Transcript Section 3 - Bonding and Structure powerpoint

Covalent bonding
Learning Objectives
Candidates should be able to:
 describe, including the use of ‘dot-and-cross’
diagrams, covalent bonding, as in hydrogen; oxygen;
chlorine; hydrogen chloride; carbon dioxide;
methane; ethene.
describe covalent bonding in terms of orbital
overlap.
Starter Activity
Forming a bond
Shared pair of electrons
Electron density map
Electron density
maps for a
hydrogen molecule
Representing a covalent bond
This single σ covalent bond can
be simply represented as:
or H – H
Overlap of atomic orbitals
Overlap of two p-orbitals – e.g. F2
Overlap of one s and one p orbital – e.g. HF
Dot-cross diagrams
Oxygen
Carbon dioxide
Chlorine
Methane
Hydrogen
chloride
Ethene
Breaking the octet rule
Unexpected
structures !!
Bonding in CH4 – promotion of an
electron
C 1s2 2s2 2p2
Bonding in CH4 – hybridisation
Hybridisation in PCl5
Co-ordinate
bonding
Learning Objectives
Candidates should be able to describe, including the use
of ‘dot-and-cross diagrams, co-ordinate (dative covalent)
bonding, as in the formation of the ammonium ion and in
the Al2Cl6 molecule.
A co-ordinate bond (also called a dative covalent bond) is
a covalent bond (a shared pair of electrons) in which both
electrons come from the same atom.
Starter Activity
Dot-cross diagrams
H2S
SiCl4
CH3OH
SF6
CO
Reaction between NH3(g) and HCl(g)
Reaction between NH3(g) and HCl(g)
Reaction between H2O and HCl
Electron-deficient BF3
Aluminium chloride vapour
Aluminium chloride vapour
Electronegativity
Learning Objectives
Candidates should be able to explain the origin of polar
bonds, with reference to electronegativity differences
between atoms.
Electronegativity is a measure of the tendency of an
atom to attract a bonding pair of electrons.
Electronegativity values
Starter Activity
Why are bonds like
bears?
…’cos some of them are POLAR!
Electronegativity differences
What
happens
if
two
atoms
electronegativity bond together?
What
happens
if
B
electronegative than A?
is
of
slightly
equal
more
Representing polar bonds
Representing polar bonds
Molecule
Electronegativity
difference
Dipole/debye
HCl
0.9
1.03
HBr
0.7
0.78
HI
0.4
0.38
Large electronegativity difference
What happens if B is a lot more electronegative
than A?
An ionic bond is formed!!!
Type of bond
As a rough guide:
Type of bond
Non-polar Covalent
Polar covalent
Ionic
Electronegativity
difference
 0.5
Between 0.5 and 1.7
 1.7
Ionic with increasing covalent
character
Polarisation of Anions
Truly ionic
Positive cation
Negative anion
Ionic with some
covalent
character
Polarisation of Anions
Property
Charge
Radius
Cation is the most
powerful polarising
agent when....
Anode is most easily
polarised when...
High
High
Small
Large
Do polar bonds make polar
molecules?
CCl4 molecule is tetrahedral - the partial negative
charges on the Cl atoms are distributed pretty
symmetrically around the molecule. The partial positive
charge on the C is buried in the center of the molecule.
The most electronegative element is
fluorine.
If you remember that fact, everything becomes easy,
because electronegativity must always increase towards
fluorine in the Periodic Table.
Trends in electronegativity
Ionic bonding
Learning Objectives
Candidates should be able to describe
ionic (electrovalent) bonding, as in sodium
chloride and magnesium oxide, including
the use of ‘dot-and-cross’ diagrams.
“The name’s Bond, Ionic Bond
– taken, not shared!!!
Starter Activity
Approaching atoms
Ionic bonding
Ionic bonding
MgO
K2O
CaCl2
Al2O3
Why CaCl2 and not CaCl or CaCl3?
Why CaCl2 and not CaCl or CaCl3?
Polarisation of Anions
Truly ionic
Positive cation
Negative anion
Ionic with some
covalent
character
Polarisation of Anions
Property
Charge
Radius
Cation is the most
powerful polarising
agent when....
Anode is most easily
polarised when...
High
High
Small
Large
Shapes of
molecules
Learning Objectives
Candidates should be able to explain the shapes of,
and bond angles in, molecules by using the model of
electron-pair repulsion.
Starter Activity
Starter Activity
180o
Linear
BeCl2
Balloon molecules
2 Balloons give a linear geometry
3 Balloons give a trigonal planar geometry
4 Balloons give a tetrahedral geometry
5 Balloons give a trigonal bipyramidal geometry
6 Balloons give an octahedral geometry
Shapes of molecules
Electrons
in Electrons
No. of pairs No.
of
Diagram of
outer shell of added
from of electrons
bonding pairs
molecule
central atom
other
atoms
(including bond
(and
any
angles)
charge on an
ion)
Description of
shape
BF3
3
3
3
3
Trigonal
planar
CCl4
4
4
4
4
Tetrahedral
NH3
5
3
4
3
Trigonal
pyramidal
H2O
6
2
4
2
Bent (or Vshaped)
Shapes of molecules
SF6
6
6
6
6
Octahedral
CO2
4
4
2
2
Linear
C2H6
4
4
4
4
Tetrahedral
C2H4
4
4
3
3
Trigonal
planar
ClF4-
7
5
6
4
Square
planar
Metallic bonding
Lesson Objectives
Candidates should
be able to describe
metallic bonding in
terms of a lattice
of
positive
ions
surrounded
by
mobile electrons.
Starter Activity
Metallic Bonding
This is sometimes described as
"an array of positive ions in a sea
of electrons".
Close packed structures?
Dense metals
Group 1 metals
Malleability
Metal grains
Intermolecular
forces
Lesson Objectives
Candidates should be able to describe intermolecular
forces (van der Waals’ forces), based on permanent and
induced dipoles, as in CHCl3(l), Br2(l) and the liquid noble
gases.
Starter Activity
Particles in solids,
liquids and gases
In a liquid or a solid there
must be forces between the
molecules causing them to be
attracted to one another,
otherwise they would move
apart from each other and
become a gas.
Intermolecular forces
Intermolecular attractions are attractions between one
molecule and a neighbouring molecule.
How do intermolecular (or van der
Waals) forces arise?
Temporary or instantaneous dipoles
The lozenge-shaped diagram represents a small
symmetrical molecule - H2, perhaps, or Br2. The even
shading shows that on average there is no electrical
distortion (i.e. the molecule is non-polar).
Temporary or instantaneous dipoles
Electrons are mobile. The constant
"sloshing around" of the electrons in
the molecule causes rapidly fluctuating
dipoles even in the most symmetrical
molecule.
It even happens in monatomic molecules
- molecules of noble gases, like helium,
which consist of a single atom.
Temporary dipole - induced dipole interaction
How temporary dipoles give rise to
intermolecular attractions
van der Waals’ forces
This diagram shows how a whole lattice of molecules could
be held together in a solid using van der Waals’ forces.
An instant later, of course, you would have to draw a quite
different arrangement of the distribution of the
electrons as they shifted around - but always in
synchronisation.
van der Waals lattice
How molecular size affects the
strength of the dispersion forces
The boiling points of the noble gases are:
helium
-269°C
neon
argon
krypton
xenon
radon
-246°C
-186°C
-152°C
-108°C
-62°C
How molecular size affects the
strength of the dispersion forces
There is a gradual increase in the very low boiling
temperatures of the noble gases with increasing atomic
size.
As the size of the atoms increases the number of
electrons increases and the magnitude of the van der
Waals forces increases.
How molecular shape affects the strength of
the temporary dipole interactions
Butane has a higher boiling point because the
intermolecular forces are greater. The molecules are longer
and can lie closer together than the shorter, fatter 2methylpropane molecules.
Permanent dipoles
Permanent dipoles
Hydrogen bonding
Lesson Objectives
Candidates should be able to describe hydrogen
bonding, using ammonia and water as simple examples of
molecules containing N-H and O-H groups.
Starter Activity
Boiling points of the Group 4 hydrides
The increase in boiling point happens because the
molecules are getting larger with more electrons, and so
van der Waals forces become greater.
Boiling points of the hydrides in Groups
5, 6 and 7.
The origin of hydrogen bonding
The origin of hydrogen bonding
Hydrogen
bonding
is
a
particularly
strong
intermolecular force that involves three features:
a large dipole between an H atom and the highly
electronegative atoms N, O or F;
the small H atom which can get very close to other
atoms;
a lone pair of electrons on another N, O or F, with
which the positively charge H atom can line up.
Drawing hydrogen bonds
1 mark for showing
H-bond
1 mark for showing
lone pair
1 mark for indicating
bond polarity
Hydrogen bonding in water
Hydrogen bonding accounts for
many of the other unusual
properties of water including:
its high specific heat capacity
its very high surface tension
its high viscosity and
the low density of ice compared
to water
Which type of intermolecular force?
A summary
Bonding,
structure and
properties
Lesson Objectives
Candidates should be able to:
describe, interpret and/or predict the effect of
different types of bonding on the physical properties
of substances.
describe, in simple terms, the lattice structure of a
crystalline solid which is ionic, simple molecular, giant
molecular, hydrogen-bonded and metallic.
suggest from quoted physical data the type of
structure and bonding present in a substance.
Starter Activity
Bonding, structure and properties
The properties of substances are decided by their bonding
and structure.
Bonding means the way the particles are held
together: ionic, covalent, metallic or weak intermolecular
bonds.
Structure means the way the particles are arranged
relative to one another. You have already met the major
types of structure at IGCSE.
Bonding, structure and properties
Bond energies
Bond
Average
bond Bond length/nm
enthalpy/kJmol-1
C – C
+347
0.154
C = C
+612
0.134
C ≡ C
+838
0.120
C – H
+413
0.108
O – H
+464
0.096
C – O
+358
0.143
C = O
+805
0.116
GIANT LATTICE
COVALENT MOLECULAR
Ionic
Covalent network
Metallic
What substances have Compounds of metals Some
elements
in Metals
this
type
of with non-metals.
Group 4 and some of
structure?
their compounds.
Examples
What type of particle
does it contain?
How are the particles
bonded together?
NaCl
ions
SiO2
atoms
Strong
ionic
bonds; attraction
between
oppositely charged
ions
Strong
covalent
bonds; attraction
of atoms’ nuclei
for
shared
electrons
high
very high
generally high
low
hard but brittle
hard but
malleable
conduct when
(s) or (l)
soft
conduct when
(l) or (aq)
very hard (if
3D)
do not normally
conduct
often soluble
insoluble
insoluble (but
insoluble
insoluble
insoluble
What are the typical properties?
M. pt and b.pt.
Hardness
Electrical conductivity
Solubility in water
Solubility in non-polar
solvents (e.g. hexane)
Cu
positive ions
and delocalised
electrons
Simple molecular
Macromolecular
Some
non-metal Polymers
elements and usually
some non-metal/nonmetal compounds.
H2 O
molecules
Poly(ethene)
molecules
Strong
metallic Weak intermolecular bonds between
bonds; attraction molecules; strong covalent bonds
of atoms’ nuclei between atoms within each molecule.
for
delocalised
electrons
some react)
moderate (often
decompose on heating)
variable
do not conduct
do not normally
conduct
sometimes
usually insoluble
(but some H-bond)
soluble
sometimes
usually soluble
soluble
Structure table
The modern use of
materials
Lesson Objectives
Candidates should be able to:
Explain the strength, high melting point and
insulating properties of ceramics in terms of their
giant molecular structure.
Relate the uses of ceramics to their properties.
Describe and interpret the uses of the metals
aluminium and copper (and their alloys) in terms of
their physical properties.
Understand that materials are a finite resource and
the importance of recycling processes.
Starter Activity
Five most common metals
Aluminium
Brass
Zinc
Copper
Steel
Uses of aluminium
Low density, corrosion resistance
and strength make it ideal for
construction of aircraft, lightweight
vehicles, and ladders.
Malleability, low density, corrosion
resistance
and good thermal
conduction make it a good material for
food packaging.
Good
electrical
conduction,
corrosion resistance and low density
leads to its use for overhead power
cables hung from pylons (low density
gives it an advantage over copper).
Uses of copper
Copper
is
an
excellent
conductor of electricity and heat.
Copper is soft and malleable.
Copper is very unreactive and
therefore corrosion resistant.
Alloys of copper
Copper Alloy Other metal
it contains
Brass
Zinc
Bronze
Tin
Main
properties
Uses
Fairly soft
and malleable
Screws and
hinges
Strong
Propellors
and bearings
Ceramics
A mineral is a naturally occurring solid
formed through geological processes that
has
a
characteristic
chemical
composition, a highly ordered atomic
structure,
and
specific
physical
properties.
Ceramic: Any of various hard,
brittle,
heat-resistant
and
corrosion-resistant materials made
by shaping and then firing a
nonmetallic mineral.
Ceramics
furnace: an enclosed chamber in which heat is produced to
heat buildings, destroy refuse, smelt or refine ores, etc.
heat shields
glass and crockery
brake pads
furnace linings
electrical insulators
Recycling
Raw materials extracted (removed by chemical means) from the Earth cannot last
forever. Although some materials are more (present in great quantity) abundant than
others, they are all finite (have a limit) resources. Increasing demand for raw
materials (items used to produce something else), coupled with ever growing problems
of waste disposal, have led to considerable interest in recycling (processing for reuse)
waste.
Recycling has a number of possible advantages (beneficial factors):
It leads to reduced demand for new raw materials;
It leads to a reduction in environmental damage (harm to the
surroundings);
It reduces the demand for landfill sites (a place for burying waste) to dump
waste;
It reduces the cost of waste disposal;
It may reduce energy costs.
The kinetic-molecular
model of liquids
Lesson Objectives
Candidates should be able to describe using a kineticmolecular model, the liquid state; melting; vaporisation
and vapour pressure.
Starter Activity
The Kinetic-Molecular Model of Liquids
Solid
Arrangement of very orderly
particles
Liquid
Gas
almost complete
disorder
high
Movement of
particles
vibrate about
fixed positions
Proximity of
particles
Compressibility
of substance
Conduction of
heat
close (~10-10m)
short-range
order, longer
range disorder
some movement
from place to
place
close (~10-10m)
very low
very low
poor except
metals and
graphite
metals very good; very poor
others poor
continuous, rapid,
random movement
far apart (~10-8m)
The Kinetic-Molecular Model of Liquids
Liquids do not have a fixed __________ because the
particles can move about. However, they remain very
__________ together. This shows that the inter-particle
forces have not been __________ broken. If sufficient
__________ is supplied, the particles overcome the interparticle forces almost completely and __________ from
the liquid. This is called __________ or boiling. The
energy required to boil a liquid is always __________ than
that required to melt the same substance and is a better
__________ of the strength of inter-particle forces.
Vapour pressure
Vapour pressure is the pressure of a vapour over a liquid at
equilibrium.
Vapour pressure
Even at low temperature there are particles with
high energy.
Vapour pressure
At equilibrium, the rate at which molecules leave the
liquid equals the rate at which molecules join the liquid.
Measuring vapour pressure
More on ideal gases
Lesson Objectives
Candidates should be able to explain qualitatively in
terms of intermolecular forces and molecular size the
limitations of ideality at very high pressures and very low
temperatures.
Starter Activity
More on ideal gases
More on ideal gases
Y is hydrogen. It is closest to ideal under all conditions.
Hydrogen has the weakest intermolecular forces and is
the smallest molecule.
Z is ammonia. It is the least ideal at lower pressures.
Ammonia molecules can hydrogen bond.
X is nitrogen. Deviates greatly from ideality at high
pressures where its larger molecular volume becomes
important.