensions, but lower vapor pressures.

Takeaway: Particle-level attractions and kinetic energy together explain why a substance occupies a particular phase and why related substances have different physical properties.

Liquid Properties: Flow, Surfaces, and Capillaries

Macroscopic liquid properties can be explained by how strongly particles attract one another and how readily they move.

measures resistance to flow. Strong attractions and complicated or elongated molecules make it more difficult for molecules to slide past one another, increasing . Heating generally decreases a liquid’s because faster-moving particles more readily overcome attractions. Honey, for example, flows more slowly than water because its molecular interactions and structures resist motion more strongly.

describes the energy required to increase a liquid’s surface area or the tendency of the surface to resist stretching. A molecule inside the liquid is attracted in all directions, whereas a molecule at the surface experiences a net inward attraction. The liquid therefore tends to minimize its surface area, which helps explain why small droplets are approximately spherical.

Cohesion is attraction between molecules of the same substance. Adhesion is attraction between a liquid and another material. results from the balance of cohesion, adhesion, , and gravity. Water rises in clean glass when adhesion to the glass draws the liquid upward and cohesion helps pull additional water molecules along. Mercury behaves differently because its cohesion is stronger relative to its adhesion to glass.

Takeaway: Stronger cohesion generally increases , while depends on both intermolecular attractions and how easily the molecular structures can move past one another.

Solid Structures and Properties

The particles and attractions in a solid determine many of its physical properties.

Molecular solids contain molecules held together by dispersion forces, , or . They often have relatively low melting points and are usually poor electrical conductors.
Ionic solids contain cations and anions held by electrostatic attractions. They typically have high melting points, are brittle, and conduct electricity when molten or dissolved.
Metallic solids contain metal atoms or cations with delocalized electrons. They are generally conductive, malleable, and ductile.
Network covalent solids contain atoms joined in an extended covalent network. They are generally very hard and have very high melting points.

A crystalline solid has a regular, repeating arrangement of particles and usually melts over a narrow temperature range. An amorphous solid lacks long-range order and may soften over a range of temperatures.

Melting requires disruption of the forces or bonds that maintain the solid structure. Consequently, stronger attractions or stronger bonding generally lead to higher melting points, although crystal structure and particle arrangement also matter.

Takeaway: Classifying a solid by its particles and the forces holding them together helps predict conductivity, hardness, flexibility, and melting behavior.

Energy and Phase Changes

A changes the physical state while preserving chemical identity. The direction of energy transfer depends on whether particles must move farther apart or come closer together.

Melting or fusion: solid to liquid; endothermic.
Vaporization: liquid to gas; endothermic.
Sublimation: solid to gas; endothermic.
Freezing or solidification: liquid to solid; exothermic.
Condensation: gas to liquid; exothermic.
Deposition: gas to solid; exothermic.

For a at equilibrium, calculate heat with:

q=n\Delta H

Here, $q$ is heat transferred, $n$ is the amount in moles, and $\Delta H$ is the molar enthalpy of the . For melting ice:

\mathrm{H_2O(s)\rightarrow H_2O(l)}\qquad \Delta H_{\mathrm{fus}}=+6.01\ \mathrm{kJ\ mol^{-1}}

Freezing has the opposite sign:

\mathrm{H_2O(l)\rightarrow H_2O(s)}\qquad \Delta H_{\mathrm{frz}}=-6.01\ \mathrm{kJ\ mol^{-1}}

The enthalpy of vaporization is normally much larger than the enthalpy of fusion because vaporization separates particles almost completely, whereas melting only partially disrupts the attractions in a solid.

Takeaway: Endothermic phase changes absorb energy to overcome attractions; exothermic phase changes release energy as attractions become more effective.

Heating Curves and Heat Calculations

A tracks temperature as heat is added. Sloped portions correspond to temperature changes within one phase. Horizontal portions correspond to phase changes, during which temperature remains constant.

For a single phase, use:

q=mc\Delta T

Here, $m$ is mass, $c$ is specific heat capacity, and $\Delta T=T_{\mathrm{final}}-T_{\mathrm{initial}}$ .

A typical sequence is:

The solid warms and its temperature increases.
Melting occurs while solid and liquid coexist; temperature remains constant.
The liquid warms and its temperature increases.
Boiling occurs while liquid and gas coexist; temperature remains constant.
The gas warms and its temperature increases.

During a horizontal segment, added energy increases the potential energy associated with separating particles rather than increasing their average kinetic energy. Because temperature measures average kinetic energy, temperature remains constant until the is complete.

For a multistep calculation, determine the heat for each segment and add the results:

q_{\mathrm{total}}=q_1+q_2+q_3+\cdots

For example, vaporizing $2.00\ \mathrm{mol}$ of water at its when $\Delta H_{\mathrm{vap}}=40.7\ \mathrm{kJ\ mol^{-1}}$ requires:

q=n\Delta H_{\mathrm{vap}}

q=(2.00\ \mathrm{mol})(40.7\ \mathrm{kJ\ mol^{-1}})=81.4\ \mathrm{kJ}

The result is positive because vaporization is endothermic.

Takeaway: Use $q=mc\Delta T$ on sloped regions and $q=n\Delta H$ on phase-change plateaus.

, Boiling, and Equilibrium

When molecules escape from a liquid surface, they enter the gas phase. In a closed container, gas molecules can return to the liquid. Once vaporization and condensation occur at equal rates, the system has , and the vapor exerts an equilibrium .

increases with temperature because more particles have enough kinetic energy to escape. Stronger intermolecular attractions make escape more difficult, producing lower . The is reached when equals external pressure.

Lower external pressure lowers the .
Higher external pressure raises the .
The normal is defined at an external pressure of $1\ \mathrm{atm}$ , approximately $101.3\ \mathrm{kPa}$ .

The can estimate how changes with temperature:

\ln\left(\frac{P_2}{P_1}\right)=\frac{\Delta H_{\mathrm{vap}}}{R}\left(\frac{1}{T_1}-\frac{1}{T_2}\right)

Use temperatures in kelvins and consistent pressure units. At the same temperature, ethanol has a higher than water. This indicates that ethanol molecules escape more readily, whereas water’s extensive hydrogen-bonding network holds its molecules together more strongly.

Takeaway: Higher temperature raises , stronger attractions lower , and boiling occurs when matches external pressure.

Phase Diagrams and Critical Behavior

A phase diagram maps the stable phase of a substance as a function of temperature and pressure. Each boundary represents conditions where two phases coexist in equilibrium.

The solid–liquid boundary separates solid and liquid regions and represents melting or freezing.
The liquid–gas boundary separates liquid and gas regions and represents vaporization or condensation.
The solid–gas boundary separates solid and gas regions and represents sublimation or deposition.

The three boundaries meet at the , where solid, liquid, and gas coexist. The liquid–gas boundary ends at the . Above the critical temperature, no amount of pressure can convert the substance into an ordinary liquid; the substance is a supercritical fluid.

Water has an unusual solid–liquid boundary that slopes slightly to the left. Ice is less dense than liquid water, so increasing pressure can favor melting under appropriate conditions.

Takeaway: Phase diagrams show how temperature and pressure determine stable phases and identify the conditions for coexistence, critical behavior, and unusual transitions.

Applying the Particle-to-Property Framework

The central strategy is to connect observations to particle-level explanations.

A higher generally suggests stronger effective attractions, although molecular size, shape, polarity, and crystal structure must also be considered.
A lower suggests that particles are held more strongly in the liquid.
A slower flow rate suggests greater , but temperature and molecular shape also affect the result.
A larger suggests stronger cohesive forces at the liquid surface.
A high melting point may reflect strong intermolecular attractions, strong ionic or metallic interactions, or an extended covalent network.

For quantitative work:

Control relevant variables such as mass, temperature, container geometry, and atmospheric pressure.
Use $q=mc\Delta T$ for temperature-change regions.
Use $q=n\Delta H$ for phase-change regions.
Convert Celsius to kelvins for gas-law and Clausius–Clapeyron calculations.
Match units, such as grams with specific heat in $\mathrm{J\ g^{-1}\ ^\circ C^{-1}}$ , or kilograms with compatible SI units.
Interpret trends rather than relying on a single isolated measurement.

The strongest explanation combines the observed macroscopic property with particle spacing, kinetic energy, molecular polarity, polarizability, shape, and the strength of relevant attractions.

Final takeaway: Intermolecular attractions, particle motion, and energy transfer form one connected framework for explaining phases of matter and their measurable properties.

Particle-Level View and Intermolecular Forces