Liquids, Solids, Solutions, and Colloids

This is Unit Eight of Chemistry, Part One: intermolecular forces, liquids and phase behaviour. Unit Seven treated gases, where molecules are far enough apart that the forces between them can mostly be ignored. This unit is what happens when they cannot. The through-line is a single balance: whether a substance is a solid, a liquid or a gas depends on how its intermolecular forces compare against its particles' kinetic energy — and almost everything in this part is a consequence of that comparison. There is also one word doing damage throughout, which is "bond", and the first topic disarms it. By the end you will be able to say what "breaking bonds" gets wrong about boiling; name the intermolecular force every substance has; say what a hydrogen bond requires and how strong it actually is; say why a boiling point needs a pressure stated alongside it; and say what water's phase diagram does that most do not.

Three entries. What decides which phase a substance is in, a prefix that separates two very different force strengths, and a common way of describing boiling that silently swaps one for the other.

Forces against motion.

Chem 1361 Condensed Phases and the IMF-KE Balance

A substance exists as a solid, a liquid or a gas depending on the balance between its intermolecular forces and its particles' kinetic energy. Phase is a balance rather than a fixed property — the same substance shifts phase as temperature or pressure moves it to the other side of that comparison.

One prefix apart.

Chem 1364 Intermolecular vs. Intramolecular Forces

Intramolecular forces hold one molecule's own atoms together; intermolecular forces are the weaker attractions between neighbouring molecules. The two words differ by a single prefix and describe forces roughly twenty-five times apart in strength — which is why confusing them is not a small slip.

What boiling breaks.

Chem 1367 The "Breaking Bonds" Misreading of Boiling

Calling boiling water "breaking its bonds" misapplies the covalent sense of the word to the far weaker hydrogen bonds actually broken between molecules. Water molecules leaving a boiling pot are still H2O — only the attractions between them gave way. If boiling broke covalent bonds, steam would not be water.

Six entries, working from the weakest force to the strongest. Each one requires something specific of a molecule, they stack rather than replace each other, and the topic closes with the procedure this whole unit rests on — identifying which force dominates in a given substance.

Everywhere, always.

Chem 1370 Dispersion Forces

A dispersion force is a weak, temporary attraction arising from momentary, fluctuating electron distributions that induce a matching dipole in a neighbour. Dispersion forces exist between every pair of atoms or molecules — the one intermolecular force no condensed phase ever lacks, and the only one nonpolar substances have.

Size, and shape.

Chem 1373 Ranking Boiling Points by Dispersion-Force Strength

Ranking a series of nonpolar molecules by boiling point means comparing their molar masses and their shapes, since larger molecules hold more polarisable electrons and more elongated ones offer more surface contact for the temporary dipoles to act across. Ranking by molar mass alone works only when shape does not also vary — same-formula isomers separate by boiling point through shape alone.

Positive to negative.

Chem 1376 Dipole-Dipole Attraction

A dipole-dipole attraction is the electrostatic pull between the positive end of one polar molecule and the negative end of a neighbour. It layers on top of dispersion forces rather than replacing them — every polar molecule keeps its dispersion forces too, and has both acting at once.

Three elements only.

Chem 1379 Hydrogen Bonding

Hydrogen bonding is a strong dipole-dipole attraction between a hydrogen bonded to fluorine, oxygen or nitrogen and a lone pair on a neighbouring fluorine, oxygen or nitrogen atom. It requires that hydrogen to be bonded directly to one of those three — a carbon-hydrogen hydrogen does not qualify.

The name misleads.

Chem 1382 Why a Hydrogen Bond Is Not a Covalent Bond

A hydrogen bond is an intermolecular attraction only five to ten percent as strong as the covalent bond its name evokes. The word "bond" is borrowed, and the two differ sharply in strength — which is the same error the boiling misdescription makes, in its specific hydrogen-bonding form.

Which one dominates.

Chem 1385 Predicting the Dominant IMF and Ranking Boiling Points

Predicting which intermolecular force dominates means checking in order: whether hydrogen bonding is possible, meaning hydrogen bonded to fluorine, oxygen or nitrogen; then whether the molecule is polar, for dipole-dipole; and treating dispersion forces as present regardless. The strongest force found sets the boiling-point ranking against substances of similar size. Stopping at polar-or-nonpolar without asking about hydrogen bonding specifically misses the largest jump in the ranking.

Four entries, all consequences of the forces covered above. How readily a liquid flows, whether it sticks to itself or to a surface, why its surface behaves like a skin, and why some liquids climb a narrow tube while others sink in it.

Resistance to flow.

Chem 1390 Viscosity

Viscosity is a liquid's resistance to flow, set by its intermolecular force strength together with its molecules' size and shape. Treating it as a fixed property ignores temperature: raising temperature lowers viscosity, since faster molecules overcome the same intermolecular forces more easily.

To itself, or to something else.

Chem 1393 Cohesive vs. Adhesive Forces

Cohesive forces are attractions between like molecules of one liquid; adhesive forces are attractions between a liquid and a different surface. Whether a liquid spreads across a surface or beads up on it depends on which of the two wins, not on the liquid's identity alone.

An unbalanced pull.

Chem 1396 Surface Tension

Surface tension is the energy needed to increase a liquid's surface area, arising from the unbalanced inward pull on molecules at the surface, which have neighbours on one side only. It tracks intermolecular force strength directly — weaker cohesive forces mean a weaker skin, not a thicker one.

Climbing a tube.

Chem 1399 Capillary Action

Capillary action is a liquid's rise up a narrow tube against gravity, driven by adhesion to the walls outcompeting the liquid's own cohesion. A liquid rises only where adhesion exceeds cohesion — mercury's much stronger cohesion makes it dip below the surrounding level instead of climbing.

Four entries. What happens in a sealed container above a liquid, the pressure that establishes, what boiling actually is in terms of that pressure, and why quoting a boiling point without a pressure leaves something unsaid.

Both directions at once.

Chem 1404 Vaporization, Condensation, and Dynamic Equilibrium

In a closed container, a liquid reaches dynamic equilibrium with its vapour when molecules leave and rejoin the liquid at equal rates. Equilibrium means the two rates match, not that the process has stopped — molecules keep crossing between phases the whole time.

The pressure it settles at.

Chem 1407 Equilibrium Vapor Pressure

A liquid's equilibrium vapour pressure is its vapour's pressure at dynamic equilibrium, and it falls as intermolecular forces strengthen — stronger attractions hold more molecules in the liquid. Vapour pressure depends on temperature and identity, never on container size or exposed surface area; those change only how fast equilibrium is reached.

When the two match.

Chem 1410 Boiling Point from Vapor Pressure

A liquid boils at the temperature where its equilibrium vapour pressure equals the pressure of its surroundings. Boiling point is therefore not a fixed property of a substance alone — it moves with the surrounding pressure, and is only a fixed number once that pressure has been specified.

At one atmosphere.

Chem 1413 The Normal Boiling Point

A liquid's normal boiling point is its boiling point specifically at one atmosphere of surrounding pressure. Citing "the boiling point" without stating a pressure silently assumes that normal, one-atmosphere value — which is why water boils below a hundred degrees at altitude without anything being wrong.

Four entries. Why a thermometer stalls while ice melts, the two energies involved and why one is always larger, the equation relating vapour pressure to temperature, and how to get an enthalpy out of two measurements.

The thermometer stalls.

Chem 1416 Heating Curves and Phase-Change Plateaus

A substance's temperature stays constant while it changes phase, because added heat goes into overcoming intermolecular forces rather than raising kinetic energy. The heat added during a plateau is not wasted — it separates particles against attraction, which is work a thermometer cannot register.

Melting, and boiling.

Chem 1419 Enthalpy of Fusion and Enthalpy of Vaporization

Enthalpy of fusion melts a substance at its melting point; enthalpy of vaporisation vaporises it at its boiling point, and is always the larger of the two. Melting only loosens a solid's structure, while vaporisation overcomes the attractions nearly in full — a gap that is easy to underestimate.

Pressure against temperature.

Chem 1424 The Clausius-Clapeyron Relation

The Clausius-Clapeyron equation relates a liquid's vapour pressure to its kelvin temperature and its enthalpy of vaporisation, plotting as a straight line in logarithmic form. Temperature must enter in kelvin — using Celsius produces a result with no physical meaning at all.

Two readings, one enthalpy.

Chem 1427 Estimating Enthalpy of Vaporization from Two Data Points

Estimating a liquid's enthalpy of vaporisation from two vapour-pressure readings means converting both temperatures to kelvin and substituting both temperature-pressure pairs into the two-point form of the equation, then solving for the enthalpy term. Mixing Celsius into one term, or swapping which pair is point one and which is point two, produces a sign error that flips the enthalpy negative.

Five entries closing the part. The map that puts every phase on one plot, two points on it where something unusual happens, how to read a state off it, and then the one substance whose diagram runs the other way.

One plot, three phases.

Chem 1432 Constructing a Phase Diagram

A phase diagram plots pressure against temperature, combining the liquid-gas, solid-liquid and solid-gas equilibrium curves into one map of which phase is stable where. The curves themselves show phase boundaries, where two phases coexist — it is the regions between them that show the phases.

All three at once.

Chem 1435 The Triple Point

A phase diagram's triple point is the single temperature and pressure at which solid, liquid and gas all coexist in equilibrium. Below the triple point's pressure no temperature at all permits a liquid state — that is a structural fact about the diagram, not a rare or delicate condition.

Where the line ends.

Chem 1438 The Critical Point and the Supercritical Fluid

A phase diagram's critical point is the temperature and pressure above which the liquid-gas boundary vanishes entirely, forming a supercritical fluid. Above the critical temperature no amount of added pressure condenses a gas back into a distinct liquid — the two phases have genuinely merged rather than becoming hard to separate.

Find the point.

Chem 1441 Reading a Phase Diagram for a Given State

Determining a substance's state from its phase diagram means locating the point at the given temperature and pressure, then reading which labelled region — solid, liquid or gas — or which boundary curve that point falls on. Misreading which axis is which, or forgetting that these axes are often not drawn to a constant scale, places the same coordinates in the wrong region entirely.

Water runs backwards.

Chem 1446 Water's Anomalous Melting-Point Slope

Water's solid-liquid curve slopes slightly negative, so its melting point falls as pressure rises — the opposite of most substances' positive slope, and traceable to ice being less dense than liquid water. This is a genuine anomaly rather than a rule being violated: most substances, carbon dioxide among them, show the ordinary positive slope.

Five things to carry out of this part. Describing boiling as "breaking bonds" gets it wrong because it borrows the covalent sense of the word: what boiling breaks are the intermolecular attractions between molecules, which are roughly twenty-five times weaker than the forces holding a molecule together. Steam is still water, and if boiling broke covalent bonds it would not be. The intermolecular force every substance has is the dispersion force — temporary, weak, arising from fluctuating electron distributions, and present between every pair of atoms or molecules, which makes it the only force a nonpolar substance has and one that polar substances keep alongside their others. A hydrogen bond requires a hydrogen bonded directly to fluorine, oxygen or nitrogen, interacting with a lone pair on another such atom — a carbon-hydrogen hydrogen does not qualify — and despite its name it is an intermolecular attraction only five to ten percent as strong as a covalent bond. A boiling point needs a pressure stated because boiling happens where a liquid's vapour pressure equals its surroundings' pressure, so the temperature moves with the surroundings; the normal boiling point is the value at one atmosphere, which is what an unqualified figure assumes. And water's phase diagram is anomalous in its solid-liquid curve sloping negative, so its melting point falls as pressure rises, which traces back to ice being less dense than the liquid. Part Two takes up solids and their crystal structures.