Periodicity Flashcards
(17 cards)
Define first ionisation energy
The energy required to remove 1 mole of electrons from 1 mole of gaseous atoms of an element to form 1 mole of gaseous 1+ ions
How can you identify when a shell has been removed from a successive ionisation energy graph?
A large jump in the ionisation energy value
What is meant by sheilding?
The effect of the inner electrons shielding the outer electrons from the effect of the charge on the nucleus
What is meant by nuclear charge?
The positive charge on the nucleus
Describe the trend in atomic radius across period 3.
- The shielding stays the same
- The nuclear charge increases
- These cause the nuclear attraction to increase so atomic radius decreases
What is the shape of the general trend graph for first ionisation energies across any period?
A straight line originating at the origin
Why does first ionisation energy increase across period 3?
- The nuclear charge increases
- Atomic radius decreases
- Nuclear attraction increases
- Shielding stays the same
- It takes more energy to remove the first electron so first ionisation energy increases
Describe the actual trend in first ionisation energy across period 3
- Increases slightly from sodium to magnesium
- Drops slightly from magnesium to aluminium
- Increases significantly between aluminium and sillicon
- Increases slightly between sillicon and phosphorus
- Decreases between phosphorus and sulphur
- Increases until the end of the period
Explain why aluminium has a lower first ionisation energy than magnesium.
- Aluminium has 1 electron in a higher subshell
- This electron is removed more easily as it is further away from the nucleus
- Therefore first ionisation energy is lower
What is the general trend in shielding, nuclear charge and atomic radius down a group?
All 3 increase
Why does nuclear attraction decrease down a group?
- Shielding and nuclear charge both increase, cancelling each other out
- The nuclear attraction decreases
Why does first ionisation energy decrease down a group?
- There is increased shielding and nuclear charge
- Nuclear attraction decreases
- Atomic radius increases
- There is less energy to remove the outer electron
Why does melting point increase from Na - Al in period 3?
(Metallic bonding)
- The charge on the metal ions increases from +1 to +3
- The number of delocalised electrons increases so the strength of the metallic bond increases
- More energy is needed to break stronge metallic bonds so the melting point increases
Why does Si have a much higher melting point than all other period 3 elements?
(Giant covelant)
- Si has a similar structure to the carbon atoms in diamond
- Each Si atom is bonded to 4 other atoms
- The covelant bonds are incredibly strong so a high amount of energy is needed to break them
Why do elements P - Ar in period 3 have a low melting point?
(Simple covelant)
- When the 4 elements melt or boil, the weak London forces break.
- Little energy is needed to overcome the forces so the melting point is low
What order does the melting points of period 3 elements decrease in and why?
- S > P > Cl > Ar
- There is more surface contact between sulphur molecules (S8) than phospurus molecules (P4) leading to stronger London forces and a higher melting point
Why does reactivity increase down groups 1 and 2?
- Sheilding and nuclear charge increase
- Nuclear attraction decreases
- Atomic radius increases so it is easier to remove electrons
