FIRST IONISATION ENERGY
Energy required to remove 1 mole of electrons from 1 mole of gaseous atoms tp from 1
mole of gaseous +1 ions
X(g) →X+(g) + e-
Decreases down the group:
● Atomic radius increases
● Shielding by inner shells increases
● Nuclear attraction on outer electrons decreases
● Outer electron is held more loosely to the nucleus so it’s easier to remove it
Increases across the period:
● Nuclear charge increases
● Similar shielding (same number of electron shells)
● Nuclear attraction increases
● Atomic radius decreases
● Outer electron is held more tightly to the nucleus so it’s harder to remove it
Decreases from Be to B: 2p sub shell has a higher energy than the 2s sub shell in be so
the 2p electron is easier to remove in B
Decreases from N to O: paired electrons in 2p sub shell repel each other, making it
easier to remove an electron from O than N
PERIODIC TREND IN MELTING POINTS
Increases from group 1 to group 4 :
Groups 1 -3 have metallic bonding increasing in strength due to increased forces of
attraction between more electrons in the outer shell taht are released to the sea of
delocalised electrons and the higher metal ion charge = more energy to break bonds
and melt the solid metallic lattice
Group 4 has a giant covalent structure with many strong covalent bonds requiring a lot
of energy to overcome
Decreases from group 5 to 0:
They have simple molecular structures with weak London forces between molecules,
requiring little energy to break
S8 has a higher melting point: more electrons = stronger London forces
ALKALI EARTH METALS (GROUP 2)
Reactivity increases down the group (first ionisation energy decreases):
- As atomic radius increases, more shielding
- Nuclear attraction decreases = easier to remove outer electrons
Down group 2, sulfates become less soluble
Solubility of group 2 hydroxides increases down the group:
Magnesium hydroxide is almost insoluble
Slow reaction
Energy required to remove 1 mole of electrons from 1 mole of gaseous atoms tp from 1
mole of gaseous +1 ions
X(g) →X+(g) + e-
Decreases down the group:
● Atomic radius increases
● Shielding by inner shells increases
● Nuclear attraction on outer electrons decreases
● Outer electron is held more loosely to the nucleus so it’s easier to remove it
Increases across the period:
● Nuclear charge increases
● Similar shielding (same number of electron shells)
● Nuclear attraction increases
● Atomic radius decreases
● Outer electron is held more tightly to the nucleus so it’s harder to remove it
Decreases from Be to B: 2p sub shell has a higher energy than the 2s sub shell in be so
the 2p electron is easier to remove in B
Decreases from N to O: paired electrons in 2p sub shell repel each other, making it
easier to remove an electron from O than N
PERIODIC TREND IN MELTING POINTS
Increases from group 1 to group 4 :
Groups 1 -3 have metallic bonding increasing in strength due to increased forces of
attraction between more electrons in the outer shell taht are released to the sea of
delocalised electrons and the higher metal ion charge = more energy to break bonds
and melt the solid metallic lattice
Group 4 has a giant covalent structure with many strong covalent bonds requiring a lot
of energy to overcome
Decreases from group 5 to 0:
They have simple molecular structures with weak London forces between molecules,
requiring little energy to break
S8 has a higher melting point: more electrons = stronger London forces
ALKALI EARTH METALS (GROUP 2)
Reactivity increases down the group (first ionisation energy decreases):
- As atomic radius increases, more shielding
- Nuclear attraction decreases = easier to remove outer electrons
Down group 2, sulfates become less soluble
Solubility of group 2 hydroxides increases down the group:
Magnesium hydroxide is almost insoluble
Slow reaction
