Periodicity Module 3 Flashcards
What are periods in the periodic table?
Horizontal row.
All the elements in the same period have similar electron core (which is the same as electron configuration of group 0 element).
Periodicity - elements in period 2 closely resemble elements in period 3 in terms of repeating chemical behaviour and physical properties.
When all end in same electron configuration ( e.g. p2) they will react in same way.
What are the diff groups called?
1- alkali metals
2- alkaline earth metals
7- halogens
8- noble gases
What is the first ionisation energy trend?
The first ionisation of an element is the energy required for the removal of 1 mole of electrons from 1 mole of gaseous atoms. (They are endothermic as energy needs to be put in to overcome electrostatic forces.)
E.g. Na (g) —> Na+ (g) + e-
O (g) —> O+ (g) +e-
( remember, its always a gas even if its normally not and don’t include 2 in diatomic ions.
What factors influence ionisation energy?
Nuclear charge
Electron shielding ( number shells)
Atomic radius ( distance between outer electron and nucleus )
How does nuclear charge influence ionisation energy?
More protons, the more nuclear charge. The greater nuclear charge, the greater attraction between outer electron so more energy required so has more ionisation energy.
How does distance from nucleus affect first ionisation energy?
As outer electron becomes further from nucleus , (atomic radius increases) attraction between outer electron and nucleus weaker so less energy required to remove as its easier to remove.
(Smaller atomic radius, stronger attraction so more energy so higher ionisation energy)
How does electron shielding affect first ionisation energy?
The more electron shielding (more electron shells) the lower the ionisation energy
What is the general first ionisation energy trend across a period?
First ionisation energy increase as there are more protons (so more nuclear charge).
Shielding remains constant.
So outer electron experiences more nuclear attraction so more energy required.
Why is there a small decrease in first ionisation energy between groups 2 + 3 across a period?
Electron in group 3 being removed from P sub shell which is slightly higher energy than S sub shell which group 2 removed from. So group 3 outer electron experiences less nuclear attraction so less energy required as its easier to remove.
Why is there a small decrease in first ionisation energy between groups 5+6 across a period?
Electron being removed from group 6 is a full orbital. Repulsion between 2 electrons in same orbital means electron easier to remove so has a lower nuclear attraction so lower first ionisation energy.
What is the general trend in first ionisation energy down a group?
First ionisation decrease because as go down, electrons fill new shell so electron shielding increases. So nuclear attraction decreases as electron further from nucleus and shielded more so less energy required lowering first ionisation energy. Although nuclear charge increases, increase in distance and electron shielding far outweighs the nuclear charge.
What is succession ionisation energies?
The energy required to remove each electron in turn.
What happens to size of atomic radius as move across period?
Atomic radius Decreases because:
- nuclear charge increases
- shielding remains constant
- nuclear attraction increases so more energy required to remove electron
What happens to atomic radius down a group?
Atomic radius Increase:
- more shells so more shielding
- outer electron further from nucleus so less nuclear attraction as go down group.
Explain why groups 1-3 have a high mp +bp?
Metallic bonding
Strong forces of attraction between positive ions and delocalised electrons which require high temperature to break the metallic bond.
More positively charged ion is, the stronger the force of attraction so higher mp.
