Ionic, Covalent and Metallic structures of solids

Introduction• Where does the concept of bonding come from and how do

we know what structures will be formed??

• The Atom

• Valence electrons• Now we know bonding refers to the forces which hold atoms

together• There are three types of bonding Ionic, Covalent and Metallic

Vocab:

• Crystalline Solid:– Lattice:– Unit Cell:

• Amorphous solid:

• Alloy:

• Conductivity:

• Boiling Point:

Ionic• Ionic compounds are Solids at room temp

– Metal and Nonmetal– Held together by electrostatic attraction

between cations and anions– Crystaline solid … crystal lattice structure

Ionic • Ionic lattices are extremely difficult structures to break apart. As a result, all ionic substances are solids with high

melting points.• Potassium iodide, magnesium chloride and calcium oxide. All of these ionic compounds have melting points over 500 °C.• There is one method of breaking up the lattice - dissolve the ionic compound in water. Water has the ability to separate

the ions from the lattice and allow them to move freely as a solution.

Ionic• Ionic Solids

– Shape of Unit Cell determines breaking angle of crystal.

• 3 unit cells: simple cubic, body centered cubic, face centered cubic

– High Lattice Energy makes them rigid, high melting and boiling points, etc.

– Rigid structure of charges means they do not conduct electricity when solid (but do when dissolved in water or liquid)

Metallic– Only Metals– Held together by electrostatic attraction

between positive nucleus and a sea of mobile valence electrons

• Also called “delocallized covalent”

– Covalent = sharing

– Delocallized = not staying in one area

– Mobile electrons = good

conductors of electricity

and heat

Metallic

• Bonding:– Close Packing– Strong and nondirectional

• Difficult to separate metal atoms– High melting and boiling points

– Solid at room temperature (except mercury)

• Easy to move their relative position– Malleable and ductile

– Can be Amorphous Solid

Metallic

• Alloys: 2 or more metallic elements– Substitutional Alloy

• All similar size

• Structure of Alloy similar to each pure metal

• Ex: brass = copper and zinc, sterling = Ag & Cu

– Interstitial Alloy• Very different sized elements

• Holes between larger atoms filled by smaller atoms

• Ex: steel = iron & carbon; harder, stronger, less ductile than pure iron

Covalent molecules• Only nonmetal elements• Covalent compounds have no ions. Their atoms join up by sharing

electrons with their neighbours in small groups or clusters of atoms called molecules.

• For example:

• You will find that covalent compounds have low melting points. They exist either as gases (methane), liquids (water) or as easily melted solids (paraffin wax).

Covalent molecules• The covalent bonds inside the molecules are very strong. The

molecules don't break apart easily. However the forces attracting neighbouring molecules to each other are very weak. It is therefore very easy to separate molecules from one another: e.g. ammonia

• When a covalent compound melts or boils, it is the forces between the molecules which are broken. Very little energy is needed to make this happen, so covalent substances have low melting and boiling points.

Covalent Network Structures• Only Nonmetal atoms

• Giant network of localized covalent bonds• Rigid like ionic crystal lattices (crystalline structures)

• Bonded by electrostatic attraction between shared, localized valence electrons and positive nucleus.

• Since e- are localized– Poor conductors of heat– Insulators (do not conduct electricity)– Rigid (will break before it changes shape)

Covalent network structures• Diamond, graphite, the element silicon and silicon dioxide

are examples

• Diamond and Silica (glass)

Covalent network structures

• These covalent network structures need a large amount of energy to break all of the strong covalent bonds in them, so these substances have very high melting points.

• Unlike the ionic lattices, the covalent networks are insoluble in water - they are not broken apart by trying to dissolve them.

Covalent network structures

• Diamond vs Graphite

• Diamond: crystalline– hard, colorless, electrical insulator– Bonded in 6-membered rings

• Graphite: amorphous-like– slippery, black, electrical conductor– Bonded in sheets of trigonal planar arrangements– Pi bonds between the layers account for the conductivity

Semiconductors

• Silicon network structures– Metallic: temp & conductivity inversely

proportional– Silicon Network Structures: temp & conductivity

directly proportional

• Doping increases conductivity– Replace a small fraction of silicon with elements

containing 1 more valence electron (n-type) or 1 less valence electron (p-type)

Compare Properties• Sulphur is easily ground, it is

brittle, so it is a nonmetal. In fact it is a covalent molecule

• Tin can be flattened but not broken.

• Magnesium is not easily broken up

SummaryMelting/boiling points

Strength of bonds Examples

Ionic High The ionic bond is strong

Sodium Chloride

Covalent Low The covalent bond is strong intramolecular

Weak intermolecular

Water

Covalent network High Covalent bonds are strong

Diamond

Metallic Varies The Strength of the metallic bond varies

Magnesium