Solids notes

MDCAT Chemistry

Solids have a definite volume and usually a definite shape because their particles are closely packed and held by strong attractive forces. This chapter explains crystalline solids, crystal lattices, ionic and molecular crystals, covalent networks, crystal properties, and lattice energy.

General Characteristics of Solids

A solid has a fixed shape and a fixed volume. Its particles are very close to one another and can only vibrate about their mean positions. The attractive forces between particles are strong, while the spaces between them are very small.

Solids are almost incompressible because their particles have little empty space. They usually have high density and diffuse very slowly. The arrangement and type of particles determine whether a solid is ionic, molecular, covalent or metallic.

Solids have definite shape and definite volume.
Particles in a solid are closely packed and vibrate around fixed positions.
Strong attractive forces exist between particles in solids.
Solids are almost incompressible because of very small intermolecular spaces.
Solids have very slow diffusion compared with liquids and gases.
A crystalline solid has a regular arrangement of particles, while an amorphous solid has an irregular arrangement.

Crystalline and Amorphous Solids

A crystalline solid contains particles arranged in a definite, repeating three-dimensional pattern. It has a definite geometrical shape, a sharp melting point and definite cleavage planes. Sodium chloride, diamond, quartz and ice are examples of crystalline solids.

An amorphous solid does not have a regular long-range arrangement of particles. It softens over a range of temperature instead of melting sharply. Glass, rubber, plastics and pitch are common examples. Table salt is crystalline, so it is not an amorphous solid.

Crystalline solids have a regular and repeating arrangement of particles.
Crystalline solids have a sharp melting point.
Crystalline solids are anisotropic, so some physical properties may differ in different directions.
Amorphous solids soften over a range of temperature.
Amorphous solids are generally isotropic, so their properties are similar in all directions.
Table salt is crystalline, whereas glass is amorphous.
The property of a crystal to grow in its usual characteristic shape is called the habit of the crystal.
X-ray diffraction is used to study crystal structures and to determine atomic or ionic radii.

Crystal Lattice and Unit Cell

A crystal lattice is a regular three-dimensional arrangement of points representing the positions of particles in a crystal. The particles may be atoms, ions or molecules. The smallest repeating portion of a crystal lattice is called a unit cell.

Repeating a unit cell in all three directions produces the complete crystal. The shape and dimensions of a unit cell depend on the arrangement of particles. In an ionic crystal, the lattice shows the regular arrangement of oppositely charged ions.

A crystal lattice is a regular, three-dimensional pattern of particles or lattice points.
A unit cell is the smallest repeating part of a crystal lattice.
Repetition of unit cells forms the complete crystal.
Particles at a corner of a unit cell are shared by eight neighbouring unit cells.
A particle at a face centre is shared by two unit cells.
A particle at an edge is shared by four unit cells.
A body-centred particle belongs completely to one unit cell.
The properties of a crystal depend on the arrangement of particles in its lattice.

Shape, Radius and Related Crystal Terms

The shape of an ionic crystal depends on the arrangement of positive and negative ions in its lattice. Ions are considered to be spherical when ionic radii are discussed. The ionic radius is the effective radius of an ion in a crystal structure.

Some substances can exist in different crystalline forms. If an element exists in more than one crystalline form, the phenomenon is called allotropy. Diamond and graphite are allotropes of carbon. If two different compounds have similar chemical formulae and the same crystalline form, they are called isomorphic substances.

Ionic radius is the radius of an ion when the ion is considered spherical.
Allotropy is the existence of an element in more than one crystalline form.
Diamond and graphite are allotropes of carbon.
Isomorphism is the existence of different compounds with analogous formulae and the same crystalline form.
The habit of a crystal is its usual external geometrical shape.
Crystal shape is related to the internal arrangement of particles in the lattice.
X-ray diffraction helps determine interparticle distances and atomic or ionic radii.

Ionic Crystals

Ionic crystals consist of positive and negative ions arranged in a regular lattice. Each ion is surrounded by ions of opposite charge. Strong electrostatic forces act in all directions, so ionic solids are hard and have high melting points.

Ionic crystals do not contain separate molecules. Their formula represents the simplest whole-number ratio of ions in the crystal lattice. For example, NaCl represents a one to one ratio of sodium ions to chloride ions, not separate NaCl molecules.

Ionic solids contain cations and anions arranged in a crystal lattice.
Ionic solids have strong electrostatic forces between oppositely charged ions.
Ionic solids are generally hard and brittle.
Ionic solids have high melting and boiling points.
Ionic solids do not conduct electricity in the solid state because their ions cannot move freely.
Molten ionic compounds and aqueous solutions conduct electricity because ions become mobile.
Ionic solids are generally soluble in polar solvents such as water.
When an ionic crystal is struck, layers may shift and like charges come close, causing repulsion and fracture.

Molecular and Covalent Crystals

Molecular crystals consist of discrete molecules held together by intermolecular forces. These forces may be London dispersion forces, dipole-dipole forces or hydrogen bonding. Molecular crystals usually have lower melting points than ionic and covalent network solids.

Dry ice, solid CO2, consists of discrete CO2 molecules. Ice and sugar are examples of polar molecular solids. The strength of London forces increases with increasing size and number of electrons, which explains the physical states of the halogens at room temperature.

Molecular crystals contain discrete molecules at lattice points.
Dry ice is a molecular crystal made of discrete CO2 molecules.
Ice is a polar molecular solid because H2O molecules are polar and hydrogen bonded.
Sugar is an example of a polar molecular solid.
London forces become stronger as molecular size and number of electrons increase.
Among halogens, iodine is solid, bromine is liquid, and fluorine and chlorine are gases at room temperature because London forces increase down the group.
Molecular solids are generally soft and have low melting points compared with ionic solids.
Covalent network crystals contain atoms joined by strong covalent bonds throughout the structure.

Covalent Network Solids and Silica

In a covalent network solid, atoms are linked by covalent bonds in a continuous three-dimensional or layered network. Quartz is a covalent crystal. Silicon dioxide has a giant covalent structure rather than separate SiO2 molecules.

In the network of silica, each silicon atom is tetrahedrally surrounded by four oxygen atoms. Each oxygen atom is bonded to two silicon atoms. Therefore, one silicon atom is associated with four silicon-oxygen connections, while each oxygen is shared between two silicon atoms. The resulting ratio is Si:O = 1:2.

Quartz is a covalent crystal.
In SiO2, each silicon atom is surrounded by four oxygen atoms.
In SiO2, each oxygen atom is bonded to two silicon atoms.
The silicon to oxygen ratio in silica is 1:2.
Silicon dioxide has a giant network structure, not discrete SiO2 molecules.
Diamond is a three-dimensional covalent network solid.
Diamond is the hardest naturally occurring substance in the usual FSc comparison.
Diamond is a bad conductor of electricity because it has no free electrons.
Graphite has layers of carbon atoms and conducts electricity because it has delocalised electrons.

Comparison of Diamond, Graphite and Molecular Solids

Diamond and graphite both contain carbon atoms, but their structures are different. In diamond, every carbon atom forms four covalent bonds in a three-dimensional network. In graphite, each carbon atom forms three covalent bonds in a planar layer, leaving one electron delocalised.

The layers in graphite are held together by weak forces and can slide over one another. This makes graphite soft. Diamond has strong covalent bonds throughout its structure, making it extremely hard. The carbon-carbon bond distance in graphite is approximately 1.426 Å.

Diamond is very hard, while graphite is soft.
Diamond is a bad conductor because it has no free electron available for conduction.
Graphite conducts electricity because it contains delocalised electrons.
Diamond has a three-dimensional covalent network.
Graphite has planar layers of carbon atoms.
Graphite is used as a lubricant because its layers slide over one another.
The approximate carbon-carbon bond distance in graphite is 1.426 Å.
Both diamond and graphite are allotropes of carbon.

Hydrates, Sublimation and Other Crystal Examples

Some crystalline compounds contain a definite number of water molecules as part of their crystal structure. Such compounds are called hydrates. The water is called water of crystallisation. It is chemically associated with the crystal and is not merely moisture on the surface.

Sublimation is the direct conversion of a solid into vapour without passing through the liquid state. It can be used to separate a sublimable substance from a non-sublimable substance. A mixture of sand and naphthalene can be separated by sublimation because naphthalene sublimes while sand does not.

Compounds containing a definite number of water molecules in their crystals are called hydrates.
Water present in a hydrate is called water of crystallisation.
Sublimation is the direct change of a solid into vapour.
Naphthalene sublimes on heating, while sand does not.
Sand and naphthalene can be separated by sublimation.
Dry ice is solid CO2 and changes directly into gaseous CO2 on warming.
A crystal may lose water of crystallisation on heating.

Lattice Energy and Born-Haber Cycle

Lattice energy is the energy released when one mole of an ionic solid is formed from its gaseous ions. It may also be described as the energy required to separate one mole of an ionic solid into its gaseous ions. The sign depends on the convention being used, but the magnitude represents the strength of attraction in the lattice.

Lattice energy depends mainly on ionic charge and the distance between ion centres. Higher ionic charges produce stronger electrostatic attraction. Increasing ionic size increases the distance between oppositely charged ions, so they become less tightly held and lattice energy decreases.

Lattice energy measures the strength of attraction between oppositely charged ions in an ionic crystal.
Lattice energy increases when the charges on the ions increase.
Lattice energy decreases when ionic size increases.
As ion size increases, oppositely charged ions become less tightly held.
Lattice energy is associated with the formation and breakdown of an ionic lattice.
The Born-Haber cycle is used to calculate lattice energy indirectly.
The Born-Haber cycle applies Hess's law to a series of thermochemical steps.
The steps may include sublimation, bond dissociation, ionisation energy, electron affinity and formation of the ionic lattice.

Key terms

Solid: A state of matter with definite shape and volume whose particles are closely packed.
Crystalline solid: A solid whose particles have a regular, repeating long-range arrangement.
Amorphous solid: A solid with an irregular particle arrangement that softens over a range of temperature.
Crystal lattice: A regular three-dimensional arrangement of particles in a crystal.
Unit cell: The smallest repeating part of a crystal lattice.
Ionic crystal: A crystal made of regularly arranged cations and anions held by electrostatic forces.
Molecular crystal: A crystal made of discrete molecules held by intermolecular forces.
Covalent crystal: A solid in which atoms are joined throughout a giant network by covalent bonds.
Ionic radius: The effective radius of an ion considered to be spherical.
Allotropy: The existence of an element in more than one crystalline form.
Isomorphism: The occurrence of different compounds with analogous formulae and the same crystalline form.
Habit of crystal: The usual external geometrical shape in which a crystal grows.
Hydrate: A compound containing a definite number of water molecules in its crystal structure.
Water of crystallisation: Water chemically incorporated in the crystal structure of a hydrate.
Lattice energy: The energy associated with forming an ionic crystal from gaseous ions or separating it into gaseous ions.
Born-Haber cycle: A Hess's law cycle used to calculate lattice energy indirectly.
London force: A temporary attraction between particles caused by instantaneous and induced dipoles.
Sublimation: The direct conversion of a solid into vapour without formation of a liquid.

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Chemistry shortcuts

Finding the limiting reactant and percentage composition

Convert every given mass or volume into moles first. The reactant that produces the least amount of the required product is the limiting reactant.

Write the balanced equation and calculate moles using n = mass/Mr.
Use the mole ratio to calculate the product. For percentage composition, use percentage = mass of element in one mole of compound divided by molar mass, multiplied by 100.
Example: Percentage of nitrogen in KNO3 = 14/101 × 100 = 13.86%.
Answer: 13.86% nitrogen.

Use gas volume at molar volume only when the gas conditions are stated or are standard conditions.

Using gas volume, pressure and temperature relations

At the same temperature and pressure, gas volume is directly proportional to the number of molecules. For changing conditions, use P1V1/T1 = P2V2/T2.

At constant temperature and pressure, divide or multiply the volume in the same ratio as the number of molecules.
Example: 10 mL H2 contains 2 × 10^3 molecules. Oxygen in 200 mL contains 20 × 2 × 10^3 = 4 × 10^4 molecules.
Answer: 4 × 10^4 molecules.
For a rigid container, increasing temperature increases molecular speed and mean free path if the gas remains in the same phase.

The direct volume to molecule ratio does not apply when temperature or pressure changes.

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