m2,3 - ionisation energy + periodicity

Question 1

first ionisation energy def

Answer 1

the amount of energy required to remove
one electron from
each atom
of one mole of gaseous atoms
(to give one mole of gaseous 1+ ions)

Question 2

800, 2400, 3700, 25000
explain why 2nd IE is greater than 1st IE

Answer 2

electrons experience more attraction to the +ve nucleus, as each electron is removed there is less repulsion and each shell is drawn closer to the nucleus, so the electrons require more energy to release them.

Question 3

800, 2400, 3700, 25000
which group is this element likely to be in?

Answer 3

group 3, the large difference in IE between the 3rd and 4th electron shows the 4th electron was taken from a shell closer to the nucleus.

Question 4

what effects ionisation energy

Answer 4

nuclear charge - greater the nuclear charge the stronger the attractive force experienced by outer electrons
atomic radius - greater the atomic radius the weaker the nuclear attraction experienced by outer electrons
electron shielding - the repulsion between e- in different inner shells. inner shell e- repel outer shell e-. the more inner shells the stronger the shielding effect and so the less the attractive force experienced by the outer electron

Question 5

trend of first ionisation energy for first 20 elements…
noble gases
across periods
Be-B and Mg-Al
N-O and P-S

Answer 5

• noble gases have the highest IE in their period, because atoms have a full outer shell so a higher positive attraction from the nucleus
• increase across periods because: nuclear charge inc, so higher attraction on electrons. e- are added to the same shell so outer shell is drawn in slightly
• small decrease from Be-B and Mg-Al, because they have a p orbital (has higher energy than s orbitals) so are marginally further from the nucleus (so e- are easier to remove so have lower IE)
• small decrease from N-O and P-S, because outermost e- is now spin paired. these e- experience some repulsion which makes first outer electron easier to remove

Question 6

ionisation energy down the groups:
atomic radius
nuclear charge
electron shielding

Answer 6

atomic radius increases (more shells)
nuclear charge increases
electron shielding increases

ionisation energy decreases (e- further from nucleus, more shielding, e- held less strongly)

Question 7

ionisation energy across the periods:
atomic radius
nuclear charge
electron shielding

Answer 7

atomic radius decreases (increased nuclear charge attracts e- (and pulls them closer))
nuclear charge increases
electron shielding stays the same (no new shells)

ionisation energy increases (increased nuclear charge attracts e-, become harder to remove)

Question 8

first ionisation energy equation for eg Al

Answer 8

Al (g) -> Al+ (g) + e-

Question 9

periodicity def

Answer 9

a repeating trend in chemical or physical properties along a period