Factors Influencing Ionisation Energy
- Nuclear charge - the more protons in the nucleus of the atom, the greater the nuclear charge.
- Distance of the electrons from the nucleus - the more shells there are, the greater the distance between the outermost electrons and the nucleaus so the easier the electrons are to remove.
- Sheilding effect - any inner shells between the electrons and the nucleus repel outer electrons. The more repulsion there is, the less attraction and the easier the electrons are to remove.
Trends In Ionisation Energy
Ionisation energy increases across a period:
- Proton number increases, therefore the attractions for the outer electrons increases.
- Electrons enter the same shell so there is no increase in shielding.
- Outer electrons are held in more tightly (smaller radius).
- Nuclear charge is the dominant factor.
- Therefore electrons are harder to remove so ionisation energy increases.
Ionisation energy decreases down a group:
- Proton number increases.
- Electrons fill new shells and so shielding increases.
- Outer electrons are futher from the nucleus (larger radius).
- The increased nuclear charge is outweighed by the change in shielding and atomic radius (both increased) which are the dominant factors.
- Therefore it is easier to remove and electron so ionisation energy decreases.
Sub-Shell Trends In First Ionisation Energy
Explaining Sub-Shell Trends In First Ionisation En
Explain why boron has a lower first ionisation energy than beryllium
The first electron in Be is removed from the 2s subshell where as in B it's removed from the 2p. 2p is higher in energy than 2s so requires less energy to remove the first electron.
Explain why oxygen has a lower first ionisation energy than nitrogen
Highest energy electrons in the 2p subshell. Oxygen has a pair of electrons in one orbital which repel so it's easier to remove an electron.
Successive Ionisation Energies
Each time an electron is removed the ionisation energy increases (requires more energy to remove a second electron than the first one):
The electron is being removed from a charged partical (increasing the proton to electron ratio) so the same number of protons are attracting fewer electrons, thus meaning a higher ionisation energy is needed.
There is a large difference in value between some ionisation energies:
Some electrons are removed from a different shell closer to the nucleus with less shielding, so require a higher ionisation energy.
The Periodic Table
- Periods (rows) and groups (columns)
- There are eight groups
- There are four periods of transitions metal elements that slot into the table between groups two and three
- The elements are arranged in order of atomic number
Names for the groups:
- Group one = the alkali metals
- Group two = the alkaline earth metals
- Group seven = the halogens
- Group eight = the noble gases
As the atomic radius increases so does the ionisation energy.
Periodic Trends In Bonding And Structure
- Fixed metal cations surrounded by mobile, delocalised electrons
- These electrons mean metals can conduct as a solid and as a liquid
- The strong metallic bonds mean that high temperatures are needed to provide the large amount of energy to overcome the electrostatic attractions between cations and electrons
- Metals have high boiling and melting points
Giant Covalent Structures
- Giant covalent lattices are held together by a network of strong covalent bonds (eg. diamond and silicon)
- The strong covalent bonds need a large amount of energy to overcome them: they have high boiling and melting points
- They are also insoluble in most solvents - the covalent bonds are too strong to be broken by interactions with solvent molecules
- Most do not conduct electricity - all outer shell electrons are involved in covalent bonding. However, graphite and graphene conduct as they have mobile, delocalised electrons
Melting Points And Boiling Points
- The melting and boiling points of a particular element is dependent on the bonding present in the structure of the element
- Metals have high boiling points because they contain metallic bonding
- Non metals with giant covalent structures like silicon have high melting points because the covalent bonds are strong and difficult to break
- Simple covalent molecules like phosphorus and chlorine have low melting and boiling points because they are held together by intermolecular forces which are weaker and easily broken down
- Metallic elements are good conductors of electricity
- Their structures contain delocalised electrons which are free to move through the structure
- Metallic character and therefore conductivity increases to a maximum in group three
- Metallic character also increases down the group
- Elements beyond silicon show no conduction as all the electrons are involved in bonding and no free charge carriers are present