Download File - Mr Weng`s IB Chemistry

IB CHEMISTRY
Topic 4 Bonding
Higher level
4.1 Ionic bonding and structure
OBJECTIVES
• Positive ions (cations) form by metals losing valence electrons.
• Negative ions (anions) form by non-metals gaining electrons.
• The number of electrons lost or gained is determined by the electron
configuration of the atom.
• The ionic bond is due to electrostatic attraction between oppositely
charged ions.
• Under normal conditions, ionic compounds are usually solids with
lattice structures.
• Deduction of the formula and name of an ionic compound from its
component ions, including polyatomic ions.
• Explanation of the physical properties of ionic compounds (volatility,
electrical conductivity and solubility) in terms of their structure.
Ionic bonding
Ionic Bond – electrostatic attraction between oppositely
charged ions
• Ions with a negative charge are called anions.
• Ions with a positive charge are called cations.
Ionic bonding – other points
• Non directional bond - strength of bond equal in all
directions
• Conducts electricity when molten or in solution (aq)
• High melting point and boiling point hard solids
• Low volatility (tendency of a substance to vaporize)
• Brittle
Formation of ions - cations
• Elements with 1, 2, or 3 valence electrons (metals or
‘losers’) react chemically by losing their valence electrons
and become positive ions (cations)
• Element takes on the noble gas electron
configuration/structure.
• Example: Na has one valence electron but once removed
it’s electron structure becomes that of Ne.
2.8.1  2.8
or
1s22s22p63s1  1s22s22p6
Lewis structure - formation of sodium ion
Sodium has 1 electron in its outer shell
If it loses this it will have no partially filled shells.
Na
Sodium atom (2.8.1)
Loses 1
electron
Na
Sodium 1+ ion (2.8.0)
Formation of ions - cations
Group 1
+1
Li+ Na+ K+
Cations
Group 2
+2
Mg2+ Ca2+
Group 3
+3
Al3+
Formation of ions - anions
• Elements with 5, 6, or 7 valence electrons (non-metals) react
chemically by gaining enough electrons to become negative ions
(anions) and take on the next noble gas electron structure.
• Result: Element takes on the next noble gas electron
configuration/ structure.
• Example: Cl has seven valence electrons but once one is gained
it’s electron structure becomes that of Ar.
2.8.7  2.8.8
or
1s22s22p63s23p5  1s22s22p63s23p6
Lewis structure - formation of chloride ion
Chlorine has 7 electrons in its outer shell.
If it gains 1 electron it can achieve a full outer electron
shell. It is, therefore, going to be able to accept the
electron that the sodium wants to lose.
Cl
Chlorine atom (2.8.7)
Gains 1 electron
(from sodium)
Cl
Chlorine 1- ion (2.8.8)
Formation of ions - anions
Group 5
-3
N3- P3-
Anions
Group 6
-2
O2- S2-
Group 7
-1
F- Cl- Br-
Formation of ions – transition metals
• Transition elements can form more than one ion.
• Example: Fe can form Fe2+ and Fe3+
• Ground state: 1s22s22p63s23p64s23d6
• Loss of 2 e-s: 1s22s22p63s23p63d6
• Loss of 3 e-s: 1s22s22p63s23p63d5
Ionic bond formation
Use the periodic table (metal + non-metal) or from the
electronegativity values in your data booklet.
The degree of polarity can be determined by looking at the
difference in electronegativity.
Roughly it breaks down like this:
0.0 – 0.4
Nonpolar covalent bond
0.4 – 1.8
Polar covalent bond
>1.8
Ionic bond
Electronegativity and ionic bonding
• The ability of atoms in a
molecule to attract
electrons to itself.
• On the periodic chart,
electronegativity increases
as you go…
– …from left to right across a
row.
– …from the bottom to the top
of a column.
Some polyatomic ions
• Common polyatomic ions formed by non-metals in period 2
and 3 that you must know!
Chemical name
Chemical formula
nitrate
phosphate
hydroxide
ammonium ion
sulfate
hydrogen carbonate
carbonate
NO3PO43OHNH4+
SO42HCO3CO32-
Describing the lattice structure
• Example: Small sodium chloride ionic lattice
Describing the lattice structure
• Endlessly repeating lattice of ions…depends on the
size of the crystal lattice...although arrangement of
ions stays the same
• The coordination number is (6,6) meaning each Na+
is surround by 6 Cl-, each Cl- is surrounded by 6 Na+.
4.2. Covalent bonding
OBJECTIVES
• A covalent bond is formed by the electrostatic attraction
between a shared pair of electrons and the positively charged
nuclei.
• Single, double and triple covalent bonds involve one, two and
three shared pairs of electrons respectively.
• Bond length decreases and bond strength increases as the
number of shared electrons increases.
• Bond polarity results from the difference in electronegativities
of the bonded atoms.
• Deduction of the polar nature of a covalent bond from
electronegativity values.
Covalent Bonding
A covalent bond is the
electrostatic attraction
between a shared pair
of electrons and
positively charged
nuclei
• These electrons are called
bonding electron pairs.
• Electrons not involved in
bonding are called nonbonding electron pairs.
100% covalent
100% ionic
Bonding spectrum
+
A B
A B
A
Increasing DEN
Increasing polarity
Transfer
-
B
Problem 1: Which is more ionic, NaCl, or LiCl?
Δχ(NaCl) =│χa - χb│
= 0.9 – 3.2 = 2.3
Ʃχ(NaCl) = (χa + χb)/2
= 4.1/2 = 2.05
Δχ(LiCl) =│χa - χb│
= 1.0 – 3.2 = 2.2
Ʃχ(LiCl) = (χa + χb)/2
= 4.2/2 = 2.1
Therefore NaCl is more
ionic.
POLARITY
Polarity
• The
partial
positive
charge is written as δ+
and the partial negative
charge is written as δ-.
• The overall charge is
called a dipole moment
and is represented by a
vector as shown
opposite.
Polarity
• Just because a molecule
possesses polar bonds it
does not mean that the
molecule as a whole will
be polar.
Polarity
The molecule is polar if the electron densities are
not symmetrical such that there is a net dipole
moment.
Polarity
BONDS STRENGTHS
Bonds strengths
• Each bond contains 2 electrons.
• Triple bonds are shorter and stronger than double
bonds.
• Double bonds are shorter and stronger than single
bonds.
<
<
Bonds strengths
A strong attraction between two nuclei for the electrons
creates a short bond length. Therefore short bond lengths
have high bond dissociation enthalpy.
Type of
bonding
Nature of
bonding
Strength of
bonds
Ionic
Electrostatic attraction
between positive and
negative ions.
Covalent
Electrostatic
attraction of a
shared pair of
electrons between
atoms.
The smaller the ions
The shorter the
and the greater the
bond, the stronger
charge on the ions, the the bond (usually).
stronger the attraction
Double bonds are
between the ions
stronger than single
(usually). This is due to bonds, while triple
a greater charge
bonds are stronger
density within the
than double bonds.
structure.
Metallic
Electrostatic attraction
between lattice of positive
metal ions and delocalised
outer shell electrons.
The smaller the metal ions, the
greater the charge on the ions,
and the more delocalised outer
shell electrons there are, the
stronger the attraction between
the ions and electrons
(usually). This is due to a
greater charge density within
the structure.
4.3 Covalent structures
OBJECTIVES
• Lewis (electron dot) structures show all the valence electrons in a covalently bonded
species.
• The “octet rule” refers to the tendency of atoms to gain a valence shell with a total of 8
electrons.
• Some atoms, like Be and B, might form stable compounds with incomplete octets of
electrons.
• Resonance structures occur when there is more than one possible position for a
double bond in a molecule.
• Shapes of species are determined by the repulsion of electron pairs according to
VSEPR theory.
• Carbon and silicon form giant covalent/network covalent structures.
• Deduction of Lewis (electron dot) structure of molecules and ions showing all valence
electrons for up to four electron pairs on each atom.
• The use of VSEPR theory to predict the electron domain geometry and the molecular
geometry for species with two, three and four electron domains.
• Prediction of bond angles from molecular geometry and presence of non-bonding
pairs of electrons.
• Prediction of molecular polarity from bond polarity and molecular geometry.
• Deduction of resonance structures, examples include but are not limited to C6H6, CO32and O3 .
• Explanation of the properties of giant covalent
LEWIS STRUCTURES
Covalent Bonding – Lewis structures
• Lewis structures are representations of molecules
showing all electrons, bonding and nonbonding.
• Each bond contains 2 electrons.
• Elements pair up according to the octet rule – elements
gain or lose electrons in order to have the electron
configuration of a noble gas.
Cl2
How to draw Lewis structures
deductively
1. Draw out the atoms with the least occurring atom as
the central atom (if applicable).
2. Draw in the valence electrons.
3. Connect up one bond from each of the outside atoms
to the central atom.
4. Count up the electrons to see if all the atoms follow the
octet rule (here oxygens only have 7).
5. If not, then add extra bonds until they do.
6. If this cannot be done, a coordinate covalent bond –
where one atom donates both electrons to the bonding
pair, may need to be added. Eg. CO
(Free radicals are the other option)
7. Clean up the structure making it easy to read (electrons
in pairs, o and x).
Structural formula:
Structural formula:
O=O
Exceptions to the Octet: Boron
Boron forms stable compounds with just 3
valence electrons
These compounds are highly reactive:
Exceptions to the Octet: Beryllium
Beryllium forms stable compounds with just 2
valence electrons
Compounds that break the octet rule are often toxic and dangerous, ready to
react so the octet rule is obeyed!
O2
N2
CO2
HCN
O
C
O
C2H4 (ethene)
NH4+
CO
Draw this one now
H 3O +
C2H2 (ethyne)
Draw this one now
or
VSEPR THEORY
Shapes of molecules
• Determined by number of valence electrons of
the central atom
• 3-D shape a result of bonded pairs and lone
pairs of electrons
• VSEPR theory (valence-shell-electron-pair
repulsion) states the best arrangement of a
given number of electron domains (negative
charge centres) is the one that minimizes the
repulsions among them.
Shapes of molecules
• Simply put, electron pairs, whether they be
bonding or nonbonding, repel each other.
• By assuming the electron pairs are placed as
far as possible from each other, we can predict
the shape of the molecule.
Shapes – electron domains
• An electron domain is the number of
bonds/electron pairs there are on the central
atom
Eg. SO2 has 3 electron domains
CH4, NH3, NH4+ all have 4 electron domains
Using the VSEPR Model
1. Draw the electron-dot structure
2. Identify the central atom
3. Count the total number of electron pairs around
central atom
4. Predict the electron shape
5. Predict the shape of the molecule using the
bonding atoms
Rule of thumb: Let each non-bonding electron pair bend by a further 2-3⁰
Shapes (if no non-bonding electron pairs)
Molecules, or ions, possessing ONLY BOND PAIRS of
electrons fit into a set of standard shapes. All the bond
pair-bond pair repulsions are equal.
All you need to do is to count up the number of bond
pairs and chose one of the following examples...
BOND
PAIRS
SHAPE
C
A covalent bond will repel
another covalent bond
BOND
ANGLE(S)
EXAMPLE
2
LINEAR
180º
BeCl2
3
TRIGONAL PLANAR
120º
AlCl3
4
TETRAHEDRAL
109.5º
CH4
5
TRIGONAL BIPYRAMIDAL
90º & 120º
PCl5
6
OCTAHEDRAL
90º
SF6
Shapes (with non-bonding electron pairs)
If a molecule, or ion, has lone pairs on the central atom, the shapes are slightly
distorted away from the regular shapes. This is because of the extra repulsion
caused by the lone pairs.
BOND PAIR - BOND PAIR
O
<
LONE PAIR - BOND PAIR
O
<
LONE PAIR - LONE PAIR
O
As a result of the extra repulsion, bond angles tend to
be slightly less as the bonds are squeezed together.
Methane CH4
H
H
H
C
C
H
H
Carbon - has four electrons to pair up
Four covalent bonds are formed
Hydrogen - 1 electron to complete shell
C and H now have complete shells
BOND PAIRS
4
LONE PAIRS
0
H
109.5°
C
BOND ANGLE...
SHAPE...
109.5°
TETRAHEDRAL
H
H
H
Ammonia NH3
H
N
H
H
N
H
BOND PAIRS
3
LONE PAIRS
1
TOTAL PAIRS
4
• Nitrogen has five electrons in its outer shell
• It cannot pair up all five - it is restricted to eight electrons in its outer shell
• It pairs up only three of its five electrons
• 3 covalent bonds are formed and a pair of non-bonded electrons is left
• As the total number of electron pairs is 4, the shape is BASED on a tetrahedron
Ammonia NH3
H
H
N
H
H
N
BOND PAIRS
3
LONE PAIRS
1
TOTAL PAIRS
4
• The shape is based on a tetrahedron but not all the repulsions are the same
• LP-BP REPULSIONS > BP-BP REPULSIONS
• The N-H bonds are pushed closer together
• Lone pairs are not included in the shape
N
H
H
N
N
H
H
H
H
H
107°
H
H
ANGLE... 107°
SHAPE... PYRAMIDAL
Ammonia NH3
H
N
H
H
N
H
BOND PAIRS
3
LONE PAIRS
1
TOTAL PAIRS
4
Water H2O
H
O
H
H
O
BOND PAIRS
2
LONE PAIRS
2
TOTAL PAIRS
4
• Oxygen has six electrons in its outer shell
• It cannot pair up all six - it is restricted to eight electrons in its outer shell
• It pairs up only two of its six electrons
• 2 covalent bonds are formed and 2 pairs of non-bonded electrons are left
• As the total number of electron pairs is 4, the shape is BASED on a tetrahedron
Water H2O
H
H
O
H
O
BOND PAIRS
2
LONE PAIRS
2
TOTAL PAIRS
4
• The shape is based on a tetrahedron but not all the repulsions are the same
• LP-LP REPULSIONS > LP-BP REPULSIONS > BP-BP REPULSIONS
• The O-H bonds are pushed even closer together
• Lone pairs are not included in the shape
O
H
O
O
H
H
H
H
104.5°
H
ANGLE... 104.5°
SHAPE... ANGULAR
Ammonium NH4+
NH4+
N+
BOND PAIRS
4
LONE PAIRS
0
TETRADHEDRAL
H
H-N-H 109.5°
H
N+
H
H
Hydronium ion H3O+
Bromine triflouride BrF3
F
F
Br
F
F
BOND PAIRS
3
LONE PAIRS
2
’T’ SHAPED
ANGLE <90°
F
Br
F
Sulphur dioxide SO2
S has 2 bonded atoms , 1 lone pair (electron cloud),
120°, bent, V-shaped
.. .. ..
:O:: S:O:
..
S
O
O
Ethane C2H3, Ethene C2H4, Ethyne C2H2
Carbon dioxide CO2
The shape of a compound with a double bond is calculated in the same way.
A double bond repels other bonds as if it was single e.g. carbon dioxide
C
O
O
C
O
Carbon - needs four electrons to complete its shell
The atoms share two electrons
Oxygen - needs two electron to complete its shell
each to form two double bonds
DOUBLE BOND PAIRS
2
LONE PAIRS
0
Double bonds behave exactly as single
bonds for repulsion purposes so the
shape will be the same as a molecule with
two single bonds and no lone pairs.
180°
O
C
O
BOND ANGLE... 180°
SHAPE... LINEAR
By Deduction
RESONANCE STRUCTURES
Resonance
When a Lewis structure allows for the same arrangement of
atoms but a different but equally valid arrangement of
electrons, resonance occurs eg. Ozone, O3
The resonance hybrid is actually the dotted line below.
The electrons are said
to be delocalized, and
are spread over two
bonding orbitals.
These electrons are located in the p orbitals. The
overlapping p orbitals create a new shape called a pi
bond () in a normal bonding situation. In resonance
these  bonds then all overlap as well.
Nitrate NO3
-
When writing Lewis structures for species like the
nitrate ion, we draw resonance structures to more
accurately reflect the structure of the molecule or ion.
• In reality, each of the four
atoms in the nitrate ion has
a p orbital.
• The double bond on N=O is
normally a  bond.
• The p orbitals on all three
oxygens overlap with the p
orbital on the central
nitrogen.
This means the  electrons
are not localized between the
nitrogen and one of the
oxygens, but rather are
delocalized throughout the
ion.
Benzene C6H6
The organic molecule
benzene has a p orbital on
each carbon atom.
• In reality the  electrons in benzene are not
localized, but delocalized.
• The even distribution of the  electrons in
benzene makes the molecule unusually stable.
Carbonate CO32-
CARBON STRUCTURES
Allotropes
Allotropes are compounds of the same element
that differ in structure
• Carbon can bond with itself in at least three
major different inorganic ways giving us 3 very
different materials
1. Diamond
2. Graphite (graphene and pyrolytic carbon)
3. Fullerenes - Buckyballs and nanotubes
Diamond
• Each carbon bonded to 4
other carbons
HL: Carbons are bonded via
sp3 hybridization to 4 other
carbon atoms forming a
giant network covalent
compound.
Properties of Diamond
• High melting point due to strong directional covalent
bonds (3550 C)
• Extremely hard because it is difficult to break atoms
apart or move them in relation to one another
• No electrical conductivity because electrons are
localized in specific bonds
• Insoluble in polar and non-polar solvents because
molecular bonds are stronger than any
intermolecular forces
Graphite
• Each carbon bonded to 3
other carbons forming
sheets of graphene. These
graphene sheets slide over
each other.
HL: Carbon atoms are bonded
via sp2 hybridization.
Carbon atoms form sheets of
six sided rings with p-orbitals
perpendicular from plane of
ring
Graphite Structure
• Carbon has 4 valence eto bond with. 3 are
used for closest atoms
in rings. 1 is
delocalized in p-orbitals
• The presence of porbitals allows for
stronger London forces
that hold the sheets
together
Properties of Graphite
• Different from Diamond
– Conducts electricity because of delocalized electrons,
use as a conductor, lack of vibration between layers
however make it an insulator not a conductor of heat
– Slippery can be used as lubricant, sheets can easily slip
past each other (think of a deck of cards), use as
pencil ‘leads’
• Same as Diamond
– High melting point (higher actually because of
delocalized e-, 3653C), use as carbon fibre
– Insoluble (same reason)
Fullerenes
• Buckyballs: spherical
• Nanotubes: tube shaped
• Both have very interesting
properties
– Super strong
– Conduct electricity and
heat with low resistance
– Free radical scavenger
Buckyballs or Buckminsterfullerene
• Carbon atoms bond in units of
60 atoms (C-60) forming a
structure similar to a soccerball
with interlocking six sided and
five sided rings.
HL: sp2 hybridization
Extra p-orbitals form pi bonds
resulting in...
Properties:
– Electrical conductivity
– Stronger covalent bonds,
therefore stronger
materials
QUART
SILICON STRUCTURES
Silicon dioxide (SiO2)
• SiO2
repeating
unit as every
oxygen links
in with a
second Si
atom
• Structurally
think
tetrahedral,
SiO4
Difference between Si and C
(diamond)
• Si is larger than C so the Si-O bond length is
greater
• The greater the bond length, the lower the
bond enthalpy (energy) is
• This means it is easier to break
• Therefore Si is more reactive than C (diamond)
4.4 Intermolecular forces
OBJECTIVES
• Intermolecular forces include London (dispersion) forces,
dipole-dipole forces and hydrogen bonding.
• The relative strengths of these interactions are London
(dispersion) forces < dipole-dipole forces < hydrogen bonds.
• Deduction of the types of intermolecular force present in
substances, based on their structure and chemical formula.
• Explanation of the physical properties of covalent compounds
(volatility, electrical conductivity and solubility) in terms of their
structure and intermolecular forces.
Intermolecular forces (IMF)
Intermolecular forces – attractions between
molecules that have temporary dipoles, permanent
dipoles or hydrogen bonding
(Van der Waal’s forces include only 1 & 2 interactions)
1. London (dispersion) forces - the interactions
between non-polar molecules from temporary
dipoles (induced dipoles).
2. Dipole-dipole forces – from permanent dipoles of
polar molecules.
3. Hydrogen bonding – from bonds with H and
either N, O or F only.
London Forces
• Weakest of the three
• also called the Dispersion forces (see
animation)
• Caused by the motion of electrons
around a nucleus
• Electrons sometimes are
asymmetrical about an atom, leading
to temporary dipoles also called
instantaneous induced dipoles
• The formation of a dipole in one
molecule can cause an opposite
dipole to form in a nearby molecule
• Dispersion forces increases with
atomic radii
Increasing London forces:
1. Large radius contains more electrons
2. Size of electron cloud ie. with longer
molecules
3. Less branching to allow for more surface area
interaction.
Dipole-dipole forces
• Occurs between polar
molecules, also called
permanent dipole
• The partially charged ends
of the molecules attract
and repel each other
• Don’t forget London forces
are always acting as well
Dipole-dipole forces
• Polar molecules (polar covalent) have slightly charged ends
• Opposites attract.
• Large electronegative difference = stronger attraction.
Hydrogen Bonds
• Hydrogen bonds are a special form of
dipole attraction
• A hydrogen bond is the attractive force
that forms between an unshared
electron pair and a hydrogen atom
covalently bonded with a strongly
electronegative element (NOF).
• Hydrogen is the only reactive element
without an underlying layer of electrons
– this makes a hydrogen bond about 5%
as strong as an average covalent bond
Hydrogen Bonds
Covalent bond
Hydrogen bond
• Hydrogen Bonding (F, O or N bonded to H) with a free
lone electron pair
• Due to small size and high electronegativity of non metals
• Creates a large charge difference
• Basically a super strong dipole-dipole bond
Boiling point trends
• Phase change when IMF are overcome
• Be sure to explain using the words IMF and
how they affect the bonds BETWEEN particles.
• London Forces are ALWAYS present!!!
General physical property trends
• London forces: Lowest MP, Non polar
• Butane (C4H10)
• Dipole-dipole: Slightly miscible
• Hydrogen Bonding: Miscible with polar substances
• H2O
• Ionic Bonding: Only conducts electricity when liquid
or aqueous. (Decomposition when it does)
• NaCl
• Metallic Bonding: Conducts electricity, not water
soluble, MP regulated by, valance, size and
packing.
• Fe
• Giant Covalent: Highest MP, Insoluble in both nonpolar and polar solvents. Does not conduct
electricity except for graphite.
• Diamond and Graphite (Allotropes)
Increasing Melting Point
• Propanone C3H6O
Van der Waal’s Forces
1.
2. Dipole – Induced dipole
3.
(Induced dipole – Induced dipole)
How IMF affect the boiling points of
substances
–When a liquid turns into a gas the attractive forces between
the particles are completely broken so boiling point is a good
indication of the strength of intermolecular forces.
–Covalent macromolecular structures, such as diamond,
have extremely high melting and boiling points.
–Metals and ionic compounds also tend to have relatively
high boiling points due to ionic attractions.
–Hydrogen bonds are in the order of 1/10th the strength of a
covalent bond whereas London forces are in the order of
less than 1/100th of a covalent bond.
–The weaker the attractive forces the more volatile the
substance.
IMF – comparing compounds
The presence of hydrogen bonding can be
illustrated by comparing the boiling points of:
• HF and HCl
• H2O and H2S
• NH3 and PH3
• CH3OCH3 and CH3CH2OH
• CH3CH2CH3, CH3CHO, and
CH3CH2OH
4.5 Metallic bonding
OBJECTIVES
• A metallic bond is the electrostatic attraction between a lattice
of positive ions and delocalized electrons.
• The strength of a metallic bond depends on the charge of the
ions and the radius of the metal ion.
• Alloys usually contain more than one metal and have enhanced
properties.
• Explanation of electrical conductivity and malleability in
metals.
• Explanation of trends in melting points of metals.
• Explanation of the properties of alloys in terms of nondirectional bonding.
Metallic bonding
Metallic bonds are the electrostatic attraction between a
lattice of positive ions and delocalized electrons.
Metallic bonding
• Strength of metallic bond increases with charge
of ion
• Strength of metallic bond increases with smaller
size of the ion
Metallic bonding – electrical
conductivity and malleability
– For conductivity to occur the substance must possess
electrons or ions that are free to move.
– Metals (and graphite) contain delocalized electrons and are
excellent conductors.
– Metals are malleable which means they can be bent and
reshaped under pressure.
– This due to the close-packed layers of positive ions can
slide over each other without breaking more bonds than are
made.
See difference between covalent brittleness and metals malleability
Alloys
Alloys are mixtures of metals and one or more
other element (usually also a metal).
They produce increased strength, durability,
hardness, resistance to corrosion, and magnetic
properties.
This is due to the different elements being of
different sizes preventing the easy sliding
between atoms.
Alloy examples
Brass is a mixture of copper and zinc.
Steel is a mixture of iron and carbon (and
others).
Type of Covalent
bonding
Cause
Ionic
Metallic
Two strong One
Two weak
electronegat strong
electronegative
ive atoms electroneg atoms
ative and
one weak
electroneg
ative atom
Variation of
strength
within
group
-
Example Diamond
NaCl
Hydrogen
bonding
Dipoledipole
forces
An
Two polar
electronegative molecules
atom from one
molecule and a
hydrogen atom
from another
molecule
London
forces
Two nonpolar
molecules
DECREASING STRENGTH 
The more positive The more polar The more More
the nucleus the
the better
polar the electrons
better
better
the better
The more
electrons the
better
Cu
H2O
SO2
O2
OBJECTIVES
• Covalent bonds result from the overlap of atomic orbitals. A sigma bond (σ) is
formed by the direct head-on/end-to-end overlap of atomic orbitals, resulting
in electron density concentrated between the nuclei of the bonding atoms. A pi
bond (π) is formed by the sideways overlap of atomic orbitals, resulting in
electron density above and below the plane of the nuclei of the bonding atoms.
• Formal charge (FC) can be used to decide which Lewis (electron dot) structure
is preferred from several. The FC is the charge an atom would have if all atoms
in the molecule had the same electronegativity. FC = (Number of valence
electrons)-½(Number of bonding electrons)-(Number of non-bonding
electrons). The Lewis (electron dot) structure with the atoms having FC values
closest to zero is preferred.
• Exceptions to the octet rule include some species having incomplete octets
and expanded octets.
• Delocalization involves electrons that are shared by/between all atoms in a
molecule or ion as opposed to being localized between a pair of atoms.
Higher level
14.1 Further aspects of covalent bonding
and structure (PTO)
OBJECTIVES
• Resonance involves using two or more Lewis (electron dot) structures to
represent a particular molecule or ion. A resonance structure is one of two or
more alternative Lewis (electron dot) structures for a molecule or ion that
cannot be described fully with one Lewis (electron dot) structure alone.
• Prediction whether sigma (σ) or pi (π) bonds are formed from the linear
combination of atomic orbitals.
• Deduction of the Lewis (electron dot) structures of molecules and ions
showing all valence electrons for up to six electron pairs on each atom.
• Application of FC to ascertain which Lewis (electron dot) structure is preferred
from different Lewis (electron dot) structures.
• Deduction using VSEPR theory of the electron domain geometry and
molecular geometry with five and six electron domains and associated bond
angles.
• Explanation of the wavelength of light required to dissociate oxygen and
ozone.
• Description of the mechanism of the catalysis of ozone depletion when
catalysed by CFCs and NOx.
Higher level
14.1 Further aspects of covalent bonding
and structure (cont...)
Formal charge (FC) is:
FC = valence – ½ bonding – non-bonding
electrons
electrons
electrons
To determine the most favoured Lewis structure we
chose:
1. The structure with a FC closest to zero
2. The structure with the negative FC on the most
electronegative atom
Higher level
Formal charge (FC)
FC = valence electrons – ½ bonding electrons non-bonding electrons
FC(N) = 5 - ½ x 6 - 2 = 0
FC(C) = 4 - ½ x 8 - 0 = 0
FC(S) = 6 - ½ x 2 - 6 = -1
Higher level
Problem 1: Determine the FC on the following
compound:
Higher level
The electronegativity of N is 3.0 and S is 2.5, so
the negative FC is best suited to N.
Therefore the bottom structure is correct.
Higher level
Problem 2: Determine which is the preferred
structure.
Higher level
FC(B) = 3 – 1/2x6 – 0 = 0
FC(F)
=7–½x2–6=0
FC(B) = 3 – ½ x8 – 0 = -1
FC(F single bond)
= 7 – ½ 2 x -6 = 0
FC(F double bond)
= 7 – ½ 4 x -4 = +1
The first structure has the most atoms on zero
FC, so is the correct structure.
Elements in the third period (Al to Ar) and below
can promote an electron to have more than 4
electron domains. This is called an expanded
octet.
eg. P = [Ne] 3s13p33d1 making PCl5
5 electron domains: non-bonding electrons in the
equitorial positions of a trigonal bipyramidal
shape
6 electron domains: non-bonding electrons in the
axial positions of a octahedron or square
bipyramidal.
Higher level
VSEPR with expanded shells
• There are two distinct
positions in this
geometry:
– Axial
– Equatorial
Non-bonding electron pairs
fill the equatorial bonds
first because they are only
repelled by 2 bonds (axial).
Axial bonds are repelled
by 3 bonds (equatorial).
Higher level
Trigonal Bipyramidal Electron Domain
Don’t forget non-bonding electron pairs distort
the angles!
Higher level
Trigonal Bipyramidal Electron Domain
Higher level
Trigonal Bipyramidal Electron Domain
There are four
distinct
molecular
geometries in
this domain:
Trigonal
bipyramidal,
Seesaw, Tshaped, Linear
Higher level
Octahedral Electron Domain
All positions are
equivalent in the
octahedral domain.
There are three
molecular geometries:
Octahedral, Square
pyramidal, Square
planar
• s (sigma) bonds result from the axial overlap of
orbitals
•
(pi) bonds result from the sideways overlap of
parallel p orbitals
•
double bonds form by one s (sigma) and one (pi)
bond
•
triple bonds form by one s (sigma) and two (pi)
bonds
Higher level
Overlapping orbitals (s and )
• Sigma bonds (σ) are characterized by
– Head-to-head overlap.
– Cylindrical symmetry of electron density
about the internuclear axis.
Higher level
σ bonds
• Pi bonds (π) are
characterized by
– Side-to-side
overlap.
– Electron density
above and below
the internuclear
axis.
Higher level
π bonds
In a multiple bond one of the bonds is a s bond and
the rest are  bonds.
Higher level
Single bonds are always s bonds, because s overlap
is greater, resulting in a stronger bond and more
energy lowering.
Higher level
• Note the weakness of the extra bond –
the electron overlap is weaker, yet it is
still an extra bond so the energy required
to break them is still a bit stronger.
• Note the smaller bond lengths have
greater bond dissociation enthalpies.
Higher level
CASE STUDY CONTINUED
OZONE
Bond order is the measurement of the number of
electrons involved in bonds between two atoms in a
molecule. If the bond order is 0 then there is no
bonding.
𝑡𝑜𝑡𝑎𝑙 𝑛𝑢𝑚𝑏𝑒𝑟 𝑜𝑓 𝑏𝑜𝑛𝑑𝑖𝑛𝑔 𝑝𝑎𝑖𝑟𝑠
Bond order =
𝑡𝑜𝑡𝑎𝑙 𝑛𝑢𝑚𝑏𝑒𝑟 𝑜𝑓 𝑝𝑜𝑠𝑖𝑡𝑖𝑜𝑛𝑠
Problem 1: Determine the bonding order of ozone.
Bond order = 3/2 = 1.5
Higher level
Resonance and bond order
𝑡𝑜𝑡𝑎𝑙 𝑛𝑢𝑚𝑏𝑒𝑟 𝑜𝑓 𝑏𝑜𝑛𝑑𝑖𝑛𝑔 𝑝𝑎𝑖𝑟𝑠
Bond order =
𝑡𝑜𝑡𝑎𝑙 𝑛𝑢𝑚𝑏𝑒𝑟 𝑜𝑓 𝑝𝑜𝑠𝑖𝑡𝑖𝑜𝑛𝑠
4
=
3
Higher level
Problem 2: Determine bond order
Ozone at ground level reacts with chemicals to form smog, harms
respiratory systems and degrades materials (eg. rubber).
Ozone at atmospheric levels is
essential for life.
1. It absorbs harmful UV
radiation (which causes
cancer and inhibition of
photosynthesis).
2. Exothermic reactions cause a
temperature inversion in the
stratosphere – this warm
layer prevent convection
keeping the layers of Earth’s
atmosphere stable.
Higher level
Importance of ozone
Review: O2 requires photons with more energy
as O3 only has the energy of a 1 ½ bond.
O2(g) + UV (λ<242nm)  O•(g) + O•(g)
O3(g) + UV (λ<330nm)  O2(g) + O•(g)
Bond angle <120⁰
~ 117⁰
Higher level
Ozone formation review
Equations:
Ephoton = hf
E = energy in one photon (J)
h = planks constant (6.63x10-34Js)
f = frequency (Hz)
v = fλ
v = velocity = c = speed of light = (3.00 x 108 m/s)
λ = wavelength (nm)
Higher level
UV and O2/O3 bond calculations
Ephoton = 363 000J/mol / 6.02x1023photons/mol
= 6.03x10-19J per photon
Ephoton = hf so
f = Ephoton / h
= 6.03x10-19J / 6.63x10-34Js
= 9.09 x 1014s
v = fλ so
λ = v/f = 3.00 x 108 m/s / 9.09 x 1014s
= 3.30 x 10-7m
= 330nm
Higher level
Problem 1: The bond energy in ozone is 363kJ/mol.
Calculated the wavelength of UV radiation needed to
break this bond.
UV light breaks the bonds in oxygen to create free
radicals (atoms/molecules with unpaired valence
electrons and so highly reactive):
• is the symbol to
O2(g) + UV (λ<242nm)  O•(g) + O•(g)
denote a radical.
O• actually has 6
electrons in the
valence shell (not
one).
This then undergoes an exothermic reaction:
O•(g) + O2(g)  O3(g)
Higher level
Ozone formation
The reverse also occurs creating an O3 cycle.
Depletion is also exothermic.
O3(g) + UV (λ<330nm)  O2(g) + O•(g)
O3(g) + O•(g)  2O2(g)
Higher level
Ozone depletion
1. CFCs – chloroflourocarbons, catalyze the
breakdown of ozone
2. NOx – also catalyze the breakdown of ozone
These compounds are free radicals (not complete
octets) and hence are highly reactive.
Higher level
Ozone destroying chemicals
The balance of the ozone cycle is disrupted by the
depletion of ozone from:
Nitrogen monoxide free radical
(unpaired valence electrons)
1. nitrogen oxides:
NO•(g) + O3(g)  NO2•(g) + O2(g)
NO2•(g) + O•(g)  NO•(g) + O2(g)
Nitrogen dioxide free radical
(unpaired valence electrons)
Nitrogen oxides slowly diffuse their way from the
troposphere. Aircraft fly in the lower stratosphere causing
direct injection of NOx. NO can last from 22 to 111 years
before breaking down.
Higher level
Catalytic ozone destruction
2. chlorofluorocarbons (CFCs):
CCl2F2(g)  CClF2•(g) + Cl•(g) by UV radiation
Cl•(g) + O3(g)  O2(g) + ClO•(g)
ClO•(g) + O•(g)  O2(g) + Cl•(g)
ClO•(g) + O3 (g)  2O2(g) + Cl•(g)
As with NO, the Cl free radical catalyzes the decomposition
reaction thousands of times:
Summary: O3(g) + O•(g)  2O2(g)
Higher level
Catalytic ozone destruction
OBJECTIVES
• A hybrid orbital results from the mixing of different types of
atomic orbitals on the same atom.
• Explanation of the formation of sp3, sp2 and sp hybrid orbitals
in methane, ethene and ethyne.
• Identification and explanation of the relationships between
Lewis (electron dot) structures, electron domains, molecular
geometries and types of hybridization.
Higher level
14.2 Hybridization
OBJECTIVES
• A hybrid orbital results from the mixing of different types of
atomic orbitals on the same atom.
• Explanation of the formation of sp3, sp2 and sp hybrid orbitals
in methane, ethene and ethyne.
• Identification and explanation of the relationships between
Lewis (electron dot) structures, electron domains, molecular
geometries and types of hybridization.
Higher level
14.2 Hybridization
Higher level
It’s hard to imagine tetrahedral, trigonal bipyramidal,
and other geometries arising from the atomic orbitals
we recognize.
Hybridization is the mixing of different types of orbitals to
produce new types of orbitals. It is a mathematical
procedure.
The most common hybrid orbitals are combinations of s and
p orbitals that then form sigma bonds.
(The number of orbitals mixed equals the number of hybrid
orbitals produced. Not all orbitals in a level are hybridized.
Count the number of sigma bonds to determine the number
of hybridized orbitals.)
Higher level
Orbital hybridization
Higher level
Orbital hybridization
In triple bonds, as
in acetylene, two sp
orbitals form a s
bond between the
carbons, and two
pairs of p orbitals
overlap to form the
two  bonds.
Higher level
Two sp orbitals
• In a molecule like
formaldehyde (shown
at left) an sp2 orbital
on carbon overlaps to
form a s bond with
the corresponding
orbital on the oxygen.
• The unhybridized p
orbitals overlap in 
fashion.
Higher level
Three sp2 orbitals
• In a molecule like
methane (shown at
left) the sp3 orbitals
on carbon overlap to
form a s bond with
the corresponding
orbitals on hydrogen.
Higher level
Four sp3 orbitals
• Pi bonds form from overlapping p orbitals
• C and O form hybrid orbitals to make the sigma bond
Higher level
Bonding in CO2
• Consider beryllium:
– In its ground electronic
state, it would not be able
to form bonds because it
has no singly-occupied
orbitals.
Higher level
BeF2 Hybridization – orbital diagrams
But if it absorbs the small
amount of energy
needed to promote an
electron from the 2s to
the 2p orbital, it can
form two bonds.
σ bond
π bond
Higher level
BeF2 Hybridization – orbital diagrams
• Mixing the s and p orbitals yields two degenerate
orbitals that are hybrids of the two orbitals.
– These sp hybrid orbitals have two lobes like a p orbital.
– One of the lobes is larger and more rounded as is the s
orbital.
Higher level
BeF2 Hybridization
• These two degenerate orbitals would align themselves
180 from each other.
• This is consistent with the observed geometry of
beryllium compounds: linear.
Higher level
BeF2 Hybridization
σ
bond
σ
bond
• With hybrid orbitals the orbital diagram for
beryllium would look like this.
• The sp orbitals are higher in energy than the
1s orbital but lower than the 2p.
Higher level
BeF2 Hybridization – orbital diagrams
Using a similar model for boron leads to…
σ
σ
σ
bond bond bond
Higher level
BF3 Hybridization – orbital diagrams
…three
degenerate
sp2 orbitals.
Higher level
BF3 Hybridization
With carbon we get…
Higher level
CH4 Hybridization – orbital diagrams
σ
σ
σ
σ
bond bond bond bond
• …four degenerate
sp3 orbitals.
Higher level
CH4 Hybridization
H
C
N
C
N
Higher level
CHN Hybridization – animation
C
N
Higher level
C
N
Higher level
C
N
Higher level
H
C
N
Higher level
Hybrid energy state
Lone
pair
3σ
bonds
•Number of electron
domains will tell you
what hybidization
there is.
•4 electron domains,
therefore 4 sp3
hybyridization
Higher level
Hybridization of nitrogen in ammonia
Higher level
Hybridization of oxygen in CO2
π bond
•Number of electron
domains will tell you
what hybidization
there is.
• Pi bonds form from
normal p orbitals
Hybrid energy state
Lone Lone
σ
pair pair bond
Higher level
Problem 1: Analyse the SO42- molecule.