Bonding and Structure
Ions and Ionic bonding
An ionic bond is the strong electrostatic force of attraction between oppositely
charged ions. They form between ions of atoms with significantly different
electronegativities.
The greater the ionic radius, the weaker the ionic bond. The greater the ionic charge,
the stronger the ionic bond.
Ions are formed by the gain or loss of electrons.
Going down a group, the ionic radius increases.
Isoelectronic ions are ions which have the same electronic structure, for example N 3-
and Al3+. These are also isoelectronic with the ions O2-, F-, Na+ and Mg2+. In
increasing atomic number of these ions (from N 3- with an atomic number of 7 and Al3+
with an atomic number of 13), the ionic radii decrease because the number of
protons increases from ion to ion, and so the nuclear charge increases from ion to
ion, and so with the number of electrons staying the same for each of the listed ions,
the attraction between the nucleus and the orbiting electrons (in the quantum shells)
increases from ion to ion, thus decreasing the atomic radius from ion to ion.
The evidence for the existence of ions is that ions are able to flow free in solution
across an applied potential, with the flow being a current caused by ions. Moving
free in solutions allows negatively charged ions (anions) to move towards the anode
and for positively charged ions (cations) to move towards the cathode during
electrolysis. Electrolysis shows a solution which conducts electricity (when a solid
form of the substance in solution wouldn’t conduct electricity) whilst products are
being formed at the anode and cathode.
Covalent bonding
A covalent bond is a strong electrostatic force of attraction between a shared pair of
electrons and their adjacent nuclei.
In a dative covalent bond (also known as a co-ordinate bond), the bonding pair of
electrons both come from one atom. Below on the left shows a co-ordinate bond in
NH4+ (the ammonium ion), and on the right Al2Cl6 (aluminium chloride in the gas
phase). In a Lewis structure diagram, the co-ordinate bond is represented by an
arrow from the atom which the bonding pair of electrons are coming from, to the
atom which that atom is bonding with by the co-ordinate bond.
, The longer the bond length, the weaker the bond. Bond length is affected by the
atomic radius – the larger the atomic radius, the higher the bond length. The stronger
the bond, the higher the bond energy. An exception is fluorine (F 2), where the bond
length is so small that there is significant repulsion between the lone electron pairs
on one atom and the lone electron pairs on the other, increasing the length of the
bond, and so making it weaker than chlorine (Cl 2), for example.
Shape of molecules
The shape of a simple molecule or ion is determined by the repulsion between the
electron pairs (both bonding and lone) that surround the central atom.
Total Number Number Shape of molecule Bond Example
number of of of lone angle
electron bonding pairs (o)
pairs regions
2 2 0 Linear 180 MgCl2
CO2
3 3 0 Trigonal planar 120 BCl3
SO3
CO32-
NO32-
..
2 1 Bent 119 SO2
NO2-
4 4 0 Tetrahedral 109.5 CH4
3 1 Trigonal pyramid 107.5 NH3
SO32-
H3O+
2 2 Bent 104.5 H2O
.. ..
Ions and Ionic bonding
An ionic bond is the strong electrostatic force of attraction between oppositely
charged ions. They form between ions of atoms with significantly different
electronegativities.
The greater the ionic radius, the weaker the ionic bond. The greater the ionic charge,
the stronger the ionic bond.
Ions are formed by the gain or loss of electrons.
Going down a group, the ionic radius increases.
Isoelectronic ions are ions which have the same electronic structure, for example N 3-
and Al3+. These are also isoelectronic with the ions O2-, F-, Na+ and Mg2+. In
increasing atomic number of these ions (from N 3- with an atomic number of 7 and Al3+
with an atomic number of 13), the ionic radii decrease because the number of
protons increases from ion to ion, and so the nuclear charge increases from ion to
ion, and so with the number of electrons staying the same for each of the listed ions,
the attraction between the nucleus and the orbiting electrons (in the quantum shells)
increases from ion to ion, thus decreasing the atomic radius from ion to ion.
The evidence for the existence of ions is that ions are able to flow free in solution
across an applied potential, with the flow being a current caused by ions. Moving
free in solutions allows negatively charged ions (anions) to move towards the anode
and for positively charged ions (cations) to move towards the cathode during
electrolysis. Electrolysis shows a solution which conducts electricity (when a solid
form of the substance in solution wouldn’t conduct electricity) whilst products are
being formed at the anode and cathode.
Covalent bonding
A covalent bond is a strong electrostatic force of attraction between a shared pair of
electrons and their adjacent nuclei.
In a dative covalent bond (also known as a co-ordinate bond), the bonding pair of
electrons both come from one atom. Below on the left shows a co-ordinate bond in
NH4+ (the ammonium ion), and on the right Al2Cl6 (aluminium chloride in the gas
phase). In a Lewis structure diagram, the co-ordinate bond is represented by an
arrow from the atom which the bonding pair of electrons are coming from, to the
atom which that atom is bonding with by the co-ordinate bond.
, The longer the bond length, the weaker the bond. Bond length is affected by the
atomic radius – the larger the atomic radius, the higher the bond length. The stronger
the bond, the higher the bond energy. An exception is fluorine (F 2), where the bond
length is so small that there is significant repulsion between the lone electron pairs
on one atom and the lone electron pairs on the other, increasing the length of the
bond, and so making it weaker than chlorine (Cl 2), for example.
Shape of molecules
The shape of a simple molecule or ion is determined by the repulsion between the
electron pairs (both bonding and lone) that surround the central atom.
Total Number Number Shape of molecule Bond Example
number of of of lone angle
electron bonding pairs (o)
pairs regions
2 2 0 Linear 180 MgCl2
CO2
3 3 0 Trigonal planar 120 BCl3
SO3
CO32-
NO32-
..
2 1 Bent 119 SO2
NO2-
4 4 0 Tetrahedral 109.5 CH4
3 1 Trigonal pyramid 107.5 NH3
SO32-
H3O+
2 2 Bent 104.5 H2O
.. ..
