Csonn t3 chemical bonding

Chemical bonding
and structure

Chemical bond
• Chemical bonds are electrostatic forces (attraction between
positive charge and negative charge) that bind particles together to
form matter.
• Different types of chemical bonds:
• a) Metallic bond
• b) Ionic bond
• c) Covalent bond
• d) Intermolecular forces (NEED TO SPECIFY WHICH ONE - LATER!)

Ionic bonding
• Ionic compound has lattice of cations + anions.
• Ionic bond is the electrostatic force of attraction
between oppositely charged ions formed by
electron transfer.
• Octet rule - general but not 100% always!

Ionic bonding - ‘dot-and-
cross’
Draw dot-and-cross diagrams for the following ionic
compounds.
a) MgO
b) CaCl2
c) Na2O

Ionic bonding
• NOT electron transfer
• INVOLVES electron transfer
• Also known as electrovalent bonding (CIE syllabus)
• Metals lose electrons and form cations => think about
ionisation energies (topic 2)
• Across a period, Zeff, hence ﬁrst IE?
• Down a group, Zeff, hence ﬁrst IE?

Ionic bonding
• Non-metals gain electrons
• (A2 syllabus: electron affinity - 1st electron affinity -
enthalpy change when one mole of electrons is added
to one mole of gaseous atoms to form one mole of
gaseous singly charged anions)
• O (g) + e- —> O- (g) 1st EA = -140 kJmol-1
• O+ (g) + e- —> O2- (g) 2nd EA = + 798 kJmol-1
• WHY difference? - think of electronic configuration?

Ionic bonding - Lattice
(formation) enthalpy
• the enthalpy change
• when 1 mole of ionic solid crystal
• is formed from its scattered gaseous ions.
• under standard conditions (298 K and 1 atm pressure)
• Lattice formation enthalpies are always negative (ΔH <
0)
• Eqn example?

Ionic bonding - Lattice
(formation) enthalpy
• Greater charge densities of ions, the more they
attract each other, the larger the lattice enthalpy
• The more exothermic the lattice enthalpy, for ionic
compounds, the higher the m.p.

Ionic bonding
• Most important factor is LATTICE ENERGY
• The lattice (formation) enthalpy is the enthalpy change
when 1 mole of solid crystal is formed from its
scattered gaseous ions.
• E.g. Na+ (g) + Cl- (g) —> NaCl (s)
• E.g. Write an equation to show lattice formation
enthalpy change for formation of MgO and Al2O3.
• Lattice formation enthalpy is always EXOTHERMIC

Ionic bonding
• Two factors affect lattice energy:
• a) charges on the ion
• b) size of the ion
• X-ray diffraction studies - absolute proof
• For simple ions,
• a) charge determines the balance between numbers of cations
and anions
• b) radii determine the way ions pack in lattice

Ionic bonding - NaCl
• Called the “rock salt structure”
• NaCl and MgO has this ionic lattice structure
• 6:6 coordination number - meaning?
• Can sketch?

Ionic bonding - Properties
• High melting and boiling points
• Good electrical conductivity only when molten
• Generally soluble in polar solvents

Ionic bonding - Q1
Which of the following statements are correct for the
sequence of compounds below considered from left to right?
NaF MgO AlN SiC
(1) The electronegativity difference between the elements in
each compound increases.
(2) The formulae-units of these compounds are isoelectronic
(have the same number of electrons).
(3) The bonding becomes increasingly covalent.

Covalent bonding
• The covalent bond is the electrostatic attraction
between the localised shared electrons and the two
positively charged nuclei.
• Covalent bonds are directed in space.
• NOT sharing electrons
• INVOLVES sharing of electrons
• Simple molecular or giant molecular? Single, double or
triple bonds? Sigma or pi bond?

Covalent bonding
• Sigma bonds - formed when orbitals overlap
head-on

Covalent bonding
• Pi bonds - formed when orbitals overlap side-on

Covalent bonding - sigma
and pi bonds
• Extent of bonding depends on orbital overlap
• similar in energy (i.e. similar sized orbitals)
• similar in symmetry
• Can you form pi bond without sigma bond?
• Absolutely not; if given the options of forming a sigma or a pi
bond, the two atoms “prefer” to form a sigma bond.
• A sigma bond is formed from more effective overlap of the
atomic orbitals compared to pi bond.
• A sigma bond is thus more stable and stronger than a pi bond.

and pi bonds
• Can there be two sigma bonds formed between two
atoms?
• Absolutely not; formation of two sigma bonds will result
in too much accumulation of electron density within
the inter-nuclei region.
• The inter-electronic repulsion will be too great.
• The different ways of forming sigma and pi bonds
spread out the inter-nuclei electron density to
minimise inter-electronic repulsion but maximising
bond strength.

and pi bonds
Sigma (σ) bond Pi (π) bond
Formed due to the axial overlap of two
orbitals (‘s-s’, ‘s-p’or’p-p’).
Formed by the lateral (sideways) overlap of
two ‘p’ orbitals.
Only one sigma bond exists between two
atoms.
There can be more than one pi bonds between
the two atoms.
The electron density is maximum and
cylindrically symmetrical about the bond axis.
The electron density is high along the
direction at right angles to the bond axis.
Free rotation about the sigma bond is possible. Free rotation about the pi bond is not possible.
This bond can be independently formed, i.e.,
without the formation of a pi bond.
The pi bond is formed after the sigma bond
has been formed,
Sigma bond is relatively strong. Pi bond is a relatively weaker bond.

and pi bonds - Q1
Which diagram describes the formation of a pi bond
from the overlap of its orbitals?5
5 Which diagram describes the formation of a π bond from the overlap of its orbitals?
D
A
B
C
6 For an ideal gas, the plot of pV against p is a straight line. For a real gas, such a plot shows a

and pi bonds - Q2
Which statements about covalent bonds are correct?
(1) A triple bond consists of one pi bond and two
sigma bonds.
(2) The electron density in a sigma bond is highest
along the axis between the two bonded atoms.
(3) A pi bond restricts rotation about the sigma bond
axis.

and pi bonds - Q2

and pi bonds - Q3
What is always involved in a carbon-carbon pi bond?
(1) a shared pair of electrons
(2) a sideways overlap of p orbitals
(3) delocalised electrons

and pi bonds - Q3

and pi bonds - Q4
Carvone is found in spearmint.
How many sigma and pi bonds are present in this
molecule?
r
proton number
r
proton number
r
proton number
r
proton number
20 Carvone is found in spearmint.
C
C
C
C
CH2
H3C
H2C
CH2
CH3
C
H
HO
carvone
How many σ and π bonds are present in this molecule?
σ π
A 13 3

and pi bonds - Q4

and pi bonds - Q5
The diagram shows a molecule that has sigma bonds
and pi bonds.
How many sigma and pi bonds are present in this
molecule?
at is compound X?
CH3CO2C2H5
CH3CH2COCH3
(CH3)3COH
CH3CH2CHOHCH3
e diagram shows a molecule that has σ bonds and π bonds.
C
O
OCH2CH2 CH2 CH2CHCH
w many σ bonds are present in this molecule?
15 B 17 C 18 D 21

and pi bonds - Q5

and pi bonds - Q6
4
8 A covalent molecule contains
● 14 electrons,
● one lone pair of electrons,
● two π bonds.
What is the molecule?
A C2H4 B HCN C H2O2 D N2
9 Which value is essential to calculate the lattice energy of the compound NaH?

and pi bonds - Q6

and pi bonds - Q7
Which statement about bond formation is not correct?
A. A triple bond consists of one σ bond and two π bonds.
B. A π bond restricts rotation about the σ bond axis.
C. Bonds formed from atomic s orbitals are always σ
bonds.
D. End-to-end orbital overlap results in a bond with
electron density above and below the bond axis.

and pi bonds - Q7

Covalent bonding - Dot-and-cross
diagrams & Lewis diagrams
m.p./°C b.p./°C Type of bonding
LiCl 605 1382 ?
NaCl 801 1413 ?
BeCl2 405 520 ?
MgCl2 714 1412 ?
AlCl3 sublimes at 180°C ?

m.p./°C b.p./°C Type of bonding
Li2O 1438 - ?
BeO 2507 - ?
Na2O 1132 - ?
MgO 2852 - ?
Al2O3 2072 2977 ?
SiO2 1600 2230 ?
P4O6 24 173 ?
P4O10 340 360 ?
SO2 -72 -10 ?
SO3 17 45 ?

H2 Cl2 O2 N2
BeO CO2 CS2 BH3
HCl CH4 C2H4 C2H2
CCl4 CFCl3 CHCl3 CHI3
SiH4 SiCl4 HClO HClO3
NH3 NCl3 PF3 PF5
H2O H2S H2Se OF2

SO2 SO3 SCl2 SF4
SF6 P4O6 P4O10 ICl
OCS HCN NH4+ BH4-
CO AlCl4- AlH4- N(CH3)4+
CH3+ CH3- SbCl5 SbF6-
H2SO4 CO32- PO43- NO3-

• Atoms share electrons in order to complete their “octet” of
electrons
• Some don't achieve an octet as they don’t have enough
electrons - Al in AlCl3, B in BCl3
• Others share only some - if they share all valence
electrons, their octet is exceeded - ie NH3, NCl3 and H2O,
OF2 (cf PCl5, SF6)
• Atoms of elements in the 3rd period onwards can exceed
their octet because they are not restricted to 8 electrons in
their outer shell (cf n=3 shell can use 3d for bonding).

Dative (coordinate) covalent
bonding
• A dative covalent (or coordinate) bond is a
covalent bond in which both the electrons shared
come from the same atom.
• Donor species will have lone pairs in their outer
shells
• Acceptor species will be short of their “octet” or
maximum

bonding
Al2Cl6 molecule - dot and cross - Lewis diagram - arrow
important - properties - common exam questions

bonding
• Ammonia + proton —> Ammonium ion
• Ammonia + borane (BH3) —> Ammonium borane
• Ammonia + BeCl2 —> ?
• Draw Lewis diagram & “dot-and-cross” - count
electrons!

bonding

bonding - Q1
solutions have the same effect on litmus.
What is element X?
A sodium
B magnesium
C aluminium
D phosphorus
15 Aluminium chloride sublimes at 178o
C.
Which structure best represents the species in the vapour at this temperature?
Cl
Cl
Cl
Cl
Al
Cl
Cl
Al
Cl
Cl
Cl
Cl
Al
Cl
Cl
Al
Al + 3Cl Al3+
(Cl –
)3
A B C D
16 Use of the Data Booklet is relevant to this question.
What mass of solid residue can be obtained from the thermal decomposition of 4.10 g of
anhydrous calcium nitrate?

bonding - Q2
AlCl3 reacts with LiAlH4 and (CH3)3N to give (CH3)3NAlH3.
Which statement about (CH3)3NAlH is correct?
A. It contains hydrogen bonding.
B. It is dimeric.
C. The Al atom has an incomplete octet of electrons.
D. The bonds around the Al atom are tetrahedrally
arranged.

bonding - Q2

bonding - Q3
Trimethylamine, Me3N, reacts with boron triﬂuoride, BF3, to form
a compound of formula Me3N.BF3. How may this reaction be
written in terms of the shapes of the reactants and products?
When barium metal burns in oxygen, the ionic compound barium peroxide, BaO2, is formed.
Which dot-and-cross diagram represents the electronic structure of the peroxide anion in BaO2?
A C DB
electron from
first oxygen atom
electron from
second oxygen atom
electron from
barium atom
key
In this question, the methyl group, CH3, is represented by Me.
Trimethylamine, Me3N, reacts with boron trifluoride, BF3, to form a compound of formula
Me3N.BF3.
How may this reaction be written in terms of the shapes of the reactants and products?
Me Me F B
Me F
N F
Me Me
Me
Me
Me
Me
N
+
N
F B
F
F
A
Me Me
BMe
F F
N F
B
F F
F
+B
Me Me F B
Me F
Me Me
Me
N
+C
© UCLES 2008 9701/01/M/J/08
5 In this question, the methyl group, CH3, is represented by Me.
Trimethylamine, Me3N, reacts with boron trifluoride, BF3, to form
Me3N.BF3.
How may this reaction be written in terms of the shapes of the reactants
Me Me F B
Me F
N F
Me Me
Me
Me
Me
Me
N
+
N
F B
F
F
A
Me Me
BMe
F F
N F
B
F F
F
B
F F
F
B
F F
F
+B
Me Me F B
Me F
N F
Me Me
Me
N
+
B
F F
F
C
Me Me
Me
N
Me Me
Me
N
+D

bonding - Q3

bonding - Q4
Which diagram correctly shows the bonding in the
ammonium ion, NH4
+?
What is the same in an atom of 4
He and an atom of 3
H?
A the number of electrons
B the number of neutrons
C the number of protons
D the relative atomic mass
2 Which diagram correctly shows the bonding in the ammonium ion, NH4
+
?
N
H
H
H H
+
A
N
H
H
H H
+
B
N
H
H
H H
+
C
N
H
H
H H
+
D
key
N electron
H electron

bonding - Q4

bonding - Q5
Why does aluminium chloride, Al2Cl6, sublime at the
relatively low temperature of 180°C?
(1) The intermolecular forces between the Al2Cl6
molecules are weak.
(2) The co-ordinate bonds between aluminium and
chlorine are weak.
(3) The covalent bonds between aluminium and
chlorine are weak.

bonding - Q5

bonding - Q6
In the gas phase, aluminium chloride exists as the dimer,
Al2Cl6. By using this information, which of the following
are structural features of the Al2Cl6 molecule?
(1) Each aluminium atom is surrounded by four chlorine
atoms.
(2) There are twelve non-bonded electron pairs in the
molecule.
(3) Each aluminium atom contributes electrons to four
covalent bonds.

bonding - Q6

bonding - Q7
Aluminium chloride catalyses certain reactions by forming
carbocations (carbonium ions) with chloroalkanes as shown.
Which property makes this reaction possible?
A. AlCl3 is a covalent molecule.
B. AlCl3 exists as the dimer Al2Cl6 in the vapour.
C. The aluminium atom in AlCl3 has an incomplete octet of
electrons.
D. The chlorine atom in RCl has a vacant p orbital.

bonding - Q7

Covalent bonding - simple
covalent ions
• Ammonium salt has two types of bonding
• Ionic bonding between ammonium ion and anion
• Covalent bonding (including dative covalent) within
ammonium ion
• Other examples include: nitrates, sulfates, phosphates,
carbonates, etc.
• How we know this? - easy - X-ray crystallographic data!
- bond length - compare single vs double bond length?

Covalent bonding -
delocalisation of electrons

Molecular geometry -
VSEPR
• Valence shell electron pair repulsion theory
• What is an electron pair?
• What is a bonding electron pair?
• What is a non-bonding electron pair/lone pair?
• Electron clouds repel each other. Why?
• Extent of repulsion: LP/LP > BP/LP > BP/BP

VSEPR
H
H
O
Bent line 2 2 104.5 OCl2, H2S, OF2 ,
SCl2
.. ..
O
H H
P
F
F
F
F
F
S
FF
FF
F
F
SCl2
Trigonal
Bipyramidal
5 0 120 and 90 PCl5
Octahedral 6 0 90 SF6
Occasionally more complex shapes are seen that are variations of octahedral and trigonal
bipyramidal where some of the bonds are replaced with lone pairs.
e.g XeF4 e.g. BrF5 e.g I3 e .g.ClF3 e.g. SF4
Remember lone pairs repel more than bonding pairs and so reduce bond angles
X
:
X:X
:
:X
::

VSEPR
In this order,
a) state number of bonding pairs and lone pairs of electrons
b) state that electron pairs repel and try to get as far apart as
possible (or to a position of minimum repulsion)
c) IF there are no lone pairs, state that the electron pairs repel
equally
d) IF there are lone pairs of electrons, then state that lone
pairs repel more than bonding pairs.
e) state actual shape and bond angle.

VSEPR
• Species with an odd number of valence electrons,
e.g. NO
• Electron deﬁcient compound, e.g. BH3, BF3
• Species with expanded valence shells, e.g. SF6,
PCl5

VSEPR - Q1
XeF4 HCN H2Se ICl4
+
SO4
2- SbF5
-
N2H4 H3O+
C2Cl4 PH3 BH4
- AlH4
-
Al(CH3)2Cl

VSEPR - Q1

VSEPR - Q2
In which sequences are the molecules quoted in
order of increasing bond angle within the molecule?
1 H2O NH3 CH4
2 H2O SF6 BF3
3 CH4 CO2 SF6

VSEPR - Q2

VSEPR - Q3
Chloroethene, CH2=CHCl, is the monomer of pvc.
What are the C-C-C bond angles along the polymeric
chain in pvc?
A. They are all 109 °.
B. Half are 109 ° and half are 120 °.
C. They are all 120 °.
D. They are all 180 °.

VSEPR - Q3

VSEPR - Q4
Which molecule contains only six bonding electrons?
A. C2H4
B. C2F6
C. H2O
D. NF3

VSEPR - Q4

VSEPR - Q5
Which of the following molecules and ions have a
regular trigonal planar shape?
(1) AlCl3
(2) CH3
+
(3) PH3
(4) BCl3
(5) NH3

VSEPR - Q5

VSEPR - Q6
Methyl isocyanate, CH3NCO, is a toxic liquid which is used in
the manufacture of some pesticides. In the methyl isocyanate
molecule, the sequence of atoms is H3C-N=C=O.
What is the approximate angle between the bonds formed by
the N atom?
3
4 Methyl isocyanate, CH3NCO, is a toxic liquid which is used in the manufacture of some
pesticides.
In the methyl isocyanate molecule, the sequence of atoms is H3C—N C O.
What is the approximate angle between the bonds formed by the N atom?
A
104
B
109
C
120
D
180
N CH3
C O
N C
H3
C
ON C
H3
C
ON C
H3
C
O
5 At room temperature and pressure chlorine does not behave as an ideal gas.
At which temperature and pressure would the behaviour of chlorine become more ideal?

VSEPR - Q6

VSEPR - Q7
Which statements about bond angles are correct?
(1) The bond angle in SO2 is smaller than the bond
angle in CO2.
(2) The bond angle in H2O is smaller than the bond
angle in CH4.
(3) The bond angle in NH3 is smaller than the bond
angle in BF3.

VSEPR - Q7

VSEPR - Q8
The diagram shows an example of an organic nitrate
molecule.
What is the correct order of the bond angles shown in
ascending order?
8
B H3O+
C OD–
D OH–
s in photochemical smog can cause breathing difficulties.
hows an example of an organic nitrate molecule.
H C
H
H
O
C O O NO2
1
2
3
rect order of the bond angles shown in ascending order (smallest
B 2 → 1 → 3 C 3 → 1 → 2 D 3 → 2 → 1

VSEPR - Q8

VSEPR - Q9
The CN- is widely used in the synthesis of organic
compounds. What is the pattern of electron pairs in
this ion?
No of bonding pairs of electrons = ?
No of lone pairs on carbon atom = ?
No of lone pairs on nitrogen atom = ?

VSEPR - Q9

VSEPR - Q10
The antidote molecule shown can help to prevent liver
damage if someone takes too many paracetamol tablets.
What is the order of decreasing size of the bond angles
x, y and z?
A ionic radius
B ionisation energy
C neutron/proton ratio
D rate of reaction with water
he antidote molecule shown can help to prevent liver damage if someone takes too
aracetamol tablets.
H S C
H H H
H H H
C N
x y
z
represents a
lone pair
What is the order of decreasing size of the bond angles x, y and z?
largest smallest
A x y z
B x z y

VSEPR limits
17 HF Hydrogen fluoride Hydrogen fluoride
17 HCl Hydrogen chloride Hydrogen chloride
17 HBr Hydrogen bromide Hydrogen bromide
17 HI Hydrogen iodide Hydrogen iodide
The electron-deficient hydrides are those that cannot complete an octet of electrons around the central atom. They are
chiefly the Group 13 elements, although the gas-phase only species BeH2 also fits this description. The electron precise
compounds are those that have an octet of electrons, while the electron-rich elements have additional electrons belonging to
the central atom that function as lone pairs. Note that although they are all electron rich, this group of compounds vary
greatly in the availability of these extra electrons.
Note that while the structures of these hydrides follow similar patterns, i.e. all the EH3 in group 15 are pyramidal, the
detailed structures differ significantly. Consider the following table of bond lengths for both groups 15 and 16:
Group 15 hydrogen compounds Bond angle Group 15 hydrogen compounds Bond angle
NH3 106.6° H2O 104.5°
PH3 93.8° H2S 92.1°
AsH3 91.8° H2Se 91°
SbH3 91.3° H2Te 89°
Source: A.F. Wells, Structural Inorganic Chemistry, Oxford University Press (1984)
In the last problem set,
you have explored the
origins of some of these
changes using molecular
orbital methods. The
bending of H–E–H
systems is due to a Jahn-
Teller effect that lowers
the energy of what would
be a degenerate set of E
orbitals in the linear
molecule. The greater
bending in the third and
subsequent periods is due
to a second-order Jahn-

Giant molecular compounds
• Some covalent compounds are not discrete molecular
compounds.
• Giant molecular compounds include,
(a) diamond
(b) SiO2, silicon (IV) oxide, sand - all same
(c) graphite
(d) BN (isoelectronic to carbon)
(e) silicon - same like diamond

- diamond
• mp of diamond = 3350 °C
• mp of silicon = 1410 °C
• For melting in giant molecular compounds, a
great amount of energy is required to overcome
the strong covalent bonds between the atoms.
• Bond strength: BDE(C-C) > BDE(Si-Si)

- graphite
• C-C (between layers) < C-C (same adjacent layer)
• Longer bond lengths => weaker bonding between these atoms
• VDW forces of attraction holds layers of graphite together, not
actual covalent bonds

Chemical bonding - Q1
Which solid exhibits more than one kind of chemical
bonding?
A. brass
B. copper
C. diamond
D. ice

Which of the following solids contain more than one
type of chemical bond?
1. brass (an alloy of copper and zinc)
2. graphite
3. ice

Which solid has a simple molecular lattice?
A. calcium ﬂuoride
B. nickel
C. silicon(IV) oxide
D. sulfur

Which diagrams represent part of a giant molecular
structure?
The responses A to D should be selected on the basis of
A B C D
1, 2 and 3
are
correct
1 and 2
only are
correct
2 and 3
only are
correct
1 only
is
correct
No other combination of statements is used as a correct response.
31 Which diagrams represent part of a giant molecular structure?
1 2 3
= C = C = Na
= C
32 Which reactions are redox reactions?
1 CaBr2 + 2H2SO4 → CaSO4 + Br2 + SO2 + 2H2O

Covalent bonding - Bond
dissociation enthalpy (BDE)
• Energy required

• to break one mole of covalent bond

• to give separated atoms

• with everything being in the gas state

• Eqn example?

Covalent bond - bond
energy
• Bond dissociation energy (BDE) is the energy required
to break the bond between two covalently bonded atoms.
• Polyatomic molecules?
• X-ray diffraction - covalent bond length and covalent
bond radii
• By halving the interatomic distances obtained for diatomic
elements = covalent bond radii
• Suggest when would actual covalent radii be very
different to that predicted by tabulated covalent radii.

Covalent bond - bond
polarity
• Bonding electron pair is shared equally between
two same atoms that form a covalent bond.
• For any two unlike atoms, the bonding electron pair
sharing is always unequal.
• Unequal sharing of electrons can have two
extremes - both leading to ionic bond.
bond radii, but these theoretical values often differ from the experi-
mental values; the greatest deviations occur when elements ofwidely
different electronegativities are joined together.
ELECTRONEGATIVITY
If two like atoms form a covalent bond by sharing an electron pair,
for example
x F * FxX X X X
it is clear that the pair will be shared equally. For any two unlike
atoms, the sharing is always unequal and depending on the nature
of the two atoms (A and B say)we can have two extreme possibilities
or A :B i.e. A+
B"
A + B->A : B ^
equal
sharing X
or A ; g ie A- B +
and an ionic bond is formed. There are many compounds which lie

Bond polarity -
Electronegativity
• Many compounds are intermediate between truly covalent
(equal sharing) and truly ionic.
• Electronegativity is a measure of the tendency of an atom to
attract a bonding pair of electrons.
• Fajan’s rules state that covalent character in ionic compounds
increases if,
• a) small cations - highly polarising
• b) large anions - highly polarisable
• c) high charge on cations and anions

Covalent bond - Electronegativity
• Electronegativity is a measure of the tendency
of an atom to attract a bonding pair of electrons.

Covalent bond -
Electronegativity - Arkel diagram

Covalent bond - Electronegativity
- Factors affecting it
• Think of Zeff
• The attraction that a bonding pair of electron feels
for a particular nucleus depends on:
• a) the number of protons in the nucleus
• b) the distance from the nucleus
• c) the amount of screening by inner electrons

Electronegativity - diagonal
relationship
• Boron is a Gp III non-metal with some properties like silicon.
• Beryllium is a Gp II metal with some properties resembling
aluminium.
• Electronegativity reasoning:
• Be = 1.5, B = 2.0
• B = 2.0, Al = 1.5
• Similar electronegativity, likely to form similar types of
bonds, hence similar chemistry

Ionic or covalent? - Q1
Which pairs of compounds contain one that is giant
ionic and one that is simple molecular?
(1) Al2O3 and Al2Cl6
(2) SiO2 and SiCl4
(3) P4O10 and PCl3
(4) B(NMe3)3 and BN
(5) C(CH3)4 and diamond

Which chlorine compound has bonding that can be
describe as ionic with some covalent character?
A. NaCl
B. MgCl2
C. AlCl3
D. SiCl4

Which element shows the greatest tendency to form
some covalent compounds?
A. Aluminium
B. Magnesium
C. Neon
D. Potassium

When barium metal burns in oxygen, the ionic
compound barium peroxide, BaO2 is formed. Which
dot-and-cross diagram represents the electronic
structure of the peroxide anion in BaO2?
3
4 When barium metal burns in oxygen, the ionic compound barium peroxide, BaO2, is formed.
Which dot-and-cross diagram represents the electronic structure of the peroxide anion in BaO2?
A C DB
electron from
first oxygen atom
electron from
second oxygen atom
electron from
barium atom
key
5 In this question, the methyl group, CH3, is represented by Me.

Dipole moment
• Unequal distribution of charge produced when
elements of different electronegativities polarises a
covalent bond joining them.
• Unless this polarity is balanced by an equal and
opposite polarity, the molecule will be a dipole and have
a dipole moment (i.e. hydrogen halide - HF, HCl, HBr,
HI)
• Examples where dipole moments get cancelled out -
there is no net dipole moment => non-polar molecule.

How to determine if a
molecule is polar?
As long as a molecule has one of the standard shapes shown
below, with identical bonding atoms and no lone pairs on the
central atom, then the molecule has no net dipole moment.
=> Linear
=> Trigonal planar
=> Tetrahedral
=> Trigonal bipyramidal
=> Octahedral

molecule is polar?

molecule is polar?
• If the molecule contains one lone pair on the central atom, then the
molecule has a net dipole moment.
• But, a molecule with more than one lone pair on the central atom
may or may not be polar.
• Eg. NF3, H2O, XeF4

molecule is polar?
• Polar bonds ≠ Polar molecule
• A polar molecule has a net dipole moment -
hence permanent dipole intermolecular forces
between their molecules.

Polar molecule - Q1
Which of the following molecules has no permanent
dipole?
A. CCl2F2
B. CHCl3
C. C2Cl4
D. C2H5Cl

Polar molecule - Q2
Which molecule has the largest overall dipole?
4
5 Which molecule has the largest overall dipole?
C C
C C
C C
A B
O C
C
O C
D
O C O
C
C
H
H
6 The first stage in the industrial production of nitric acid from ammonia can be represented by the
following equation.
4NH3(g) + 5O2(g) 4NO(g) + 6H2O(g)
Using the following standard enthalpy change of formation data, what is the value of the standard
enthalpy change, ∆Ho
, for this reaction?

Polar molecule - Q3
Which molecule has the largest overall dipole?
CO2(g) + 3H2(g) CH3OH(g) + H2O(g) ∆H = –49kJmol
What would increase the equilibrium yield of methanol in this process?
A adding a catalyst
B adding an excess of steam
C increasing the pressure
D increasing the temperature
10 Which molecule has the largest overall dipole?
CH3
CH3
H
H
C C
A
CH3
CH3
O C
B
Cl
Cl
O C
C
H3C
Cl
Cl
CH3
C C
D
11 In which substance does nitrogen exhibit the highest oxidation state?
A NO B N2O C N2O4 D NaNO2
12 Red lead oxide, Pb3O4, is used in metal priming paints. It can be made by heating PbO in air.

Polar molecule - Q4
The greater the difference between the electronegativities of
the two atoms in a covalent bond, the more polar is the bond.
Which pair will form the most polar covalent bond between
the atoms?
A. chlorine and bromine
B. chlorine and iodine
C. ﬂuorine and chlorine
D. ﬂuorine and iodine

Intermolecular forces
• Intermolecular attractions are attractions between
one molecule and a neighbouring molecule.
• The forces of attraction which hold an individual
molecule together are the intramolecular attractions
(chemical bonds - like covalent bonds!)

Intermolecular forces - van
der Waals forces
• Between simple covalent molecules and also separate atoms in
noble gases

• VDW (monatomic noble gas atoms) < VDW (simple covalent
molecules) despite similar electronic count [eg. He vs H2]

• In any molecule, the electrons are moving constantly and
randomly. Electron density can ﬂuctuate and parts of the
molecule become more or less negative - small temporary
dipoles

• These instantaneous dipole induces dipoles of opposite sign
on neighbouring molecules.

der Waals forces
• Two factors affecting strength of VDW:
a) the number of electrons in the molecule
=> the more electrons, the bigger the electron cloud, the
more polarisable is the electron cloud, the stronger is the id-
id interactions
b) the surface area for contact of the molecule
=> the greater the surface area of contact possible between
the molecules, the greater is the extent of the id-id
interactions

der Waals forces
Room temperature is 298 K (25 °C). bp of Br2 is 59 °C
while bp of Cl2 is -34 °C.
Account for why bromine is a liquid at r.t.p. while
chlorine is a gas.

der Waals forces

der Waals forces
Pentane has lower boiling point than its structural
isomer, dimethylpropane.
bp of pentane is 36 °C while bp of dimethylpropane
is 10 °C.

der Waals forces
RMM of CH4 is 16, that for N2 is 28. Explain why CH4
has higher bp than N2.

Intermolecular forces -
permanent dipoles
• Only between polar molecules
• Uneven distribution of electrons in polar bonds,
permanent separation of charges (dipoles) found
within polar molecules
• Two criteria required,
• a) there must be polar bonds within the molecules.
• b) there must be a net dipole moment for the
molecule.

hydrogen bond
• Present between molecules that have at least one
highly electronegative atom - F, O or N -
covalently bonded to an H atom.
• Sketch hydrogen bonding.

hydrogen bond
• Strength of a hydrogen bond depends on:
a) dipole moment of the H-X bond [X = F, O, N]
F-H - - - F-H > O-H - - - O-H > N-H - - - N-H
b) ease of donation of a lone pair on Y [Y= F, O, N]
N-H - - - N-H > O-H - - - O-H > F-H - - - F-H
• Overall, hydrogen bond strength is in this order,
F-H - - - F-H > O-H - - - O-H > N-H - - - N-H

hydrogen bond
• In terms of boiling point,
HF: 20 °C [But HF forms the strongest hydrogen bonds, so why?]
H2O: 100 °C
H3N: - 33 °C
• H2O has the highest bp because it can form more extensive
hydrogen bonding than NH3 and HF.
• H2O can form two hydrogen bonds per water molecule
whereas both NH3 and HF can only form one hydrogen bond
per molecule (why? sketch?)

hydrogen bond

Hydrogen bonding -
implications/importance
• Dimerisation of short-chain carboxylic acid

Hydrogen bonding -
• High solubility of ammonia and short-chain
alcohols (e.g. methanol, ethanol) in water

Hydrogen bonding -
• Intermolecular vs. intramolecular hydrogen
bonding
• bp of 2-nitrophenol (214 °C) vs. 4-nitrophenol (259 °C)

Hydrogen bonding -
• Ice ﬂoats on water - ice is less dense than water.

Hydrogen bonding -
• Stabilisation of structure of proteins - alpha helix and beta
sheets
• Different base pairings (A’ Level biology - Adenine-Thymine
Guanine-Cytosine)

Intermolecular forces - Q1
A crystal of iodine produces a purple vapour when
gently heated. Which pair of statements correctly
describes this process?
type of bond broken formula of purple species
A covalent I
B covalent I2
C induced dipole-dipole I2
D permanent dipole-dipole I2

Which types of intermolecular forces can exist between
adjacent urea molecules?
1. hydrogen bonding
2. permanent dipole-dipole forces
3. temporary induced dipole-dipole forces
e responses A to D should be selected on the basis of
A B C D
1, 2 and 3
are
correct
1 and 2
only are
correct
2 and 3
only are
correct
1 only
is
correct
other combination of statements is used as a correct response.
Which types of intermolecular forces can exist between adjacent urea molecules?
H2N
C
O
NH2
urea
1 hydrogen bonding
2 permanent dipole-dipole forces
3 temporary induced dipole-dipole forces
Ethanol is manufactured by reacting ethene gas and steam in the presence of phosphoric(V)
acid.
C H (g) + H O(g) C H OH(g) ∆H = –45kJmol–1

What is involved when a hydrogen bond is formed
between two molecules?
1. a hydrogen atom bonded to an atom less
electronegative than itself
2. a lone pair of electrons
3. an electrostatic attraction between opposite
charges

The three statements that follow are all true.
Which of these can be explained, at least in part, by
reference to hydrogen bonding?
1. At 0 °C, ice ﬂoats on water.
2. The boiling point of propan-2-ol is 82 °C. The
boiling point of propanone is 56 °C.
3. At 20 °C, propanone and propanal mix completely.

Which compound is the only gas at room temperature
and pressure?
A CH3CH2CH2NH2 Mr = 59.0
B CH3CH2CH2OH Mr = 60.0
C CH2OHCH2OH Mr = 62.0
D CH3CH2Cl Mr = 64.5

Metallic bonding
• Metallic bond is the electrostatic attraction between the
positive ions and the delocalised valence electrons
• Strength of metallic bonding depends on:
• a) number of valence electrons available for bonding (across
a period, how?)
• b) size of the metal cation (down a group, how?)

Metallic bonding
• Malleable and ductile

=> Ability of cations to move over one another without breaking
metallic bonds

• High mp and bp

=> Large amount of energy required to overcome strong metallic bond
(electrostatic attraction) between the positively charged ions and the
“sea of delocalised valence electrons”.

• Good thermal and electrical conductivity

=> Due to presence of delocalised electrons in the metallic lattice

Metallic bonding - Q1
Which of the following are features of the structure of
metallic copper?
(1) ionic bonds
(2) delocalised electrons
(3) lattice of ions

Csonn t3 chemical bonding

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