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Ionic, Covalent and Metallic structures of solids.

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1 Ionic, Covalent and Metallic structures of solids

2 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

3 Vocab: Crystalline Solid: –Lattice: –Unit Cell: Amorphous solid: Alloy: Conductivity: Boiling Point:

4 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

5 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.

6 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)

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8 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

9 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

10 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

11 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).

12 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.

13 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)

14 Covalent network structures Diamond, graphite, the element silicon and silicon dioxide are examples Diamond and Silica (glass)

15 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.

16 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

17 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)

18 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

19 Summary Melting/boiling points Strength of bondsExamples IonicHighThe ionic bond is strong Sodium Chloride CovalentLowThe covalent bond is strong intramolecular Weak intermolecular Water Covalent networkHighCovalent bonds are strong Diamond MetallicVariesThe Strength of the metallic bond varies Magnesium


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