Aufbau Principle

1
Aufbau Principle
e- fill lowest energy levels 1st.
Half-filled and filled states are preferred.
1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 5g 6s
6p 6d 6f 6g 6h 7s 7p 7d 7f 7g 7h 7i
4f
5p
4d
5s
4p
3d
Increasing energy
4s
3p
3s
1 H Hydrogen 1s1
2 He Hydrogen 1s2
3 Li Lithium 1s22s1
4 Be Beryllium 1s22s2
5 B Boron 1s22s22p1
6 C Carbon 1s22s22p2
7 N Nitrogen 1s22s22p3
8 O Oxygen 1s22s22p4
9 F Fluorine 1s22s22p5
10 Ne Neon 1s22s22p6
11 Na Sodium Ne3s1
12 Mg Magnesium Ne3s2
13 Al Aluminum Ne3s23p1
14 Si Silicon Ne3s23p2
15 P Phosphorus Ne3s23p3
16 S Sulfur Ne3s23p4
17 Cl Chlorine Ne3s23p5
18 Ar Argon Ne3s23p6
19 K Potassium Ar4s1
20 Ca Calcium Ar4s2
21 Sc Scandium Ar3d14s2
22 Ti Titanium Ar3d24s2
23 V Vanadium Ar3d34s2
24 Cr Chromium Ar3d54s1
24 Cr Chromium Ar3d54s1
25 Mn Manganese Ar3d54s2
26 Fe Iron Ar3d64s2
27 Co Cobalt Ar3d74s2
28 Ni Nickel Ar3d84s2
29 Cu Copper Ar3d104s1
29 Cu Copper Ar3d104s1
30 Zn Zinc Ar3d104s2
31 Ga Gallium Ar3d104s24p1
32 Ge Germanium Ar3d104s24p2
33 As Arsenic Ar3d104s24p3
34 Se Selenium Ar3d104s24p4
35 Br Bromine Ar3d104s24p5
36 Kr Krypton Ar3d104s24p6
2p
2s
1s
nucleus
2
Irregular e- configurations of Cr and Cu

• Half filled and filled sublevels are preferred.

Chromium A 4s electron moves to a 3d sublevel to
half fill its entire 3d sublevel
Copper A 4s electron moves to FILL its 3d
sublevel
3
Orbital filling table
s1
s2p1
s2p2
s2p3
s2p4
s2p5
s2
s2
s2p6
4
Mendeleevs Periodic Table (1872)
Dmitri Mendeleev 1st to publish an organized
table of elements.

• Grouped elements by similar chemical properties.
• Arranged elements by increasing mass.

5
Modern Russian Table
6
Henry Moseley (1913)
Used X-Ray diffraction to determine how many
protons are in an atom of an element.

• Grouped elements by similar chemical properties.
• Arranged elements by increasing atomic number.

7
The Periodic Table
Period
Group or Family
1
2
3
4
5
6
7
8
Properties of Metalloids
Metalloids straddle the border between metals and
nonmetals on the periodic table.
Metalloids
9
Periodic Table with Group Names
Period
Group or Family
1
2
3
Noble Gases
Halogens
Transition Metals
4
Alkali Metals
Alkaline Earth Metals
Boron Family
Carbon Family
Nitrogen Family
Oxygen Family
5
6
7
Inner Transition Metals
10
Trends in Atomic Size

• First problem Where do you start measuring from?
• The electron cloud doesnt have a definite edge.
• They get around this by measuring more than 1
  atom at a time.

11
Determination of Atomic Radius
Half of the distance between nuclei in
covalently bonded diatomic molecule
(e- shared)
"covalent atomic radii"
12
Trends in Atomic Size

• Influenced by three factors
• 1. Energy Level
• Higher energy level is further away.
• 2. Charge on nucleus
• (Zeff effective nuclear charge)
• More charge pulls electrons in closer.
• 3. Shielding effect
• Electrons within level and from previous levels
  block the effects of the () nucleus.

e lt-gt e repulsion
13
Periodic TableTrend for Atomic Radii
Size of atom decreases from left to right.
Size of atom increases from top to bottom.
14
Periodic Trends in Atomic Radius

• Radius decreases across a period

Effective nuclear charge, Zeff, is increased.
Due to ? more overall charge (p and e-) ?
shielding from lower levels is constant.
15
Group trends
H
Li

• As we go down a group...
• each atom has another energy level (more
  shielding)
• so the atoms get bigger (Zeff is less effective).

Na
K
Rb
16
Periodic Trends

• As you go across a period, the radius gets
  smaller. (more effective nuclear charge, Zeff)
• Electrons are in same energy level.
• More nuclear charge. atomic (p e-)
• Outermost electrons are closer.

z 11
17
18
12
13
14
15
16
Na
Mg
Al
Si
P
S
Cl
Ar
17
Overall
Rb
K
Na
Li
Atomic Radius (nm)
Kr
Ar
Ne
H
Atomic Number
10
18
1st Ionization Energy - the energy required to
remove the 1st electron from a mole of atoms.
2372 kJ/mol
He
H
1312 kJ/mol
2081 kJ/mol
Li
Ne
Be
520 kJ/mol
N
O
F
B
C
1521 kJ/mol
Ar
Na
496 kJ/mol
1351 kJ/mol
K
Kr
419 kJ/mol
Increasing 1st Ionization energy
1170 kJ/mol
Rb
403 kJ/mol
Xe
1037 kJ/mol
Cs
Rn
376 kJ/mol
Fr
375 kJ/mol
19
Ionization Energy - the energy required to
remove an electron from an atom

• Increases for successive electrons taken from
• the same atom

• Tends to increase across a period

Electrons in the same quantum level do not
shield as effectively as electrons in inner
levels
    Irregularities at half filled and filled
sublevels due to extra repulsion of
electrons paired in orbitals, making them
easier to remove

• Tends to decrease down a group

Outer electrons are farther from the nucleus
20
Table of 1st Ionization Energies
21
Ionization of Magnesium
Mg 738 kJ ? Mg e-
1st ionization energy
Mg 1451 kJ ? Mg2 e-
2nd ionization energy
Mg2 7733 kJ ? Mg3 e-
3rd ionization energy
22
Symbol First Second Third
5247 7297 1757 2430 2352 2857 3391 3375 3963
1312 2731 520 900 800 1086 1402 1314 1681 2080
HHeLiBeBCNO F Ne
11810 14840 3569 4619 4577
5301 6045 6276
23
Symbol First Second Third
11810 14840 3569 4619 4577
5301 6045 6276
5247 7297 1757 2430 2352 2857 3391 3375 3963
1312 2731 520 900 800 1086 1402 1314 1681 2080
HHeLiBeBCNO F Ne
24
Electronegativity
A measure of the ability of an atom in a
chemical compound to attract electrons

• Electronegativities tend to increase across
• a period

• Electronegativities tend to decrease down a
• group or remain the same

25
Periodic Table of Electronegativities
Increasing electronegativity
26
Summation of Periodic Trends
27
Ionic Radii

• Positively charged ions formed when
• an atom of a metal loses one or
• more electrons

Cations

• Smaller than the corresponding
• atom

• Negatively charged ions formed
• when nonmetallic atoms gain one
• or more electrons

Anions

• Larger than the corresponding
• atom

28
Ion radius
Atom radius
29
Shielding

• The electron in the outermost energy level
  experiences more inter-electron repulsion
  (shielding).
• Second electron has same shielding, if it is in
  the same period

30
Group trends

• As you go down a group, first IE decreases
  because...
• The electron is further away.
• More shielding.

31
Periodic trends

• All the atoms in the same period have the same
  energy level.
• Same shielding.
• But, increasing nuclear charge
• So IE generally increases from left to right.
• Exceptions at full and 1/2 full orbitals.

32
He

• He has a greater IE than H.
• same shielding
• greater nuclear charge

H
First Ionization energy
electrons
1s
2s
2p
nucleus
Orbital diagram
Atomic number
33
He

• Li has lower IE than H
• Outer electron further away
• outweighs greater nuclear charge

H
First Ionization energy
electrons
Li
nucleus
1s
2s
2p
Orbital diagram
Atomic number
34
He

• Be has higher IE than Li
• same shielding
• greater nuclear charge

H
First Ionization energy
Be
electrons
Li
nucleus
1s
2s
2p
Orbital diagram
Atomic number
35
He

• B has lower IE than Be
• same shielding
• greater nuclear charge
• p orbital is slightly more diffuse and its
  electron easier to remove

H
First Ionization energy
Be
electrons
B
Li
nucleus
1s
2s
2p
Atomic number
36
He
C
H
First Ionization energy
Be
electrons
B
Li
nucleus
1s
2s
2p
Atomic number
37
He
N
C
H
First Ionization energy
Be
B
Li
nucleus
1s
2s
2p
Atomic number
38
He

• Breaks the pattern, because the outer electron is
  paired in a p orbital and experiences
  inter-electron repulsion.

N
O
C
H
First Ionization energy
Be
B
Li
nucleus
1s
2s
2p
Atomic number
39
He
F
N
O
C
H
First Ionization energy
Be
B
Li
nucleus
1s
2s
2p
Atomic number
40
He
Ne

• Ne has a lower IE than He
• Both are full,
• Ne has more shielding
• Greater distance

F
N
O
C
H
First Ionization energy
Be
B
Li
nucleus
1s
2s
2p
Atomic number
41
He
Ne

• Na has a lower IE than Li
• Both are s1
• Na has more shielding
• Greater distance

F
N
O
C
H
First Ionization energy
Be
B
Li
Na
Atomic number
42
First Ionization energy
Atomic number
43
Table of 1st Ionization Energies