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 worlds
  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/essentialche
  mistry/flash/molvie1.swf

21
(No Transcript)
22
(No Transcript)
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
Covalent bonding properties

• Giant covalent
• Ex diamond, silicon dioxide
• Very hard
• Very high MP (gt1000oC)
• Does not conduct
• Insoluble in all solvents

• Molecular covalent
• Ex CO2, alcohols, I2
• Usually soft, malleable
• Low MP (lt200oC)
• Does not conduct
• More soluble in non-aqueous solvents, unless they
  can h-bond

26
Solubility of methanol in water

• http//www.mhhe.com/physsci/chemistry/animations/c
  hang_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 its 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