H
H N H
Bond indicated
by an arrow
No charged particles
+
Power of an atom to attract the
electron density in a covalent bond
Strong
H
Shared pair of
electrons
A shared pair of electrons
where both electrons are
supplied by one atom
Coordinate
(dative covalent) bond
Covalent
bonds
Intermolecular
forces
Permanent dipole-dipole
forces
Polar covalent
bonds can cause a
permanent dipole
Uneven electron
distribution produces a
polar covalent bond
Electronegativity
Bond polarity
Covalent
Bonding
1.3 BONDING
Do not conduct
Contain a fixed number
of atoms covalently
bonded together
Weakest
Hydrogen
bonds
Strongest
Influences low
density of ice
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Strong covalent bonds
between atoms
Giant
structure
Shapes
Induced dipole-dipole
forces
Low melting and
boiling points
Very high melting
points
Molecular
structures
Macromolecular
structures
Weak intermolecular
forces
E.g. diamond,
graphite, silicon
dioxide
Shapes and bond angles of
simple molecules and ions
Molecules/ions with up to 6
electron pairs
Electron pair
repulsion
VSEPR principles
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Metal loses electron(s),
non-metal gains electron(s)
Electrostatic attraction
between oppositely
charged ions
Giant Ionic Lattice
Structure
1.3 BONDING:
IONIC BONDING
Properties
Conduct electricity when
molten or aqueous
E.g. SO42-, OH-, NO3-, NH4+
Formulas of
Common Ions
High melting and
boiling points
Very strong electrostatic
forces between ions
Lots of energy needed
to overcome forces
Ions are free to move to
carry charge
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Easily separated
Positive metal ions held in a
sea of delocalised electrons
Giant
Structure
1.3 BONDING:
METALLIC BONDING
Strong metallic bonds
Good conductors
Delocalised electrons
carry charge
Properties
High melting points
Malleable
Atoms arranged in a regular layers
which can slide over each other
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