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.

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

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

Physical PropertiesConductivity 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)

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)

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!

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)

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

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.

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 malleableAll metals conduct electricity conduct

electricity in the solid state.

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

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 propertiesSolubilities of compounds in polar and non-polar solventsSolubilities of alcohols in water being related to chain length

General physical properties

Depend on the forces between the particles

The stronger the bonding between the particles, the higher the M.P and BPMP tends to depend on the existence of a

regular lattice structure

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

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

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

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)

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.

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.

Soluble salt in water breaks up as

NaCl (s) Na+ (aq) + Cl- (aq)

http://www.mhhe.com/physsci/chemistry/essentialchemistry/flash/molvie1.swf

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.

Molten or dissolved ionic compounds conduct electricity

Insoluble in most solvents, yet H2O is polar and attracts both the + and – ions from salts

Covalent bonding propertiesGiant 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

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.

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.

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