1.2 Electron Orbitals and Configurations Flashcards
principle quantum number
- symbol: n
- defines: size and energy of the shell
- n = 1, 2, 3, 4
- symbol: n
- defines: size and energy of the shell
- n = 1, 2, 3, 4
principle quantum number
angular quantum number
- symbol: l
- defines: shape of sub shell
- n = 0 to n-1l = 0 = s orbitall = 1 = p orbitall = 2 = d orbitall = 3 = f orbital
- symbol: l
- defines: shape of sub shell
- n = 0 to n-1
angular momentum quantum number
magnetic quantum number
- symbol: ml
- defines: direction
- n = -l .. 0 .. +l
- ml will tell you how many electron pairs are in the orbital
0 = 1 pair
-1, 0, +1 = 3 pairs
- symbol: ml
- defines: direction
- n = -l .. 0 .. +l
magnetic quantum number
spin magnetic quantum number
- symbol: ms
- defines: spin
- n = -1/2 or +1/2
- symbol: ms
- defines: spin
- n = -1/2 or +1/2
spin magnetic quantum number
draw all the shapes of orbitals in order of s, p, d, f
aufbau rule
Electrons occupy orbitals in order of increasing energy, eg the lowest energy orbital is filled first
Electrons occupy orbitals in order of increasing energy, eg the lowest energy orbital is filled first
aufbau rule
hunds rule
degenerate (equal energy) orbitals are occupied singularly before pairing
degenerate (equal energy) orbitals are occupied singularly before pairing
hunds rule
degenerate
of equal energy
of equal energy
degenerate
No 2 electrons can have the same 4 quantum numbers numbers and if there are 2 electrons, their spin must be opposite
pauli exclusion principle
pauli exclusion principle
No 2 electrons can have the same 4 quantum numbers numbers and if there are 2 electrons, their spin must be opposite
atomic orbital
a region of high probability of finding an electron
a region of high probability of finding an electron
atomic orbital
how many electrons does an orbital hold
maximum of two
how many electrons do s, p, and d orbitals hold each
- s orbital will hold 2 electrons
- p orbital will hold 6 electrons
- d orbital will hold 10 electrons
electron configuration of sodium
1s2 2s2 2p6 3s1
electron configuration of calcium
1s2 2s2 2p6 3s2 3p6 4s2
electron configuration of Sr2+
1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6
electron configuration of Ca2+
1s2 2s2 2p6 3s2 3p6
electron configuration of Cs2+
1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p5
Cr2+
1s2 2s2 2p6 3s2 3p6 4s2 3d2
what is unusual about chromium
when electrons occupy chromium, instead of being 4s2 3d4, it will be 4s1 3d5
an electron will move from the 4s orbital to the 3d orbital because there is a special stability associated with half filled orbitals.
atoms would rather have 2 half filled than 1 filled.
what is unusual about copper
copper will be 4s1 3d10 rather than 4s2 3d9 due to the stability associated with filled orbitals
what electrons are lost when atoms become ions
electrons from the 4s orbital
ionisation energy
the energy required to remove one mole of electrons from one mole of atoms in their gaseous state
the energy required to remove one mole of electrons from one mole of atoms in their gaseous state
ionisation energy
ionisation energy and half filled/filled shells
there is a stability associated with filled and half filled orbitals. this means that it requires more energy to remove electrons
nuclear charge and ionisation energy
an increase in protons across periods means there is a stronger attraction for electrons, so requires more energy to remove
atomic size and ionisation energy
as the atomic size increases down groups, the further away electrons are less attracted to the nucleus so require less energy to remove electrons
sheilding and ionisation energy
inner electrons shield outer electrons from the nuclear charge of the nucleus, making them easier to remove. increases down groups
electron configuration of aluminium (13 electrons)
1s2 2s2 2p6 3s2 3p1
