Transcript 4.4 Metallic bonding

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
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Occurs between
atoms with low
electronegativities
Metal atoms pack
close together in 3-D,
like oranges in a box.
Close-packed lattice
formation
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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 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)
Malleability
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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
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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
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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.
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They are generally less malleable and ductile
Some alloys are made by melting and mixing
two or more metals
 Bronze
 Steel =
= copper and zinc
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
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Who needs it?
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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:
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Use the commonly accepted model of
metal bonding to explain why:
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.
 The
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Reading on pages 369-371
 Page
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 properties
Solubilities of compounds in polar and non-polar solvents
Solubilities 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 BP
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 MP
tends to depend on the existence of a
regular lattice structure
Impurities and Melting points
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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
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A qualitative measure of how readily a liquid or
solid is vaporised upon heating or evaporation
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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)
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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
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 C12H22O11(s)  C12H22O11(aq)
Dissolving sugar (covalent
compound)
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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.
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Soluble salt in water breaks up as
NaCl (s)  Na+ (aq) + Cl- (aq)
http://www.mhhe.com/physsci/chemistry/essentialchemis
try/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.
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Molten or dissolved ionic compounds
conduct electricity
 Insoluble in most solvents, yet H2O is
polar and attracts both the + and – ions
from salts
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Covalent bonding properties
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 nonaqueous solvents,
unless they can hbond
Solubility of methanol in water
http://www.mhhe.com/physsci/chemistry/a
nimations/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.
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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.
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http://www.chemguide.co.uk/atoms/bonding/metall
ic.html
 Metallic
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bonding review
http://chemed.chem.purdue.edu/genchem/topicrev
iew/bp/ch18/soluble.php
 Solubility
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review
http://wwwcsi.unian.it/educa/inglese/kevindb.html
 History
involved with dissolving ionic compounds