Download Bonding-and-Intermolecular-Forces

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Types of bonding
There are three types of bond that can occur between atoms:

an ionic bond
occurs between
a metal and
non-metal atom
(e.g. NaCl)
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
a covalent
bond occurs
between two
non-metal atoms
(e.g. I2, CH4)

a metallic bond
occurs between
atoms in a metal
(e.g. Cu)
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Ionic bonding
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Charge on the ions
Metals lose electrons to form positive ions while
non-metals gain electrons to form negative ions.
The number of electrons gained or lost by an atom is
related to the group in which the element is found.
Group
1
2
3
4
5
6
7
Charge
1+
2+
3+ N/A 3-
2-
1- N/A
Example Na+ Mg2+ Al3+ N/A N3- O2-
8/0
F- N/A
The elements in groups 4 and 8 (also called group 0) do
not gain or lose electrons to form ionic compounds.
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Representing ionic bonding
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Covalent bonding
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Co-ordinate bonding
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Examples of co-ordinate bonds
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Co-ordinate bonds: true or false?
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Metallic bonding
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Strength of metallic bonding: ion charge
The strength of metallic bonding depends on two factors:
1. the charge on the metal ions
2. the size of the metal ions.
1. The charge on the metal ions
The greater the charge on the metal ions, the greater the
attraction between the ions and the delocalized electrons,
and the stronger the metallic bonds. A higher melting
point is evidence of stronger bonds in the substance.
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Element
Na
Mg
Al
Charge on ion
1+
2+
3+
Melting point (K) 371
923
933
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Strength of metallic bonding: ion size
2. The size of the metal ions
The smaller the metal ion, the closer the positive
nucleus is to the delocalized electrons. This means
there is a greater attraction between the two, which
creates a stronger metallic bond.
Element
Li
Ionic radius
0.076
(nm)
Melting
point (K)
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454
Na
K
Rb
Cs
0.102
0.138
0.152
0.167
371
337
312
302
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Types of bonding
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What is electronegativity?
In a covalent bond between two different elements, the
electron density is not shared equally.
This is because different elements have differing abilities to
attract the bonding electron pair. This ability is called an
element’s electronegativity.
Electronegativity
values for some
common elements.
Values given here
are measured on
the Pauling scale.
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Electronegativity and atomic radius
The electronegativity of an element depends on a
combination of two factors:
1. Atomic radius
As radius of an atom increases, the bonding pair of
electrons become further from the nucleus. They are
therefore less attracted to the positive charge of the
nucleus, resulting in a lower electronegativity.
higher
electronegativity
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lower
electronegativity
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Electronegativity, protons and shielding
2. The number of unshielded protons
The greater the number of protons in a nucleus, the
greater the attraction to the electrons in the covalent
bond, resulting in higher electronegativity.
However, full energy levels of electrons shield the
electrons in the bond from the increased attraction of the
greater nuclear charge, thus reducing electronegativity.
greater nuclear
charge increases
electronegativity…
…but extra shell of
electrons increases
shielding.
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Electronegativity trends: across a period
Electronegativity increases across a period because:
1. The atomic radius decreases.
2. The charge on the nucleus increases without
significant extra shielding. New electrons do not
contribute much to shielding because they are added
to the same principal energy level across the period.
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Electronegativity trends: down a group
Electronegativity decreases down a group because:
1. The atomic radius increases.
2. Although the charge on the nucleus increases,
shielding also increases significantly. This is
because electrons added down the group fill new
principal energy levels.
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Non-polar bonds
If the electronegativity of both atoms in a covalent bond is
identical, the electrons in the bond will be equally attracted
to both of them.
This results in a symmetrical
distribution of electron
density around the two
atoms.
Bonding in elements (for
example O2 or Cl2) is always
non-polar because the
electronegativity of the atoms
in each molecule is the same.
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cloud of electron density
both atoms are
equally good at
attracting the
electron density
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Polar bonds
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Effect of electronegativity on polarization
The greater the electronegativity difference between the two
atoms in a bond the greater the polarization of the bond.
This can be illustrated by looking at the hydrogen halides:
Element
Pauling
elecronegativities
H
F
Cl
Br
I
2.2
4.0
3.2
3.0
2.7
Molecule
H–F
H–Cl
H–Br
H–I
Electronegativity
difference
between atoms
1.8
1.0
0.8
0.5
decreasing polarization
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Ionic or covalent?
Rather than saying that ionic and covalent are two distinct
types of bonding, it is more accurate to say that they are at
the two extremes of a scale.
Less polar bonds have
more covalent
character.
More polar bonds have more
ionic character. The more
electronegative atom attracts the
electrons in the bond enough to
ionize the other atom.
increasing polarization
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Polar molecules
Molecules containing polar bonds are not always polar.
Non-polar molecules
Polar molecules
If the polar bonds are
arranged symmetrically,
the partial charges cancel
out and the molecule is
non-polar.
If the polar bonds are
arranged asymmetrically,
the partial charges do not
cancel out and the
molecule is polar.
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Identifying polar molecules
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Types of intermolecular force
The molecules in simple covalent substances are not entirely
isolated from one another. There are forces of attraction
between them. These are called intermolecular forces.
There are three main types of intermolecular force:

van der Waals forces – for example, found between
I2 molecules in iodine crystals.

permanent dipole–dipole forces – for example, found
between HCl molecules in hydrogen chloride.

hydrogen bonds – for example, found between
H2O molecules in water.
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Van der Waals forces
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The strength of
van der Waals
forces increases
as molecular size
increases.
This is illustrated
by the boiling
points of group 7
elements.
boiling point (°C)
Strength of van der Waals forces
200
150
100
50
0
-50
-100
-150
-200
F2
Cl2
Br2
element
I2
Atomic radius increases down the group, so the outer
electrons become further from the nucleus. They are
attracted less strongly by the nucleus and so temporary
dipoles are easier to induce.
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Strength of van der Waals forces
The points of contact between molecules also affects the
strength of van der Waals forces.
butane (C4H10)
2-methylpropane (C4H10)
boiling point = 272 K
boiling point = 261 K
Straight chain alkanes can pack closer together than
branched alkanes, creating more points of contact between
molecules. This results in stronger van der Waals forces.
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Boiling points of alkanes
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Permanent dipole–dipole forces
If molecules contain bonds with a permanent dipole, the
molecules may align so there is electrostatic attraction
between the opposite charges on neighbouring molecules.
Permanent
dipole–dipole
forces (dotted
lines) occur in
hydrogen chloride
(HCl) gas.
The permanent dipole–dipole forces are approximately
one hundredth the strength of a covalent bond.
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Permanent dipole–dipole or not?
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What is hydrogen bonding?
When hydrogen bonds to nitrogen, oxygen or fluorine, a
larger dipole occurs than in other polar bonds.
This is because these atoms are
highly electronegative due to their
high nuclear charge and small size.
When these atoms bond to hydrogen,
electrons are withdrawn from the H
atom, making it slightly positive.
The H atom is very small so the positive charge is more
concentrated, making it easier to link with other molecules.
Hydrogen bonds are therefore particularly strong examples
of permanent dipole–dipole forces.
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Hydrogen bonding
In molecules with OH
or NH groups, a lone
pair of electrons on
nitrogen or oxygen is
attracted to the slight
positive charge on the
hydrogen on a
neighbouring molecule.
hydrogen
bond
lone pair
Hydrogen bonding makes the melting and boiling points of
water higher than might be expected. It also means that
alcohols have much higher boiling points than alkanes of a
similar size.
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Hydrogen bonding and boiling points
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Boiling points of the hydrogen halides
40
20
0
-20
-40
-60
-80
-100
HF
boiling point (°C)
The boiling point
of hydrogen
fluoride is much
higher than that of
other hydrogen
halides, due to
fluorine’s high
electronegativity.
HCl
HBr
The means that hydrogen bonding between molecules of
hydrogen fluoride is much stronger than the permanent
dipole–dipole forces between molecules of other
hydrogen halides. More energy is therefore required to
separate the molecules of hydrogen fluoride.
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HI
Permanent dipole–dipole forces
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Glossary
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What’s the keyword?
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Multiple-choice quiz
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