Types of Solids

Types of Solids

SCH 4U1

Types of Solids

We will classify solids into four types:1.Ionic Solids2.Metallic Solids3.Molecular Solids (Non Polar and Polar)4.Covalent network solids (3D and 2D)

Ionic Solids

An ionic solid is a solid that consists of cations and anions held together by the electrical attraction of opposite charges (ionic bonds)

This is technically considered an intra-ion force and results in ionic compounds having very high melting points.

Ionic Bonding

The forces involved in ionic compounds are derived from the creation of a crystal lattice composed of alternating negative and positive ions.

Because of the full charge on ions, ionic bonds are stronger than dipole-dipole interactions

Ionic Bonding

Ionic compounds have the following physical properties: Very high melting points Brittle Non-conductive in a solid state Conductive in a liquid state and when

dissolved in water (electrolyte) May dissolve in water, but not all

ionic compounds do!

Arrangement of Ions in a Sodium chloride crystal

Metallic solids

Metals usually have very high melting points, although this can be highly variable.

For example, mercury melts at -39oC while tungsten melts at 3410oC

Therefore, the forces in metals must be quite strong.

Metallic solids

In metals, the valence electrons of neighbouring atoms form a sort of ‘electron sea’

These electrons are called “delocalized electrons”

This ‘delocalized electron soup’ can be thought of as a ‘glue’ that holds the positive nuclei of the metal atoms together.

Metallic solids

The positive nuclei of metals form lattice structures within the electron soup.

This means that metals are composed of crystals, but they are usually too small to see.

However, formation of metals through electroplating can make the crystalline nature more obvious.

Different Types of Crystal Lattices

Gallium

Molecular solids

A molecular solid is a solid that consists of atoms or molecules held together by intermolecular forces

Examples include: Solid neon (melting point of -248oC) Solid CO2 (dry ice)

Solid H2O (ice) Iodine Sugar

Molecular solids

Consider the element Iodine, I2

Iodine is a solid at room temperature and pressure

What intermolecular forces maintain this solid state?

Only London forces! However, it has 106 electrons, giving it

sufficient intermolecular force to be a solid at room temperature.

Molecular solids

Molecular solids are held together more strongly if made up of polar molecules (dipole-dipole forces) rather than non-polar molecules (London dispersion forces).

Molecular solids tend to be soft, waxy, flexible, and have low melting points and water solubility.

Covalent Network Solid

Consider carbon dioxide (CO2) and silicon dioxide (SiO2).

What would you expect the physical properties of SiO2 to be?

CO2 b.p. = -78.5oC If London forces are the only

intermolecular force, then you might predict the b.p. of SiO2 to be slightly more than CO2.

Covalent Network Solid

However SiO2 has a m.p. of 1650oC, and b.p. of 2230oC!

SiO2 is also known as quartz, or sand, is used in the production of glass.

Clearly, SiO2 is not a molecular solid like CO2 is.

SiO2 is a covalent network solid.

Raw silica (SiO2)

Silica glass (SiO2)

Amethyst quartz (SiO2)

Covalent Network Solid A covalent network solid is a solid that

consists of atoms held together in large networks or chains by covalent bonds. Very common with Group 14 elements - can make 4 bonds.

Every atom is covalently bonded forming a 3-dimensional network

Examples include: diamond, graphite, silicon, quartz, asbestos

Diamond - 3D network solid

3D network solids are hard and are generally transparent or translucent

Diamond is different allotrope of carbon than graphite

Like graphite, it is a covalent network solid, but it is 3-dimensional lattice of carbon atoms.

This is what gives diamond its characteristic hardness.

Graphite - 2D network solid

It is a 2D covalent network solid, which means it is soft and shiny

It consists of ‘sheets’ of covalently bonded carbons

These sheets can slide over one another easily, making graphite useful for writing (i.e. it is the ‘lead’ in pencils!)

It is also an electrical conductor in one direction only because of delocalized electrons between sheets of carbon.

Graphite, Diamond and Silicon dioxide (Silica)