There are many important concepts and trends that need to be understood in Inorganic Chemistry. The understanding of these trends can lead to a much better understanding of properties on the whole.
Perodicity
The periodicity are the properties of the elements across the periodic table. On of the biggest changes is the fact that there is the movement from metal to non-metal properties across the period. This is due to a number of different reasons:
Valence electrons and electronic structure
Nearly all bonding occurs between valence electrons, these are the electrons highest in energy and are the first to be removed or shared in a reaction. Whether or not it is more favourable for an element to gain or lose electrons depends on the eletronegativities of the elements that it is in contact with.
Size of atom: Atomic radius
The effective nuclear charge is a measure of the overall positive charge that an atom has and the electrons around it that can cause shielding affects to the valence electrons. Taking argon as an example, this has a small number of shielding electrons but the largest nucleus in the period. This causes it to have a very large effective nuclear charge. This means that the valence shell is pulled inwards.
For ions a cation is pulled inwards a great deal as the effective nuclear charge has risen whereas an anion is bigger than the orginal ion as the effective nuclear charge has decreased slightly. This means that there is a slight rise in energy, although this is much smaller than the energy decrease in having a stable octet of electrons.
Electronegativity
The electronegativity is not an exact quantity, it is the power of an atom to attract electrons and takes into account many different properties of the atom such as ionic radius, effective nuclear charge and electron affinity, to give a measurement. The electronegativity increases towards the top rights of the peiodic table and decrease towards to bottom left.
- Valence electrons and electronic structure.
- Size of atom: atomic radius.
- Electronegativity.
- Electron affinity.
- Polarisability
Valence electrons and electronic structure
Nearly all bonding occurs between valence electrons, these are the electrons highest in energy and are the first to be removed or shared in a reaction. Whether or not it is more favourable for an element to gain or lose electrons depends on the eletronegativities of the elements that it is in contact with.
Size of atom: Atomic radius
The effective nuclear charge is a measure of the overall positive charge that an atom has and the electrons around it that can cause shielding affects to the valence electrons. Taking argon as an example, this has a small number of shielding electrons but the largest nucleus in the period. This causes it to have a very large effective nuclear charge. This means that the valence shell is pulled inwards.
For ions a cation is pulled inwards a great deal as the effective nuclear charge has risen whereas an anion is bigger than the orginal ion as the effective nuclear charge has decreased slightly. This means that there is a slight rise in energy, although this is much smaller than the energy decrease in having a stable octet of electrons.
Electronegativity
The electronegativity is not an exact quantity, it is the power of an atom to attract electrons and takes into account many different properties of the atom such as ionic radius, effective nuclear charge and electron affinity, to give a measurement. The electronegativity increases towards the top rights of the peiodic table and decrease towards to bottom left.
Electron affinity
Electron affinity is a direct measurement. This is the energy change when an electron is added to an atom, there can be multiple electron affinity readings for each atom although they are usually slightly higher as there is an addition of an electron to an already negatively charged species.
Electron affinity is a direct measurement. This is the energy change when an electron is added to an atom, there can be multiple electron affinity readings for each atom although they are usually slightly higher as there is an addition of an electron to an already negatively charged species.
Hard and Soft properties
The idea of hard and soft cations, anions, acids and bases is a concept that measures the interactions between different species. It is found that when all of the conditions are kept the same that soft attracts soft and hard attracts hard. This is due to similarities in energy levels between orbitals allowing stronger more thermodynamically favourable complexes to form.
Hard acids and hard bases
Examples of hard acids are: H+, Li+, Na+, K+, Ti4+, Cr6+. These are all highly charged cationic species, this means that the valence shell on these species is drawn very close to the centre and the effective nuclear charge is very high. Another example is BF3, this has large electron withdrawing effects away from the boron centre causing a very high electronegativity making it a powerful Lewis acid. Examples of hard bases are: OH-, F-, Cl-, NH3-, H2O, CO32-. These have highly electronegative groups which have a high affinity other hard acids.
Soft acids and soft bases
Soft acids and soft bases have very different peoperties to their hard counterparts.
Examples of these soft acids are: Pt2+, Ag+, Au+ Hg2+ BH3, notice these are high oxidation state cations or weak Lewis acids, the borane in this example has a high electron density centred on the boron group which causes only a small electron accepting ability. Some soft bases are: H-, R3P, SCN-, I-. The reactions here are much more covalent between soft and soft compared to the more ionic hard and hard.
Hard acids and hard bases
- Small atomic/ionic radius.
- High oxidation state.
- Low polarisability.
- High electonegativity.
Examples of hard acids are: H+, Li+, Na+, K+, Ti4+, Cr6+. These are all highly charged cationic species, this means that the valence shell on these species is drawn very close to the centre and the effective nuclear charge is very high. Another example is BF3, this has large electron withdrawing effects away from the boron centre causing a very high electronegativity making it a powerful Lewis acid. Examples of hard bases are: OH-, F-, Cl-, NH3-, H2O, CO32-. These have highly electronegative groups which have a high affinity other hard acids.
Soft acids and soft bases
Soft acids and soft bases have very different peoperties to their hard counterparts.
- Large atomic/ionic radius.
- Low or zero oxidation state bonding.
- High polarisability.
- Low electronegativity.
Examples of these soft acids are: Pt2+, Ag+, Au+ Hg2+ BH3, notice these are high oxidation state cations or weak Lewis acids, the borane in this example has a high electron density centred on the boron group which causes only a small electron accepting ability. Some soft bases are: H-, R3P, SCN-, I-. The reactions here are much more covalent between soft and soft compared to the more ionic hard and hard.
Inert pair effect
The inert pair effect is an important effect that is seen in the p-block elements. This is an important concept as it explains the differences seen in the oxidation states of the elements further down the group.
This effect is seen in group 14 for example where that these elements are found in a +4 oxidation state apart from lead which is mostly found in a +2 oxidation state. This also occurs in group 13 where indinum and thallium are usually found in a +1 oxidation state where boron, alumnium and gallium are found in a +3 oxidation state.
So why does this occur?
The concept has nothing to do with ionisation enthalpies the ionisation enthalpies of these later elements are actually very similar to other elements that are found in higher oxidation state. What does happen though is that these larger elements have more diffuse orbitals this causes a decrease in bond enthalpy. This decrease in bond enthalpy means that forming more bonds in unfavoured, so doesn't occur as there is no thermodynamic driving force.
This effect is seen in group 14 for example where that these elements are found in a +4 oxidation state apart from lead which is mostly found in a +2 oxidation state. This also occurs in group 13 where indinum and thallium are usually found in a +1 oxidation state where boron, alumnium and gallium are found in a +3 oxidation state.
So why does this occur?
The concept has nothing to do with ionisation enthalpies the ionisation enthalpies of these later elements are actually very similar to other elements that are found in higher oxidation state. What does happen though is that these larger elements have more diffuse orbitals this causes a decrease in bond enthalpy. This decrease in bond enthalpy means that forming more bonds in unfavoured, so doesn't occur as there is no thermodynamic driving force.