4.4 Metallic bonding

1. 4.4 Metallic bonding 4.4.1 Describe metallic bond as the electrostatic attraction between a lattice of positive ions surrounded by delocalized valence electrons. 4.4.2 Explain the electrical conductivity and malleability of metals Students should appreciate the economic importance of these properties and the impact that the large-scale production of iron and other metals has made on the world.

2. Metallic bond • Occurs between atoms with low electronegativities • Metal atoms pack close together in 3-D, like oranges in a box. • Close-packed lattice formation

3. Many metals have an unfilled outer orbital • In an effort to be energy stable, their outer electrons become delocalised amongst all atoms • No electron belongs to one atom • They move around throughout the piece of metal. • Metallic bonds are not ions, but nuclei with moving electrons

4. Physical Properties Conductivity • Delocalised electrons are free to move so when a potential difference is applied they can carry the current along • Mobile electrons also mean they can transfer heat well • Their interaction with light makes them shiny (lustre)

5. Malleability • The electrons are attracted the nuclei and are moving around constantly. • The layers of the metal atoms can easily slide past each other without the need to break the bonds in the metal • Gold is extremely malleable that 1 gram can be hammered into a sheet that is only 230 atoms thick (70 nm)

6. Melting points • Related to the energy required to deform (MP) or break (BP) the metallic bond • BP requires the cations and its electrons to break away from the others so BP are very high. • The greater the amount of valence electrons, the stronger the metallic bond. • Gallium can melt in your hand at 29.8 oC, but it boils at 2400 oC!

7. Alloys • Alloying one metal with other metal(s) or non metal(s) often enhances its properties • Steel is stronger than pure iron because the carbon prevents the delocalised electrons to move so readily. • If too much carbon is added then the metal is brittle. • They are generally less malleable and ductile • Some alloys are made by melting and mixing two or more metals • Bronze = copper and zinc • Steel = iron and carbon (usually)

8. Economic importance • Iron is found by certain percentages in minerals, such as iron oxides like of magnetite (Fe3O4), hematite (Fe2O3), and many others. • Hematite- up to 66% pure could be put in a blast furnace directly for the production of iron metal • 98% of iron production is destined for making steel

9. Who needs it? • China, then Japan, then Korea are the world’s largest consumer's of iron Where does it come from? • Iron rich minerals are commonly found everywhere in the world, however China, Brazil and Australia are the highest producers of iron ore mining • The main constraint is the position of the iron ore relative to market, the cost of rail infrastructure to get it to market and the energy cost required to do so.

10. Exercise: • Use the commonly accepted model of metal bonding to explain why: • The boiling points of metals in the 3rd period increase from sodium to magnesium to aluminum. • Most metals are malleable • All metals conduct electricity conduct electricity in the solid state.

11. Reading on pages 369-371 • Page 375 # 9.70, 9.74, 9.72

12. 4.5 Physical Properties 4.5.1 Compare and explain the following properties of substances resulting from different types of bonding: melting and boiling points, volatility, conductivity and solubility. Look at how impurities affect these properties Solubilities of compounds in polar and non-polar solvents Solubilities of alcohols in water being related to chain length

13. General physical properties • Depend on the forces between the particles • The stronger the bonding between the particles, the higher the M.P and BP • MP tends to depend on the existence of a regular lattice structure

14. Impurities and Melting points • An impurity disrupts the regular lattice that its particle adopts in the solid state, so it weakens the bonding. • They always LOWER melting points • Its often used to check purity of a known molecular covalent compound because its MP will be off, proving its contamination

15. How would this ideal heat curve look different if the substance was contaminated?

16. Volatility • A qualitative measure of how readily a liquid or solid is vaporised upon heating or evaporation • It is a measure of the tendency of molecules and atoms to escape from a liquid or a solid. • Relationship between vapour pressure and temperature (B.P) • Mostly dealing with liquids to gas, however can occur from solid directly to gas (dry ice). • The weaker the intermolecular bonds, the more volatile

17. Conductivity • Generally molecules have poor solubility in polar solvents like water, but if they do dissolve they do not for ions • There are no charged particles to carry the electrical charge across the solution. • Example: sugar dissolves in water • C12H22O11(s) C12H22O11(aq)

18. Dissolving sugar (covalent compound) • It takes energy to break the bonds between the C12H22O11 molecules in sucrose crystal structure. • It also takes energy to break the hydrogen bonds in water so that one of these sucrose molecules can fit into solution. • In order for sugar to dissolve, there must be a greater release of energy when the dissolution occurs than when the breaking of bonds occur.

19. Ionic compounds • The energy needed to break the ionic bond must be less than the energy that is released when ions interact with water. • The intermolecular ion-dipole force is stronger than the electrostatic ionic bond • Breaks up the compound into its ions in solution.

20. Soluble salt in water breaks up as NaCl (s) Na+ (aq) + Cl-(aq) • http://www.mhhe.com/physsci/chemistry/essentialchemistry/flash/molvie1.swf

23. Ionic compounds • Held together by strong 3-d electrostatic forces. • They are solid at room temperature and pressure • If one layer moves a fraction, the ions charges are off and now repulsion occurs. This is the reason they are strong, yet brittle.

24. Molten or dissolved ionic compounds conduct electricity • Insoluble in most solvents, yet H2O is polar and attracts both the + and – ions from salts

25. Giant covalent Ex: diamond, silicon dioxide Very hard Very high MP (>1000oC) Does not conduct Insoluble in all solvents Molecular covalent Ex: CO2, alcohols, I2 Usually soft, malleable Low MP (<200oC) Does not conduct More soluble in non-aqueous solvents, unless they can h-bond Covalent bonding properties

26. Solubility of methanol in water • http://www.mhhe.com/physsci/chemistry/animations/chang_7e_esp/clm2s3_4.swf • Alcohols generally become less soluble, the longer the carbon chain due to the decreasing tendency for hydrogen bonding to occur intermolecularly.

27. States of matter • Physical state depends on intermolecular forces • The weaker the attraction, the more likely it’s a gas, while stronger attractions indicate solid.

28. http://www.chemguide.co.uk/atoms/bonding/metallic.html • Metallic bonding review • http://chemed.chem.purdue.edu/genchem/topicreview/bp/ch18/soluble.php • Solubility review • http://wwwcsi.unian.it/educa/inglese/kevindb.html • History involved with dissolving ionic compounds