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Effective nuclear charge
Effective Nuclear charge
Effective Nuclear charge
H
• Modern periodic table: arrange elements in order of
increasing atomic number.
He
1 proton
•Effective nuclear charge (Zeff):
actual charge exerted by nucleus on e–.
•Main concept of Ch. 7
1 e–
In He atom, each e– attracted by 2 protons
•The Zeff ≠ charge on nucleus
due to shielding effect of inner e–.
Effective Nuclear charge
In H atom,
e–
attracted by 1 proton
Effective Nuclear Charge
What does cross section of e– distribution look like
(probability of finding e– for 3s orbital)?
Li
• Zeff is approx. since orbitals have differing shapes,
sizes and nodes.
Z=3
• What is full e– configuration for Mg?
1s22s22p63s2
In Li atom, valence e– attracted by 3 protons, but
repelled by e– sometimes.
Zeff ∼ Z - # core e–
Probability
(e– density)
What are the valence electrons?
which ones are closer to nucleus?
Distance from nucleus
Calculate Zeff for:
Z - S = Zeff
Mg
11-10 ∼ 1
Cl
K
Ca
Br
Rb
Sr
I
Na
Penetration of 2s and 3 s inside 1s makes Zeff approx
Zeff ∼ Z – S
Z = atomic number
S = core e–
1
Effective Nuclear Charge
summary
Calculate Zeff for:
Z - S = Zeff
11-10 = 1
Na
Mg
12-10 = 2
Cl
17-10
=7
K
19-18 = 1
Ca
20-18 = 2
Br
35-28
=7
Rb
37-36 = 1
Sr
38-36 = 2
I
53-46
=7
Sizes of Atoms and Ions
• Outer e– attracted to nucleus,
• core (inner) e– screen valence (outer) e– from the
nuclear charge.
• As atomic # in a row ↑, the Zeff ↑.
• Distance between two
nuclei is bond distance.
• Bonding radius is
estimated from known
structures of molecules.
How does Zeff help predict radii of atoms?
Which of the following atoms is larger?
Na
Z - S = Zeff
11-10 = 1
Mg
12-10 = 2
Cl
17-10 = 7
HINT: think Zeff , sizes of valence orbitals, # of protons
atomic radius (nm)
Zeff ∼ Z - S
Sizes of Atoms and Ions
0.17
0.15
0.13
0.11
0.09
0
2
4
Zeff
6
8
Atomic size varies in periodic table.
• Down a group, atoms become larger.
• Across a period, atoms become smaller.
From left to right :
# protons increases
size (n value) of orbital is constant
Zeff increases
As Zeff increases, e– pulled tightly to nucleus,
radii of atoms decrease
Atomic radius trends
Trends in the Sizes of Ions
• Ion size is distance between ions in ionic solid.
•Cations lose e– from largest orbitals;
smaller than parent atom.
• Ion size depends on Zeff, n of valence e– .
and # of e–
• Which is bigger, Na or Na+?
• Which is bigger, Cl or Cl–?
•Anions gain e– to largest orbital
larger than parent atom.
Fig 7.6
2
Trends in the Sizes of Ions
• Isoelectronic: atoms that have same e– config
He and Li+
• Which is bigger in this iso-electronic series?
O2- F- Na+ Mg2+ Al3+
1.18
Ionization energy
• 1st ionization energy, IE, is E to remove 1 e– from
1 atom in gas phase:
Na(g) → Na+(g) + e--
• :
O2- > F- > Na+ > Mg2+ > Al3+
• The 2nd IE, is the E is for 2nd e–
Na+(g) → Na2+(g) + e-• Larger IE = harder to remove e–.
Fig 7.7
According to electron configurations, from which orbitals are
electrons lost?
(hint: think e– config)
According to e– configurations, from which orbitals are e– lost?
Li:
Li:
Be:
Element
1st
electron
1s22s1
1s22s22p1
According to e– configurations, from which orbitals are e– lost?
Be:
1s22s2
B:
Li: 1s22s1
1st
electron
2nd
electron
3rd
electron
4th
electron
X
X
Li
X
Be
Be:
1s22s2
1st
electron
2nd
electron
2nd
electron
Li
3rd
electron
4th
electron
X
X
Li
X
Be
Be
Element
B
B
Element
B
X(g) → X+(g) + e−
Li
2nd
electron
a
b
3rd electron
Energies in kJ/mol to remove electrons
Be
a
b
c
B
a
b
c
Element
1st
electron
2nd
electron
4th electron
X
X
1s2
2s2
2s1
1s2
2p1
2s2
2s1
X
1s2
X(g) → X+(g) + e−
Energies in kJ/mol to remove electrons
Using what you know, talk with your
neighbor to explain the data below:
For each element, which is largest, a or b or c or d?
1st electron
4th
electron
X(g) → X+(g) + e−
Ionization energy
Energies in kJ/mol to remove electrons
3rd
electron
2s1
Ionization energy
Ionization energy
B: 1s22s22p1
B:
3rd
electron
Li
520
7,300
Be
900
1,750
14,800
B
800
2,430
3,700
4th
electron
25,000
Element
1st
electron
2nd
electron
3rd
electron
Li
520
7,300
Be
900
1,750
14,800
B
800
2,430
3,700
4th
electron
25,000
d
3
Ionization Energy
Ionization Energy
• Note HUGE increase in IE when a core e– is removed.
• IE ↓ down a group, easier to remove e– from
bigger atoms
• Easier to remove outermost e– lower in the group.
• IE generally ↑ from L to R
• As Zeff ↑, more difficult to remove e–.
• Two exceptions:
1st p e– and 4th p e–
Fig 7.10
Table 7.2
Electron Affinities
Electron Affinities
• Electron affinity: opposite of ionization energy.
∆E when an atom(g) gains an e– to form anion (g) :
Cl(g) + e- → Cl– (g) EA = –349 kJ/mol
• EA can either be exo or endothermic:
Ar(g) + e- → Ar–(g) EA = 35 kJ/mol
Fig 7.11
+ EA,
Hi IE,
No rxns
Low IE, low –EA,
Lose e–
Hi –EA, Hi IE,
Gain e–
4