REVIEWTOPIC 4/14: BONDING
A. KEY DEFINITIONS (MEMORIZE!)
*No explicit stating of definitions required in this unit.
B. YOU SHOULD KNOW:
 Ionic bonds form by electron transfer between elements of differing electronegativity, so between metals and non-metals
 Ionic bonding is due to the electrostatic attraction of oppositely charged ions
 Ionic bonding leads to the formation of a regular framework of alternating positive and negative ions described as a lattice
 Covalent bonds form when electrons are shared between nuclei
 Single bonds are longer and weaker than double bonds, and triple bonds are shorter and stronger than double bonds, so CC is more easily broken than C=C, which is more easily broken than the carbon – carbon triple bond
 Coordinate covalent bonds form when both the electrons being shared have come from the same atom
 The structure and bonding in the allotropes of carbon, silicon and silicon dioxide, as in the table below.
Diamond (C)
Graphite (C)
Fullerene (C60)
Silicon (Si)
Silicon
Graphene
dioxide
(SiO2)
Structure
Giant: all C
Giant: all C
Molecular: all
Giant: all Si
Giant: all Si
Giant: one
atoms bonded atoms bonded C atoms
atoms bonded atoms
layer of
to 4 others
to 3 others in
bonded to 3
to 4 others
bonded to 4
graphite
a layer
others in a
oxygens,
spherical
each oxygen
shape
bonded to 2
silicons
Bonding
Covalent
Covalent
Covalent
covalent
Covalent
Covalent
Shape and
Carbons
Carbons are
Carbons in
Silicons
Silicons
Carbons in
bond
arranged
trigonal planar trigonal planar arranged
arranged
trigonal planar
angles
tetrahedrally,
arrangement,
arrangement,
tetrahedrally,
tetrahedrally, arrangement,
bond angle is
bond angle is
bond angle is
bond angle is
bond angle is bond angle is
109.5°
120°
120°
109.5°
109.5°
120°
 Sigma (  ) bonds are formed by the axial overlap (along inter-nuclear axis) of orbitals
 Pi (  ) bonds are formed by the sideways overlap (above and below inter-nuclear axis) of parallel p-orbitals
 Single bonds are sigma bonds
 Double bonds contain a sigma bond and a pi bond
 Triple bonds contain a sigma bond and two pi bonds
 Metallic bonds are formed by the electrostatic attraction of delocalized electrons for the lattice of positive metallic ions
 Intermolecular forces hold molecules together and are much weaker than covalent bonds
 The strength of intermolecular forces/bonds are in the order:
hydrogen bonds > dipole-dipole forces > temporary dipole/London forces
 That physical properties depend on bonding and structure, as shown in the following table:
Property
Structure
Ionic
Simple Molecular
Giant molecular
Metallic
Melting and boiling
High
Low
High
High
point
Volatility
Low
High
Low
Low
Electrical
Good as liquid, none None
None (*exception =
Good
conductivity
as solid
graphite)
Solubility in polar
Good
Poor
Poor
Poor
solvent
Solubility in nonPoor
Good
Poor
Poor
polar solvent
C. YOU SHOULD BE ABLE TO:
 Deduce the formula of ionic compounds by knowing the details in the table:
Group number
Electrons lost
Charge on ion
Group number
1
1
1+
5
2
2
2+
6
3
3
3+
7
Electrons gained
3
2
1
Charge on ion
321-
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Deduce the Lewis structures of simple molecules and different resonance forms, where applicable.
Deduce the formal charge of atoms in molecules.
Use valence shell electron pair repulsion (VSEPR) theory to predict the bond angle(s), electron domains and molecular
geometry of a molecule or ion
Predict whether a bond will be polar by looking at the electronegativities of the atoms (the bigger the difference, the more
polar the bond)
Predict whether a molecule is polar by looking at the bond polarity and shape of the molecule – in a polar molecule the
dipoles do not cancel out (they are not symmetrical)
Explain that, when hybridization occurs, atomic orbitals merge to create new orbitals of intermediate energy
Explain the hybridization in methane, ethane and ethyne
Explain that metals conduct electricity because of their delocalized electrons
Explain that metals are malleable because there are no directional bonds to be broken – the ions can slide over each other
when a force is applied.
Deduce the intermolecular force(s) present between molecules from looking at the structure of the molecules
Predict relative boiling points.
Explain the wavelength of light required to dissociate oxygen and ozone
Describe the mechanism of the catalysis of ozone depletion when catalyzed by CFCs and oxides of nitrogen
Describe the properties of an alloy ; provide examples of alloys ; define alloy
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D. BE PREPARED
 You should know and be able to use the formulas for carbonate (CO32-), nitrate (NO3-), sulfate (SO42-), hydroxide (OH-),
phosphate (PO43-), hydrogencarbonate (HCO3-) and ammonium (NH4+).
 In Lewis structures, a pair of electrons can be represented by two dots or two crosses or a line.
 Make sure that you include all the valence electrons of all atoms.
 Know that lone pairs of electrons repel more than bonding pairs and this distorts the shape of the molecule and changes
the bond angles within it, as show in the table below.
Number of
Arrangement
Number of
Number of
Shape of
Bond angle
Examples
electron pairs of electron
bonding pairs lone pairs of
molecule
(degrees)
pairs
of electrons
electrons
2
Linear
2
0
Linear
180
CO2, HCN
3
4
5
6
Trigonal
planar
Tetrahedral
Trigonal
bipyramidal
Octahedral
3
0
2
1
4
3
0
1
2
2
5
0
4
3
1
2
6
5
0
1
Trigonal
planar
Bent
120
BCl3, AlCl3
118
SO2
Tetrahedral
Trigonal
pyramidal
Bent
109.5
>109.5
(approx. 107)
>109.5
(approx, 104)
Trigonal
bypyramidal
See-saw
T-shaped
120 and 90
PCl5
90 & 117
90
SF4
ClF3
90
90
SF6
BrF5
90
XeF4
Octahedral
Square
pyramindal
4
2
Square planar
 Know how shape, bond angle and hybridization are linked, as in the table.
Hybridization
Molecular shape
Bond angle (degrees)
3
sp
Tetrahedral
109.5
sp2
Trigonal planar
120
sp
Linear
180
 Appreciate the economic importance of iron and other metals
 Hydrogen bonding can only occur when hydrogen is directly bonded to F, O or N.
CH4, SiF4
NH3, PCl3
H2O, SCl2
Examples
CH4, NH3
C2H4, AlCl3
C2H2, CO2
PROBLEM SET #4 (TOPIC 4/14)
1. Draw the Lewis structures, state the shapes and predict the bond angles for the following species. (3 each)
(i) SiF62(ii) NO2+
2. The table below shows the approximate boiling points of the hydrides of group 5. Discuss the variation in the boiling points. (4)
PERIOD
HYDRIDE
BOILING
POINT (°C)
2
NH3
-32
3
PH3
-88
4
AsH3
-57
5
SbH3
-20
3. Explain, using diagrams, why NO2 is a polar molecule but CO2 is a non-polar molecule. (3)
4. Describe the structure and bonding in silicon dioxide. (2)
5. Consider the molecule HCONH2.
(i) State the name of the compound and draw its structural formula, showing all the bonds present. (2)
(ii) Explain the term hybridization. (1)
(iii) Describe how π and σ bonds form. (2)