Atomic Structure Flashcards
(15 cards)
angle of deflection
charge/mass
why e- in Cr occupy 3d orbital first?
more stable as inter-electronic repulsion in 4s orbital minimised
all other elements e- occupy 4s orbitals first
why e- in Cu occupy 3d orbitals first?
more stable as 3d subshell (not orbitals) are fully filled with symmetrical charge distribution
all other elements e- occupy 4s orbitals first
why atomic radius decrease across period?
- no. of electronic shell remain same
- proton no. up, nuclear charge up
- e- up but shielding eff same as e- added to same outermost shell
- eff nuclear charge up
- electrostatic attraction between valence e- and nucleus up
- size of electron cloud down
why atomic radius increase down the period?
- no. of electronic shells up
- distance between valence e- and nucleus up
- shielding experience by valence e- up
- even tho nuclear charge up, electrostatic attraction between valence e- and nucleus down
- size of electron cloud up
why radius of cation smaller than anion?
- both got same nuclear charge
- cation got one less electron shell
- electrostatic attraction between valence e- and nucleus up
- size of electron cloud down
cation (positive charge) as electrons are removed
why ionisation energy increase across the period?
- electronic shell same
- no. of proton up, nuclear charge up
- e- up but shielding eff same as e- added to same outermost shell
- eff nuclear charge up
- electrostatic attraction between valence e- and nucleus up
- energy required to remove valence e- up
why ionisation energy decrease down the group?
- no. of electronic shells up
- distance between valence e- and nucleus up
- shielding experience by valence e- up
- even tho nuclear charge up, electrostatic attraction between valence e- and nucleus down
- energy required to remove valence e- down
why does ionisation energy decrease from group 2 to group 13?
- 3p electron removed from Al is at a higher energy level than 3s e- removed from Mg
- lesser energy required to remove 3p electron than 3s electron
- first IE of Al is lower than Mg
Al (group 13) Mg (group 2)
why does ionisation energy decrease from group 15 to group 16?
- 3p electron removed from S is a paired electron
- 3p electron removed from P is unpaired electron
- inter-electronic repulsion present between paired e- in the same orbital
- lesser energy required to remove paired 3p electron from S
S (group 16) P (group 15)
why successive ionisation energy of element increase?
- first e- removed is from neutral atom
- successive e- removed is from ion of **increasing positive charge **that attracts e- more strongly
simi is electronegativity?
the ability to attract bonding electrons
why electronegativity increase across period?
- electronic shell same
- no. of proton up, nuclear charge up
- e- up but shielding eff same as e- added to same outermost shell
- eff nuclear charge up
- electrostatic attraction between bonding e- and valence e- up
- electronegativity increase across period
why electronegativity decrease down the group?
- no. of electronic shells up
- distance between valence e- and nucleus up
- shielding experience by valence e- up
- even tho nuclear charge up, electrostatic attraction between bonding e- and nucleus down
- electronegativity decrease down the group
why is LiAlH4 a stronger reducing agent than NaBH4?
There is a larger difference in electronegativity in Al–H than in B–H, which makes the H in Al–H more electron rich than in B–H. This allows LiAlH 4 to be a stronger reducing agent.
