1. .
The Onset of Quantum theory:
by Dr. Robert D. Craig, Ph.D.
In 1934, Schrödinger lectured at Princeton University; he was
offered a permanent position there, but did not accept it (just
like “Will Hunting!!!)
•!
5. Today’s Agenda
• More on Schrodinger (Who was He?
• Use other power point for history@@@
• More about his Cat!
• More Quantum numbers
• Hund’s Rule
• Pauli Exclusion principle
• Aufbah principle
6. We all make final together-2 questions
I have a sample with me!!!!
7. . Dr Poget – said you may see him!
• Dr Poget is free after this class-in office!!
• Would like to here from students after this
class!
9. .
• In 1934, Schrödinger lectured at
Princeton University; he was offered a
permanent position there, but did not accept
it (just like “Will Hunting!!!)
• Let’s look at what was going on at the time!
11. FOR OUR PURPOSES
• Principle quantum number (n)
Azimuthal quantum number (l)
Magnetic quantum number (m)
Spin quantum number (s)
12. LITHIUM SYSTEM
• Electron 1: N =1, L=0, Ml= 0 , Ms = +1/2
• Electron 2: N =1, L=0, Ml= 0 , Ms = -1/2
• Electron 3: N =2, L=0, Ml= 0 , Ms = +1/2
13. Pauli Exclusion Principle
• Pauli Exclusion Principle
• The second major fact to keep in mind is the
Pauli Exclusion Principle which states that no two
electrons can have the same four quantum
numbers. The first three (n,l, and ml) may be
similar but the fourth quantum number must be
different. We are aware that in one orbital a
maximum of two electrons can be found and the
two electrons must have opposing spins.
14. Pauli Exclusion Principle
• That means one would spin up ( +1/2) and the
other would spin down (-1/2). This tells us
that each subshell has double the electrons
per orbital. The s subshell has 1 orbital that
can hold to 2 electrons, the p subsheel has 3
orbitals that can hold up to 6 electrons, the d
subshell has 5 oribtals that hold up to 10
electro
15. BERYLLIUM
• Electron 1: N =1, L=0, Ml= 0 , Ms = +1/2
• Electron 2: N =1, L=0, Ml= 0 , Ms = -1/2
• Electron 3: N =2, L=0, Ml= 0 , Ms = +1/2
• Electron 4: N =2, L=0, Ml= 0 , Ms = +1/2
16. BORON
• Electron 1: N =1, L=0, Ml= 0 , Ms = +1/2
• Electron 2: N =1, L=0, Ml= 0 , Ms = -1/2
• Electron 3: N =2, L=0, Ml= 0 , Ms = +1/2
• Electron 3: N =2, L=0, Ml= 0 , Ms = +1/2
• Electron 3: N =2, L=1, Ml= 0 , Ms = +1/2
17. BORON
• Because the 1s and 2s orbitals are filled the
fifth electron has to be assigned to a p-orbital
20. CARBON –IS THE MOST INTERESTING
last time!!!
• Electron 1: N =1, L=0, Ml= 0 , Ms = +1/2
• Electron 2: N =1, L=0, Ml= 0 , Ms = -1/2
• Electron 3: N =2, L=0, Ml= 0 , Ms = +1/2
• Electron 4: N =2, L=0, Ml= 0 , Ms = +1/2
• Electron 5: N =2, L=1, Ml= 1 , Ms = +1/2
• Electron 6: N =2, L=1, Ml= -1 , Ms = +1/2
24. SO –IS NITROGEN- ns2 nsp3
it follows Hunds rule
• Hund's rule: every orbital in a subshell is
singly occupied with one electron before any
one orbital is doubly occupied, and all
electrons in singly occupied orbitals have the
same spin
25. Nitrogen
• Electron 1: N =1, L=0, Ml= 0 , Ms = +1/2
• Electron 2: N =1, L=0, Ml= 0 , Ms = -1/2
• Electron 3: N =2, L=0, Ml= 0 , Ms = +1/2
• Electron 4: N =2, L=0, Ml= 0 , Ms = +1/2
• Electron 5: N =2, L=1, Ml= 1 , Ms = +1/2
• Electron 6: N =2, L=1, Ml= 0 , Ms = +1/2
• Electron 7: N =2, L=1, Ml= -1 , Ms = +1/2
26. Nitrogen
• . Atoms at ground states tend to have as many
unpaired electrons as possible.
27. As far apart as possible and lined up
• When visualizing this processes, think about
how electrons are exhibiting the same
behavior as the same poles on a magnet
would if they came into contact;
• as the negatively charged electrons fill
orbitals they first try to get as far as possible
from each other before having to pair up.
28. Example
• If we look at the correct electron
configuration of Nitrogen (Z = 7), a very
important element in the biology of plants: 1s2
2s2 2p3
30. Oygen
• Oxygen (Z = 8) its electron configuration is: 1s2
2s2 2p4
• Oxygen has one more electron than Nitrogen
and as the orbitals are all half filled the
electron must pair up
31. The Aufbau Process
• Aufbau comes from the German word
"Aufbauen" which means "to build". When
writing electron configurations, we are
building up electron orbitals as we proceed
from atom to atom.
• However, there are some exceptions to this
rule
32. *Examples
• If we follow the pattern across a period from B
(Z=5) to Ne (Z=10) the number of electrons
increase and the subshells are filled.
• Here we are focusing on the p subshell in
which as we move towards Ne, the p subshell
becomes filled.
33. In order
• B (Z=5) configuration: 1s2 2s2 2p1
• C (Z=6) configuration:1s2 2s2 2p2
• N (Z=7) configuration:1s2 2s2 2p3
• O (Z=8) configuration:1s2 2s2 2p4
• F (Z=9) configuration:1s2 2s2 2p5
• Ne (Z=10) configuration:1s2 2s2 2p6
35. True picture is orbital levels adjust
themselves- due to sheilding
36. Rules for Assigning Electron Orbitals
• The order of levels filled would look like this:
• 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f,
5d, 6p, 7s, 5f, 6d, and 7p
37. Exceptions to Electron Configuration
Trends
• The reason these exceptions occur is because
some elements are more stable with less
electrons in some subshells and more
electrons within others. A list of the
exceptions to the Aufbau process can be
found below
40. Copper
• Copper: Z:27 [Ar] 3d104s1
• Example: In the following configuration,
• Cu: [Ar]4s23d9, copper's d shell is just one away
from stability, and therefore, one electron
from the s shell jumps into the d shell:
[Ar]4s13d10. This way, the d shell is full, and is
therefore stable, and the s shell is half full,
and is also stable.
42. Most stable configurations####
• Chromium has a configuration of [Ar]4s13d5,
• The stability rule applies to atoms in the same
group as chromium and copper.
• If one of these atoms has been ionized, that is,
it loses an electron, it will come from the s
orbital rather than the d orbital.
• the configuration of Cu+ is [Ar]4s03d10.
43. Loss of outer electrons
• If one of these atoms has been ionized, that is,
it loses an electron, it will come from the s
orbital rather than the d orbital. For instance,
the configuration of Cu+ is [Ar]4s03d10. If more
electrons are removed, they will come from
the d orbital
45. On final
• The spin of an electron creates a magnetic
field (albeit ridiculously weak), so unpaired
electrons create a small magnetic field. Paired
electrons have opposite spin, so the magnetic
fields cancel each other out, leading to
diamagnetism.
46. 3 terms on final
• Diamagnetism is actually a very weak repulsion to
magnetic fields. All elements have diamagnestism to
some degree. It occurs when there are pair electrons.
• Paramagnetism is an attraction to external magnetic
fields. It is also very weak. It occurs whenever there is
an unpaired electron in an orbital.
• Ferromagnetism is the permanent magnetism that we
encounter in our daily lives. It only occurs with three
elements: iron (Fe), nickel (Ni), and cobalt (Co).
48. Writing Electron Configurations
• When writing the electron configuration we
first write the energy level (the period) then
the subshell to be filled and the superscript,
which is the number of electrons in that
subshell. The total number of electrons as
mentioned before is the atomic number, Z.
Using the rules from above, we can now start
writing the electron configurations for all the
elements in the periodic table.
49. Methods
• There are three main methods used to write
electron configurations: orbital diagrams, spdf
notation, and noble gas notation. Each
method has its own purpose and each has its
own drawbacks.
50. Orbital Diagrams
• As seen in some examples above, the orbital
diagram is a visual way to reconstruct the
electron configuration by showing each of the
separate orbitals and the spins on the
electrons. This is done by first determining the
subshell (s,p,d, or f) then drawing in each
electron according to the stated rules above.
51. Example
• Electron configuration for aluminum.
• If we look at the periodic table we can see
that its in the p-block as it is in group 13. Now
we shall look at the orbitals it will fill: 1s, 2s,
2p, 3s, 3p. We know that aluminum
completely fills the 1s, 2s, 2p, and 3s orbitals
because mathematically this would be
2+2+6+2=12.
52. Example
• The block that the atom is in (in the case for
aluminum: 3p) is where we will count to get
the number of electrons in the last subshell
(for aluminum this would be one electron
because its the first element in the period 3 p-
block).
53. Aluminum system
• From this we can construct the following:
• Note that in the orbital diagram
54. • Note that in the orbital diagram
• Note that in the orbital diagram, the two
opposing spins of the electron can be seen.
This is why it is sometimes useful to think
about electron configuration in terms of the
diagram
55. Filling orbitals: In action
• http://en.wikibooks.org/wiki/General_Chemis
try/Filling_Electron_Shells
57. Electron Notation using spdf
• This is a much simpler and efficient way to
portray electron configuration of an atom. A
logical way of thinking about it is that all we
have to do is to fill orbitals as we move across
a period and through orbital blocks
58. Electron Notation using spdf
• As we move across we simply count how
many elements fall in each block. We know
that yttrium is the first element in the fourth
period d-block, thus this corresponds to one
electron in the that energy level. To check our
answer we would just add all the superscripts
to see if we get the atomic number. In this
case 2+2+6+2+6+2+10+6+2+1= 39 and Z=39
thus we have the correct answer.
59. Electron Notation using spdf
• Example
• Vanadium (V, Z=23) lies in the transition metals
at the four period in the fifth group. The noble
gas before it is argon, (Ar, Z=18) and knowing
that vanadium has filled those orbitals before it,
we will use argon as our reference noble gas. We
denote the noble gas in the configuration as the
symbol, E, in brackets: [E] configuration:
• Vanadium, V: [Ar] 4s2 3d3
61. Electron Configurations of Ions
• Writing electron configurations for ions,
whether it be cation or anion, is basically
exactly the same as writing them for normal
elements. All the same rules apply, except you
must take into account the gained or lost
electrons. For instance, when Potassium (K)
loses an electron it becomes K+ and has the
noble gas configuration of [Ar].
62. electron configuration
• K ([Ar]4s1) --> K+([Ar]) + e-
Therefore, the electron configuration for the
K+ ion is simply [Ar]
• When an atom, such as Chlorine (Cl) gains an
electron, it becomes Cl- and also has the
electron configuration of [Ar].
• Cl ([Ne]3s23p5) + e- --> Cl- ([Ar])
Yet again, the electron configuration is [Ar]
63. • For more complex ionic electron configurations, such
as an ion from the transition metals, the answer isn't
always a noble gas. Take Iron (Fe). The most common
irons for Iron are Fe2+ and Fe3+. Lets focus on Fe2+.
• Fe ([Ar]3d64s2) --> Fe2+ ([Ar]3d6) - 2e-
Here Iron loses two electrons. So thats two less electrons to fill orbitals. When you backtrack two electrons in Fe's original electron configuration you
[Ar]3d6 as Fe2+'s new configuration
get
• When writing the electron configuration for ions, treat
it like any normal element. Just remember to simply
add or subtract the gained or lost electrons when
filling out shells.
