Ionic bonding: The electrostatic force of attraction between the
oppositely charged ions due to the transfer of a electrons from the
metal to the non-metal in the lattice.
Ionic bonds occur between the metals (lose electrons-cations) and
the non-metals (gain electrons-anions) to make an ion that is
isoelectronic (same number of electrons) with the nearest noble gas.
When a metal loses an electron, it becomes isoelectronic with the
noble gas in the period before. When a non-metal gains an electron,
it becomes isoelectronic with the noble gas in the same period.
Ionic lattice is a regular arrangement of positive and negative ions
in a three-dimensional structure.
Note that:
o The number of electrons lost or gained is determined by the
electron configuration of an atom.
o The transition metals can form more than one ion.
o There are no individual units of a bond (NaCl) as it exists as a
giant structure.
1) Factors affecting the strength of the ionic bond:
a. The charge of the ion: the larger the charge, the stronger
the ionic bond.
b. The number of main energy levels (ionic radius): the smaller
the number of main energy levels, the stronger the bond.
c. Charge density (charge\ionic radius): the charge density is
directly proportional to the charge of the positive of
negative ions and inversely proportional to the ionic radius
of the ion.
2) Properties of ionic bonds:
a. High melting and boiling points: due to the strong
electrostatic force of attraction between the oppositely
charged ions in the lattice, which require a high amount of
energy to break the bonds.
b. Solids at room temperature: due to the strong electrostatic
force of attraction between the oppositely charged ions in
the lattice.
c. Low volatility (a measure of how readily a substance
evaporates): due to the strong electrostatic force of
attraction between the oppositely charged ions in the
lattice.
d. Brittle (shatter when hit): when the lattice is hit, a layer
of ions is shifted so that the ions with the same charges
are lined up together; there like charges repel each other,
splitting the lattice.
e. Electrical conductivity of ionic compounds: they do not
conduct electricity when solids because the ions are packed
together and cannot move, but only vibrate. When molten, the
, ions are free to move due the break of the lattice. As they
are charged particles, they can carry an electric current.
f. Electrical conductivity in ionic solutions: aqueous
solutions of ionic substances conduct electricity because of
the freely moving ions that can carry current.
g. Soluble in water: water molecules have a slight electrical
charge and can attract the ions away from the lattice (non-
soluble in non-polar compounds).
Metallic bonding: the electrostatic force of attraction between the
positive metal ions and the delocalized electrons.
Metallic structure: metals contain a regular lattice
arrangement formed from layers of positive metal ions
surrounded by a sea of electrons that are freely moving
(delocalized).
1) Factors affecting the strength of the metallic bond:
a. Number of valence/ delocalized electrons: as the number
of delocalized electrons increases, the attraction
between the positive metal ion and the valence electron
increases, so the metallic bonding is stronger.
b. Ionic radius: the greater the ionic radius, the less
attraction for the valence electrons, resulting in a
weaker metallic bond.
2) Properties of metallic bonding:
a. Good conductors of heat and electricity: because of the
freely moving electrons that can carry charge.
b. High melting and boiling points: because of the strong
forces of attraction between the positive metal ion and
the delocalized electron.
c. Ductile and malleable: when a force is applied, the rows
of the positive metal ions slide past each other. The
delocalized electrons move along with them and bonds
reform (the structure is exactly the same as the
original). Because of the non-directionality of the
bonding, the metal ions in the lattice attract the
delocalized electrons in all directions.
3) Alloys: homogeneous mixtures of two or more metals or of a metal
with a non-metal. They are stronger/ stiffer than pure metals
since the presence of atoms of different size will prevent the
planes of metal atoms to slide.
oppositely charged ions due to the transfer of a electrons from the
metal to the non-metal in the lattice.
Ionic bonds occur between the metals (lose electrons-cations) and
the non-metals (gain electrons-anions) to make an ion that is
isoelectronic (same number of electrons) with the nearest noble gas.
When a metal loses an electron, it becomes isoelectronic with the
noble gas in the period before. When a non-metal gains an electron,
it becomes isoelectronic with the noble gas in the same period.
Ionic lattice is a regular arrangement of positive and negative ions
in a three-dimensional structure.
Note that:
o The number of electrons lost or gained is determined by the
electron configuration of an atom.
o The transition metals can form more than one ion.
o There are no individual units of a bond (NaCl) as it exists as a
giant structure.
1) Factors affecting the strength of the ionic bond:
a. The charge of the ion: the larger the charge, the stronger
the ionic bond.
b. The number of main energy levels (ionic radius): the smaller
the number of main energy levels, the stronger the bond.
c. Charge density (charge\ionic radius): the charge density is
directly proportional to the charge of the positive of
negative ions and inversely proportional to the ionic radius
of the ion.
2) Properties of ionic bonds:
a. High melting and boiling points: due to the strong
electrostatic force of attraction between the oppositely
charged ions in the lattice, which require a high amount of
energy to break the bonds.
b. Solids at room temperature: due to the strong electrostatic
force of attraction between the oppositely charged ions in
the lattice.
c. Low volatility (a measure of how readily a substance
evaporates): due to the strong electrostatic force of
attraction between the oppositely charged ions in the
lattice.
d. Brittle (shatter when hit): when the lattice is hit, a layer
of ions is shifted so that the ions with the same charges
are lined up together; there like charges repel each other,
splitting the lattice.
e. Electrical conductivity of ionic compounds: they do not
conduct electricity when solids because the ions are packed
together and cannot move, but only vibrate. When molten, the
, ions are free to move due the break of the lattice. As they
are charged particles, they can carry an electric current.
f. Electrical conductivity in ionic solutions: aqueous
solutions of ionic substances conduct electricity because of
the freely moving ions that can carry current.
g. Soluble in water: water molecules have a slight electrical
charge and can attract the ions away from the lattice (non-
soluble in non-polar compounds).
Metallic bonding: the electrostatic force of attraction between the
positive metal ions and the delocalized electrons.
Metallic structure: metals contain a regular lattice
arrangement formed from layers of positive metal ions
surrounded by a sea of electrons that are freely moving
(delocalized).
1) Factors affecting the strength of the metallic bond:
a. Number of valence/ delocalized electrons: as the number
of delocalized electrons increases, the attraction
between the positive metal ion and the valence electron
increases, so the metallic bonding is stronger.
b. Ionic radius: the greater the ionic radius, the less
attraction for the valence electrons, resulting in a
weaker metallic bond.
2) Properties of metallic bonding:
a. Good conductors of heat and electricity: because of the
freely moving electrons that can carry charge.
b. High melting and boiling points: because of the strong
forces of attraction between the positive metal ion and
the delocalized electron.
c. Ductile and malleable: when a force is applied, the rows
of the positive metal ions slide past each other. The
delocalized electrons move along with them and bonds
reform (the structure is exactly the same as the
original). Because of the non-directionality of the
bonding, the metal ions in the lattice attract the
delocalized electrons in all directions.
3) Alloys: homogeneous mixtures of two or more metals or of a metal
with a non-metal. They are stronger/ stiffer than pure metals
since the presence of atoms of different size will prevent the
planes of metal atoms to slide.
