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ELECTRON ARRANGEMENT
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• Electrons are arranged in electrons shells (energy levels).
• The shells have sub-shells (sub-levels).
• Each shell/sub-shell is made up of electron orbitals which can each hold 2 electrons.
Shells, sub-shells & orbitals
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• Each sub-level consists of electron orbitals (region of space in which the electron spends most of its time).
• Each orbital can hold 2 electrons with opposite spins (one electron spins clockwise and one anticlockwise).
• Orbitals are regions of space that electrons are most likely to be in.
Orbitals
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s orbital p orbital
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The Orbitron
http://winter.group.shef.ac.uk/orbitron/AOs/1s/index.html
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4s of "lower" energy than 3d
Distance from nucleus
Energy
1s
2s
2p
3s
3p
3d 4s
4p
4d
4f
Ionisation energy
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Orbitals
Sub-level
Number of orbitals in sub-level
Shape (no need to learn)
Maximum number of
electrons in sub-level
s 1 2
p 3 6
d 5 10
f 7 Even more complicated! 14
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other T-shirts are available!!
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Electrons enter the lowest energy orbital available.
Aufbau Principle
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Electrons prefer to occupy orbitals on their own, and only pair up when no empty orbitals of the same energy are available .
Hund’s Rule
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4s of "lower" energy than 3d
Distance from nucleus
Energy
1s
2s
2p
3s
3p
3d 4s
4p
4d
4f
Ionisation energy
e.g. silicon 14 e- 1s2 2s2 2p6 3s2 3p2
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e.g. calcium 20 e- 1s2 2s2 2p6 3s2 3p6 4s2
4s of "lower" energy than 3d
Distance from nucleus
Energy
1s
2s
2p
3s
3p
3d 4s
4p
4d
4f
Ionisation energy
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• The highest energy electrons are lost when an ion is formed.
• Note that 4s electrons are lost before 3d (as once 4s and 3d are occupied, 4s moves above 3d).
Ions
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e.g. Ca2+ 18 e- 1s2 2s2 2p6 3s2 3p6
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4s of "lower" energy than 3d
Distance from nucleus
Energy
1s
2s
2p
3s
3p
3d 4s
4p
4d
4f
Ionisation energy
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• Cu and Cr do not have the expected electron structure.
Cr = 1s2 2s2 2p6 3s2 3p6 4s1 3d5 NOT 1s2 2s2 2p6 3s2 3p6 4s2 3d4
Cu = 1s2 2s2 2p6 3s2 3p6 4s1 3d10 NOT 1s2 2s2 2p6 3s2 3p6 4s2 3d9
Cu & Cr
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• Evidence for how the electrons are arranged in atoms comes from ionisation energies.
• 1st ionisation energy = energy required to remove one electron from each atom in a mole of gaseous atoms producing one mole of 1+ gaseous ions.
• Note that 2nd ionisation energy is the energy required to remove the second electron (not both electrons).
e.g. 1st IE of Na: Na(g) → Na+(g) + e–
2nd IE of Na: Na+ (g) → Na2+(g) + e
Ionisation Energy
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Ionisation Energy
+
e_
2) nuclear charge
1) distance from nucleus
3) shielding (repulsion) by electrons in inner shells between nucleus and outer electron
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1st ionisation energy (down group)
0
200
400
600
800
1000
Be Mg Ca Sr Ba
1st
ion
isat
ion
en
erg
y
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1st ionisation energy (down group)
• Atoms get bigger
• More shielding
• Therefore weaker attraction from nucleus to electron in outer shell
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1st ionisation energy (across period)
0200
400600
8001000
12001400
1600
Na Mg Al Si P S Cl Ar1st
ion
isat
ion
en
erg
y (k
J/m
ol)
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1st ionisation energy (across period)
General trend
• Increased nuclear charge (i.e. more protons)
• Atoms get smaller
• Therefore stronger attraction from nucleus to electron in outer shell
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1st ionisation energy (across period)
Group 2 → 3
• Electron lost from Group 3 element is from p orbital, while that lost from Group 2 element is from s orbital.
• p orbital is higher energy than s orbital, so easier to lose electron.
4s of "lower" energy than 3d
Ionisation energy
Distance from nucleus
Energy
4f 4d 4p 4s
3d 3p 3s
2p
2s
1s
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1st ionisation energy (across period)
Group 5 → 6
• Group 6 element loses electron from orbital with 2 electrons (p4)
• Group 5 element loses electron from orbital with 1 electrons (p3)
• Extra electron-electron repulsions make it easier to lose electron from p4 than p3.
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1st ionisation energy
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1st ionisation energy
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down a group(group 0)
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1st ionisation energy
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down a group(group 1)
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1st ionisation energy
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period 2period 3period 4
Across a period
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1st ionisation energy
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End of period
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Successive ionisation energies (K)
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