Liquids, Solids, Solutions, and Colloids

This is Unit Eight of Chemistry, Part Two: solids and crystal structure. Part One ended at the phase diagram, having established that intermolecular forces decide which phase a substance occupies. This part goes inside the solid phase and asks what holds it together and how its particles are arranged. The through-line is that the four kinds of solid are distinguished by what has to be overcome to melt them — and that turns out to predict hardness, conductivity and melting point together. By the end you will be able to say what distinguishes a crystalline solid from an amorphous one; say why an ionic solid shatters rather than bends; say what separates a covalent network solid from a molecular one; say why a corner particle counts as one-eighth of a unit cell; and say what iodine shows about density against force strength.

Six entries. First a distinction about arrangement that cuts across everything else, then the four bonding types in order of what melting them costs, and finally the procedure for identifying an unknown from properties you can measure.

Ordered, or frozen random.

Chem 1451 Crystalline vs. Amorphous Solids

A crystalline solid has its particles arranged in a definite, repeating pattern; an amorphous solid's particles are frozen in place randomly instead. The two forms can share an identical chemical formula — the distinction is arrangement rather than composition, which is why the same substance can occur as either.

Hard, and brittle.

Chem 1454 Ionic Solids

An ionic solid is held together by electrostatic attraction between ions, giving high melting points and hardness combined with brittleness. It shatters rather than bends because shifting one layer of ions brings like charges into alignment, and the resulting repulsion drives the layers apart.

A sea of electrons.

Chem 1457 Metallic Solids

A metallic solid is held by metallic bonding — a shared sea of delocalised electrons surrounding fixed metal atom positions, which is what makes metals conduct. Calling a solid metallic does not imply a uniformly high melting point: liquid mercury and transition metals melting above a thousand degrees share the same delocalised-electron bonding.

One giant molecule.

Chem 1460 Covalent Network Solids

A covalent network solid is held by a continuous three-dimensional network of covalent bonds spanning the entire crystal. Melting one means breaking actual covalent bonds throughout the structure rather than overcoming intermolecular attraction — which is why these are the hardest solids and demand the highest temperatures.

Molecules, held loosely.

Chem 1463 Molecular Solids

A molecular solid is held by intermolecular forces between intact neutral molecules, giving lower melting points than the other three types. Its melting point tracks intermolecular force strength rather than any internal bond — the molecules themselves stay intact regardless of temperature, which is the distinction from a covalent network solid.

Test, in order.

Chem 1466 Classifying an Unknown Solid from Its Properties

Classifying an unknown crystalline solid means checking in sequence: does it conduct as a solid, which means metallic; does it conduct only when melted or dissolved, which means ionic; does it resist melting below a very high temperature, meaning covalent network; or does it melt comparatively low without conducting in any state, meaning molecular. Hardness alone cannot decide — ionic and covalent network solids are both hard, so conductivity and melting point must be checked before an identification is final.

Three entries. What happens where the ideal pattern breaks, the smallest block whose repetition builds the whole crystal, and the count that describes how many neighbours a particle actually touches.

Where the pattern breaks.

Chem 1471 Crystal Defects

A crystal defect breaks a solid's ideal repeating pattern — a vacancy where a particle is missing, or an interstitial where one fills a gap it should not occupy. A defect is a departure from ideal structure rather than automatically a flaw: doping introduces defects deliberately, to change a material's properties.

The smallest block.

Chem 1476 The Unit Cell

A unit cell is the smallest repeating structural unit whose repetition in three dimensions generates an entire crystal lattice. Its corner and edge particles are shared with neighbouring cells, so the number belonging to any one cell is a fraction of what a drawing appears to show — which is the point the next topic's counts all rest on.

How many touch.

Chem 1479 Coordination Number

A particle's coordination number in a crystal is the number of other particles it directly contacts. It describes direct contact only — particles sitting nearby but not touching are not counted, however close they may appear in a diagram.

Four entries. Three cubic arrangements in increasing density of packing, each with its own particle count and coordination number, and then a density calculated from one of them.

Corners only.

Chem 1482 The Simple Cubic Unit Cell

A simple cubic unit cell has one particle at each of its eight corners, and because each corner particle is shared among eight neighbouring cells, only one-eighth of each belongs to this one — contributing one full particle in total. Counting a whole particle at every corner overstates the cell's contents eightfold.

One in the middle.

Chem 1485 The Body-Centered Cubic Unit Cell

A body-centred cubic cell adds one particle at its centre to the corner arrangement, giving two particles in total and a coordination number of eight. The corner atoms do not touch each other in this arrangement — each contacts the central atom, which is where the coordination number of eight comes from.

One on each face.

Chem 1488 The Face-Centered Cubic Unit Cell

A face-centred cubic cell adds one particle centred on each of its six faces, giving four particles in total and a coordination number of twelve. A face-centred particle is shared between two cells and so contributes one-half each, unlike a corner particle's one-eighth share.

Mass over volume.

Chem 1493 Computing a Metal's Density from Its Unit Cell

Computing a metal's density from its unit cell means finding the mass of the particles belonging to that cell, from their count and molar mass, finding the cell's volume from its edge length cubed, and dividing the first by the second. Leaving the edge length in picometres instead of converting to centimetres before cubing produces a density wrong by many orders of magnitude, not a small error.

Two entries to close the part. How a crystal built from two different ions arranges them, and then the capstone, which takes a single substance and shows that its vapour pressure, phase behaviour, solid type and density are all consequences of one thing.

One lattice, two ions.

Chem 1498 Ionic Crystal Unit Cells

An ionic crystal's unit cell treats one ion's positions as the lattice, with the other ion filling the gaps that lattice leaves. Sodium chloride's face-centred cubic cell contains four sodium ions and four chloride ions, keeping the required one-to-one ratio. The two ion counts must reduce to the compound's actual formula ratio, or the calculation has gone wrong somewhere.

Iodine, start to finish.

Chem 1503 Capstone: From Dominant IMF to Crystal Structure

Solid iodine is nonpolar with no hydrogen available for hydrogen bonding, so dispersion forces are its only intermolecular force — the weakest category. Consistent with that, it has a high vapour pressure for a solid and visibly sublimes at room temperature. It crystallises as a molecular solid, discrete molecules held at lattice positions by attraction alone, melting at a low 114 degrees rather than the thousands a covalent network would demand. Yet its calculated density comes out near 4.9 grams per cubic centimetre — dense, because density reflects how tightly molecules pack rather than how strongly they are held.

Five things to carry out of this part. A crystalline solid has its particles in a definite repeating pattern and an amorphous solid has them frozen randomly — a difference of arrangement rather than composition, so the same formula can give either. An ionic solid shatters rather than bends because its structure depends on opposite charges sitting adjacent: shift one layer and like charges come into alignment, and the repulsion drives the layers apart. What separates a covalent network solid from a molecular one is what melting has to overcome — a covalent network is one continuous lattice of covalent bonds, so melting it breaks those bonds and requires enormous temperatures, while a molecular solid is held only by intermolecular forces between molecules that stay intact throughout. A corner particle counts as one-eighth of a unit cell because it is shared among the eight cells meeting at that corner, which is why a simple cubic cell with eight corner particles contains only one; a face-centred particle, shared between two cells, counts as one-half. And iodine shows that density and intermolecular force strength are independent: its dispersion-only bonding is the weakest kind, yet its density is high, because density measures how tightly molecules pack rather than how strongly they are held in place. Part Three takes up solutions and the colligative properties.