Topic 4 Bonds.notebook
December 03, 2016
Topic 4 Chemical Bonding
and Structure
4.1 Ionic bonding and structure
4.2 Covalent bonding
4.3 Covalent structures
4.4 Intermolecular forces
4.5 Metallic bonding
Dec 2­1:09 PM
4.1 Ionic bonding and structure
Understandings:
4.1.1. Positive ions (cations) form by metals losing
valence electrons.
4.1.2. Negative ions (anions) form by non­metals
gaining electrons.
4.1.3. The number of electrons lost or gained is
determined by the electron configuration of the
atom.
4.1.4. The ionic bond is due to electrostatic
attraction between oppositely charged ions.
4.1.5. Under normal conditions, ionic compounds
are usually solids with lattice structures.
Dec 2­1:14 PM
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4.2 Covalent bonding
Understandings:
4.2.1. A covalent bond is formed by the electrostatic
attraction between a shared pair of electrons and the
positively charged nuclei.
4.2.2. Single, double and triple covalent bonds involve
one, two and three shared pairs of electrons
respectively.
4.2.3. Bond length decreases and bond strength
increases as the number of shared electrons increases.
4.2.4. Bond polarity results from the difference in
electronegativities of the bonded atoms.
Dec 2­1:17 PM
4.3 Covalent structures
Understandings:
4.3.1. Lewis (electron dot) structures show all the valence
electrons in a covalently bonded species.
4.3.2. The “octet rule” refers to the tendency of atoms to
gain a valence shell with a total of 8 electrons.
4.3.3. Some atoms, like Be and B, might form stable
compounds with incomplete octets of electrons.
4.3.4. Resonance structures occur when there is more than
one possible position for a double bond in a molecule.
4.3.5. Shapes of species are determined by the repulsion of
electron pairs according to VSEPR theory.
4.3.6. Carbon and silicon form giant covalent/network
covalent structures.
Dec 2­1:20 PM
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4.4 Intermolecular forces
Understandings:
4.4.1. Intermolecular forces include London
(dispersion) forces, dipole­dipole forces and
hydrogen bonding.
4.4.2. The relative strengths of these
interactions are London (dispersion) forces
< dipole­dipole forces < hydrogen bonds.
Dec 2­1:25 PM
4.5 Metallic bonding
Understandings:
4.5.1. A metallic bond is the electrostatic
attraction between a lattice of positive ions
and delocalized electrons.
4.5.2. The strength of a metallic bond
depends on the charge of the ions and the
radius of the metal ion.
4.5.3. Alloys usually contain more than one
metal and have enhanced properties.
Dec 2­1:25 PM
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4.1 The octet rule
Electron configuration of noble gases (He, Ne, Ar).
Dec 13­10:36 AM
4.1 The octet rule
The octet rule states that atoms lose or gain electrons
to achieve the electron configuration of noble gases.
Atoms can achieve the electron configuration of a noble
gas by either sharing electrons (covalent bonding) or by
losing or gaining electrons (ionic bonding).
Nov 26­6:40 PM
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Exceptions to the octet rule
Nov 26­6:44 PM
4.1 Elecrtonegativity and bonding
Electronegativity is a measure of the attraction of an atom for
a bonding pair of electrons.
It increases across a period due to increasing nuclear charge,
and decreases down a group as atomic radius increases.
0 < χρ < 0.4
0.5 < χρ < 1.8
n o n­p o l a r
p..o...l...a...r
co..va..lent
0
ionic bond
1.8
Nov 26­6:47 PM
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4.1 Elecrtonegativity and bonding
3.2
3.2
2.2
non­polar covalent bond
3.2
polar covalent bond
3.2
0.9
ionic bonding
Nov 26­6:45 PM
4.1 Elecrtonegativity and bonding
Ionic bonding takes place between metals and non­metals.
Covalent bonding takes place between non­metals.
Dec 13­10:50 AM
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4.1 Ions: Positive Ions
Dec 13­10:57 AM
4.1 Ions: Negative Ions
Dec 13­10:58 AM
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4.1 Ions and ion formation
group number
number of
charge on ion
valence electrons
example
1
1
1+
Li+
2
2
2+
Mg2+
13
3
3+
Al3+
15
5
3­
P3­
16
6
2­
O2­
17
7
1­
F­
Worksheet 1­2
Nov 26­7:14 PM
4.1 Ionic bonding
An ionic bond is the electrostatic attraction
between oppositely chaged ions.
Nov 26­7:20 PM
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4.1 Ionic bonding
Ionic compound have a lattice structure.
The lattice is held together
by the positive and negative
charges of the oppositely
charged ions.
Ionic compounds are solids
under standard conditions ­
they have high melting and
boiling points.
Nov 26­7:21 PM
4.1 Properties of the ionic compounds
The ions in the lattice structure are held
in place by strong electrostatic
attractions.
Ionic compounds do not conduct
electricity when solid, they only conduct
electricity when molten or dissolved in
water.
Worksheet 3
When molten or dissolved, the ions are
free to move and conduct electricity.
Nov 26­7:24 PM
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4.1 Properties of the ionic compounds
Ionic compounds are soluble in
polar solvents. The ions are
separated from the lattice
structure by the polar water
molecules.
The ions are then surrounded
by water molecules (hydration).
Nov 26­7:28 PM
Effect of ionic charge on melting point
ionic compound
cation charge
anion charge
melting point
(oC)
NaCl
Na+
Cl­
801
MgO
Mg2+
O2­
2800
The greater the charge on the ion, the stronger
the electrostatic attraction between the oppositely
charged ions and the higher the melting point.
Nov 26­7:33 PM
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Effect of ionic radius on melting point
ionic compound
cation radius
(x10­12m)
anion radius
(x10­12m)
melting point
(oC)
NaF
102
133
992
KF
138
133
857
The greater the ionic radius of the ion, the weaker
the electrostatic attraction between the oppositely
charged ions and the lower the melting point
Nov 26­7:36 PM
4.1 Properties of the ionic compounds
Ionic compounds only conduct electricity when
molten or dissolved in solution.
They are soluble in polar solvents (such as H2O).
Ionic compounds have high melting points because
of the strong electrostatic attractions between ions.
The greater the charge on the ion and the smaller
the ionic radius, the higher the melting point
Nov 26­7:40 PM
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Writing formulae of ionic compounds
Nov 26­7:44 PM
Formulae of ionic compounds
Lithium fluoride
Li
+
­
F
Magnesium chloride
Mg
LiF
2+
Cl
­
MgCl2
Nov 26­8:00 PM
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Formulae of ionic compounds
Aluminium bromide
Al
3+
­
Br
Iron (III) oxide
Fe3+ O2­
AlBr3
Fe2O3
Nov 26­8:00 PM
Formulae of ionic compounds
Barium hydroxide
Aluminium sulfate
Ba2+ OH­
Al3+ SO42­
Ba(OH)2
Al2(SO4)3
Nov 26­8:00 PM
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4.2 Covalent bonding
Covalent
bonding occurs
between non­
metal elements
A covalent bond is the electrostatic attraction
between positive nuclei and shared pair of
electrons.
Dec 3­2:12 PM
4.2 Covalent bonding
number of shared
electrons
C to C bond
strength (kJ mol­1)
C to C length
(10­12m)
2
347
153
4
614
134
6
839
120
Dec 3­2:14 PM
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4.2 Covalent bonding
3.2
3.2
2.2
non­polar covalent bond
3.2
polar covalent bond
+
­
H Cl
H Cl
Dec 3­2:24 PM
4.2 Covalent bonding
Covalent bonding occurs between non­metal
elements. It is the electrostatic attraction between
positive nuclei and shared pairs of electrons.
Single bonds are weaker and longer than double
or triple bonds which are stronger and shorter.
Covalent bonds can be polar or non­polar
depending on the difference in electronegativity
between the atoms.
Worksheet 4 & Worksheet 5
Dec 3­2:32 PM
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4.2 Non­polar molecules
The polarity of a molecule depends
on the presence of polar bonds
within the molecule and also on the
molecular geometry.
Dec 3­2:40 PM
4.2 Non­polar molecules
Non­polar molecule with polar bonds (CCl4)
In this molecule, the bond polarities cancel out due
to the molecular geometry, overall the molecule has
no net dipole moment.
Dec 3­2:52 PM
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4.2 Polar molecules
Polar molecule (CH2Cl2)
In this molecule, the bond polarities do not cancel out,
therefore it has a net dipole moment.
Dec 3­2:57 PM
Non­polar molecules
Polar molecules
­
+
+
­
+
+
+
Worksheet 6
Dec 3­3:10 PM
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4.3 Lewis structures
Lewis structures show all the valence electrons in a
molecule; the bonding electrons and the non­bonding
(lone pairs) of electrons.
Dec 3­3:28 PM
4.3 Lewis structures
1). Count the total number of valence electrons in all
the atoms in the molecule.
2). Determine the number of electrons needed for
each atom to achieve an octet.
3). Subtract 1 from 2 to get the number of bonding
electrons in the molecule.
4). Add electrons to each atom until it has an octet.
5). Count the total number of valence electrons, it
should be equal to the number in part 1.
Dec 3­3:38 PM
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4.3 Lewis structures
Methane (CH4)
1). 4+(4x1)=8 valence electrons
2). 8+(4x2)=16 electrons needed
to complete each atom's octet
3). 16­8=8 bonding electrons
4). Complete each atom's octet
5). Total number of electrons = 8
Dec 3­3:40 PM
4.3 Lewis structures
Dichloromethane (CH2Cl2)
1). 4+(2x1)+(2x7)=20 valence
electrons
2). 8+(2x2)+(2x8)=28 electrons
needed to complete each atom's
octet
3). 28­20=8 bonding electrons
4). Complete each atom's octet
5). Total number of electrons = 20
Dec 3­3:44 PM
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4.3 Lewis structures
Ammonia (NH3)
1). 5+(3x1)=8 valence electrons
2). 8+(3x2)=14 electrons needed
to complete each atom's octet
3). 14­8=6 bonding electrons
4). Complete each atom's octet
5). Total number of electrons = 8
Dec 3­3:45 PM
4.3 Lewis structures
Ethene (C2H4)
1). (2x4)+(4x1)=12 valence
electrons
2). (2x8)+(4x2)=24 electrons
needed to complete each atom's
octet
3). 24­12=12 bonding electrons
Worksheet 7
4). Complete each atom's octet
5). Total number of electrons = 12
Dec 3­3:50 PM
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4.3 VSEPR theory
Valence shell electron pair repulsion theory is used
to predict the geometry (shape) of molecules.
Electron pairs (bonds or lone pairs) repel each
other and spread apart as far as possible.
least repulsion
greatest repulsion
lone pair ­ lone pair > lone pair ­ bonding pair > bonding pair ­ bonding pair
The term electron domain is used to refer to bonds
or lone pairs of electrons around an atom in a
molecule.
Dec 3­3:56 PM
4.3 VSEPR theory
Single bonds, double bonds, triple bonds and
lone pairs of electrons count as one electron
domain.
4 electron domains
around the carbon atom
(4 bonding domains)
Dec 3­4:10 PM
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4.3 VSEPR theory
4 electron domains
around the oxygen
atom (2 bonding
domains, 2 lone pairs)
2 electron domains
around the carbon atom
(2 bonding domains)
Dec 3­4:13 PM
4.3 VSEPR theory
CH4
electron
domains
bonding
domains
lone pairs
electron domain
geometry
molecular
geometry
bond angle
4
4
0
tetrahedral
tetrahedral
109.5o
Dec 3­4:18 PM
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4.3 VSEPR theory
NH3
electron
domains
bonding
domains
lone pairs
electron domain
geometry
molecular
geometry
bond angle
4
3
1
tetrahedral
trigonal
pyramidal
107.8o
Dec 3­4:18 PM
4.3 VSEPR theory
H 2O
electron
domains
bonding
domains
lone pairs
electron domain
geometry
molecular
geometry
bond angle
4
2
2
tetrahedral
bent
104.5o
Water has a bent shape due to the extra repulsion from
the two lone pairs of electrons on the oxygen atom.
Dec 3­4:18 PM
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4.3 VSEPR theory
BF3
electron
domains
bonding
domains
lone pairs
3
3
0
electron domain
geometry
molecular
geometry
trigonal planar trigonal planar
bond angle
120o
Dec 3­4:18 PM
4.3 VSEPR theory
SO2
electron
domains
bonding
domains
lone pairs
electron domain
geometry
molecular
geometry
bond angle
3
2
1
trigonal planar
bent
<120o
Dec 3­4:18 PM
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4.3 VSEPR theory
CO2
electron
domains
bonding
domains
lone pairs
electron domain
geometry
molecular
geometry
bond angle
2
2
0
linear
linear
180o
Dec 3­4:18 PM
No lone
pairs of
electrons
around
central
atom.
Dec 3­4:37 PM
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Lone pairs of electrons affect the bond angles in a molecule.
Lone pairs occupy more space than bonding pairs, so they
decrease the bond angle between bonding pairs.
Dec 3­4:42 PM
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