Atomic Structure 5

1
Fred J. Grieman
Atomic Structure (5) Shells, Core and Valence
Electrons Chemistry and Valence
Electrons Periodic Table Groups and
Rows Representative Elements Transition
Metals Transition Metal Ions
2
Orbital Energy Diagram
Atom Electron configuration H 1s
unpaired electron magnetic atom!!
He 1s2
Li 1s22s
Be 1s22s2
B 1s22s22p
C 1s22s22p2
Auf Bau (Building up) principle To get GROUND
STATE electronic configuration 1) put e-s in
lowest energy orbital, 2) obey Pauli principle 3)
obey Hunds rule
Can do ions as well! O2- 1s22s22p6
N 1s22s22p3 O 1s22s22p4 F 1s22s22p5 Ne 1s22s22p6
Na 1s22s22p6 3s Mg 1s22s22p6 3s2
3
Shells, core, and valence electrons
Shells 1s 2s 2p 3s 3p 3d 4s 4p 4d
4f K L M N
subshells
subshells
subshells
Filled shells are called core electrons, the
core. These electrons are pulled in very close
to the nucleus.
Valence electrons electrons outside the core,
screened by the core important for chemistry
Ne
Consider Na 1s22s22p6 3s
? Ne 3s
core
valence
Li 1s22s ? He 2s
Similarity core shields 1 valence electron
similar chemistry Difference Ave. r(Na) gt
Ave. r(Li)
4
Another example O 1s22s22p4 ? He 2s22p4 S
1s22s22p6 3s23p4 ? Ne 3s23p4 Again similar
chemistry same of valence electrons Different
sizes different outer shell In general
Ave.r(K) lt Ave.r(L) lt
Ave.r(M) Ave.r in shell decreases with Z,
Why? Same of valence electrons similar
chemistry Periodic Table 1) Group
elements with same of valence electrons in
columns 2) Row of elements with same n atoms
of similar size History in text chemical
properties and mass, explained by
q.m. and orbital structure
5
Periodic Table Divided into types look first
at Representative Elements Representative
Elements Cores ns np valence
electrons Look at first 3 rows
Q.M. Pauli explains limits in rows
Ave.r big jumps
2s
2p
3s
3p
6
Non-metals p filling accepts e-s easily
Metals s filling give up e-s easily
Chalcogens Greek - Cu producing
Alkali Arabic ashes of saltwort
plant
Halogens Greek salt producing
Alkaline alkali like
Metalloids on border Tend to both give up
and accept e-s
Noble gases
Transition metals E(3d) lt E(4p)
E(4d) lt E(5p)
4s fill be- fore 3d E(4s)ltE(3d) E(5s)ltE(4d)
Unreactive Why? ns2np6 does not gain or
lose e-s easily Why?
7
More on transition metals later After transition
metals (d10), continue w/ p orbital filling
representative elements After Lanthanum 4f
filling Lanthanides
After Actnium 5f filling Actinides
3d
4d
5d
6d
8
Explanation of Orbital Energy Ordering
Core
E3s lt E3p lt E3d Because of penetration/less
shielding of nucleus
r2R2
For neutral atoms E4s lt E3d Due to less
screening of 4s electron
penetration
Alkali and Alkaline Metal before Transition metals
return
9
d-Transition metals
Exceptions to E(n1)s lt End
4s13d5
4s13d10
5s14d5
5s14d6
5s14d7
5s14d8
4d10
5s14d10
5s14d4
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Transition Metal Ions

• Transition metals tend to lose 4s e-s before 3d
  e-s !!?

Why?
For example V(4s23d3) ? V2 ( ? )
2e-
3d3
Consider V2 and Sc They are isoelectronic!!
For ion E 3d now lt E 4s
4
From H atom energy of orbital ? -Z2/n2 (no
shielding)
3
E
For neutral multielectron atoms, 4s lower than 3d
because decreased shielding overcomes orbital
energy difference
d
s
For V2 and Sc, shielding the same because
isoelectronic (both have 21 e-s)
But Now!! Z(V) 23 gt Z(Sc) 21 so
Orbital Energy Difference is greater
3 vs. 4 ?E Z2/n2
212/32 212/42 21.4 for Sc Z2/n2 232/32
232/42 25.7 for V2
diagram
4s
?E 25.7-21.4 4.3 increase of 20!!!
4s
3d
Next time Summary Thus Far
Decrease due to shielding not enough to overcome
orbital energy difference!! For ion, E4s gt
E3d
3d
Sc
V2
shielding
no
no
shielding
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Exam next Wednesday Closed book bring
calculator pencils Memorization Ry ao
size and energy
of atoms Chapters 15, 16,
17-1 Conceptual as well as working problems
You have to know it, not just be able to
work problems Begin Studying early! QA Session
Monday TBA