Bonding.ppt chemistry for preparatory students

Bonding

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Types of bonding
an ionic bond
occurs between
a metal and
non-metal atom
(e.g. NaCl)
a covalent
bond occurs
between two
non-metal atoms
(e.g. I
2
, CH
4
)
a metallic bond
occurs between
atoms in a metal
(e.g. Cu)
There are three types of bond that can occur between atoms:

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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 elements in groups 4 and 8 (also called group 0) do
not gain or lose electrons to form ionic compounds.
12345678/0
1+2+3+N/A3-2-1-N/A
Na
+
Al
3+
N/AN/AN
3-
O
2-
F
-
Mg
2+
Group
Charge
Example
The number of electrons gained or lost by an atom is
related to the group in which the element is found.

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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
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.
2. the size of the metal ions.
NaMg Al
1+2+3+
371923933
Element
Charge on ion
Melting point (K)

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Strength of metallic bonding: ion size
Element
Ionic radius
(nm)
Melting
point (K)
Li Na K Rb Cs
0.0760.1020.1380.1520.167
454 371 337 312302
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.

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Types of bonding

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Electronegativity
values for some
common elements.
Values given here
are measured on
the Pauling scale.
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.
What is electronegativity?

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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
lower
electronegativity
Electronegativity and atomic radius

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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
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.
Electronegativity decreases down a group because:
1. The atomic radius increases.

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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 O
2 or Cl
2) is always
non-polar because the
electronegativity of the atoms
in each molecule is the same.
both atoms are
equally good at
attracting the
electron density
cloud of 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.
decreasing polarization
This can be illustrated by looking at the hydrogen halides:
H F ClBr I
Molecule
Electronegativity
difference
between atoms
H–FH–ClH–BrH–I
1.81.0 0.8 0.5
Pauling
elecronegativities
Element
2.24.03.23.02.7

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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.
increasing polarization
More polar bonds have more
ionic character. The more
electronegative atom attracts the
electrons in the bond enough to
ionize the other atom.

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Polar molecules
Molecules containing polar bonds are not always polar.
If the polar bonds are
arranged symmetrically,
the partial charges cancel
out and the molecule is
non-polar.
Non-polar molecules
If the polar bonds are
arranged asymmetrically,
the partial charges do not
cancel out and the
molecule is polar.
Polar molecules

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Identifying polar molecules

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Types of intermolecular force
hydrogen bonds – for example, found between
H
2
O molecules in water.
permanent dipole–dipole forces – for example, found
between HCl molecules in hydrogen chloride.
van der Waals forces – for example, found between
I
2
molecules in iodine crystals.
There are three main 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.

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Van der Waals forces

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Strength of van der Waals forces
The strength of
van der Waals
forces increases
as molecular size
increases.
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.
0
50
100
150
200
-50
-100
-150
-200
b
o
i
l
i
n
g

p
o
i
n
t

(
°
C
)
Br
2
This is illustrated
by the boiling
points of group 7
elements.
F
2
Cl
2
I
2
element

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Strength of van der Waals forces
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.
butane (C
4H
10)
boiling point = 272 K
2-methylpropane (C
4
H
10
)
boiling point = 261 K
The points of contact between molecules also affects the
strength of 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.
Hydrogen bonds are therefore particularly strong examples
of permanent dipole–dipole forces.
The H atom is very small so the positive charge is more
concentrated, making it easier to link with other molecules.

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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 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.
hydrogen
bond lone pair

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Hydrogen bonding and boiling points

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Boiling points of the hydrogen halides
The boiling point
of hydrogen
fluoride is much
higher than that of
other hydrogen
halides, due to
fluorine’s high
electronegativity.
0
20
40
-20
-40
-60
-80
-100
b
o
i
l
i
n
g

p
o
i
n
t

(
°
C
)
HF HCl HBr HI
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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Permanent dipole–dipole forces

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Glossary

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What’s the keyword?

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Multiple-choice quiz