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Chapter 13
Electrons in Atoms
Quantum Mechanics
http://www.meta-synthesis.com/webbook/30_timeline/310px-Bohr-atom-PAR.svg.png
Better than any previous model,
quantum mechanics does explain
how the atom behaves.
Quantum mechanics treats electrons
not as particles, but as waves (like
light) which can gain or lose energy.
But they can’t gain or lose just any
amount of energy. They gain or lose
a “quantum” of energy.
A quantum is just an amount of energy that the electron needs to
gain (or lose) to move to the next energy level. Max Planck,
another German Nobel Prize winning scientist first came up with
this idea.
What the heck is a
Quantum?
http://www.blogcdn.com/www.slashfood.com/media/2008/08/splenda425.jpg
http://upload.wikimedia.org/wikipedia/commons/e/e9/Sucralose2.png
Think of a quantum as a “packet” of
energy, much like a sugar packet at
a restaurant. A sugar packet
contains a teaspoonful of sugar.
If the electron absorbs energy, it
moves to a higher energy level. If it
emits (loses) energy, it moves to a
lower energy level.
But like Bohr suggested in his
model, the electron has to gain or
lose exactly the right amount. That
amount is a quantum of energy.
C12H19O8Cl3 is the formula for sucralose, which is the chemical name for Splenda. That
“beast” molecule is sucralose. It’s an Organic compound.
Atomic Orbitals
http://milesmathis.com/bohr2.jpg
The energy levels in quantum mechanics describe
locations where you are likely to find an electron.
These locations are called ORBITALS. Orbitals are
“geometric shapes” around the nucleus where
electrons are found.
Quantum mechanics can CALCULATE where those
orbitals are located.
The 4 different types of orbitals are s, p, d, and f.
Sound familiar?
Atomic Orbitals
http://courses.chem.psu.edu/chem210/quantum/quantum.html
Of course, you could find an electron anywhere if you
looked hard enough.
So scientists agreed to limit these calculations to locations
where there was at least a 90% chance of finding an
electron.
Think of orbitals as sort of a "border” for spaces around the
nucleus inside which electrons are allowed. No more than 2
electrons can ever be in 1 orbital. The orbital just defines an
“area” where you can find an electron.
What is the chance of finding an electron in the nucleus?
Yes, of course, it’s zero. There aren’t any electrons in the
nucleus.
Atomic Orbitals
http://www-hep.phys.unm.edu/~gold/phys492/orbitals.gif
3s
2s
1s
Quantum mechanics doesn’t
predict SPECIFIC orbits, like the
Bohr model does.
Energy Levels
http://www.chem4kids.com/files/art/elem_pertable2.gif
Quantum mechanics has a
principal quantum number. It is
represented by a little n. It
represents the “energy level”
similar to Bohr’s model.
Red
Orange
Yellow
Green
Blue
Indigo
Violet
n=1
n=2
n=3
n=4
n=5
n=6
n=7
n=1 describes the first energy level
n=2 describes the second energy
level
Etc.
Each energy level represents a
period or row on the periodic
table. It’s amazing how all this
stuff just “fits” together.
Sub-levels = Specific
Atomic Orbitals
Each energy level has 1 or more “sublevels” which describe the specific
“atomic orbitals” for that level.
Blue = s block
n = 1 has 1 sub-level (the “s” orbital)
n = 2 has 2 sub-levels (“s” and “p”)
n = 3 has 3 sub-levels (“s”, “p” and
“d”)
n = 4 has 4 sub-levels (“s”, “p”, “d”
and “f”)
There are 4 types of atomic orbitals:
s, p, d and f
Each of these sub-levels represent the
blocks on the periodic table.
Shapes of These Orbitals
(the nucleus is ALWAYS at the center of the orbital)
The s orbital looks like a ball or sphere.
The p orbital looks like a dumb-bell.
These orbitals are all perpendicular to each other.
The d orbitals have two shapes.
4 of the 5 look like “4-leaf clovers.”
The 5th one looks like a “big dumb-bell” with a
“hula-hoop” around the middle.
The shapes of the f orbitals are complex. They
are on the next slide, but you don’t need to
remember them, nor will they be on the test.
Orbitals
http://media-2.web.britannica.com/eb-media/54/3254-004-AEC1FB42.gif
http://upload.wikimedia.org/wikipedia/commons/thumb/e/e1/D_orbitals.svg/744px-D_orbitals.svg.png
s
p
d
In the s block, electrons are going into s orbitals.
In the p block, the s orbitals are full. New electrons are going into the p orbitals.
In the d block, the s and p orbitals are full. New electrons are going into the d orbitals.
What about the f block?
f orbitals
http://antoine.frostburg.edu/chem/senese/101/electrons/faq/f-orbital-shapes.shtml
To Summarize
Energy
Level
Sublevels
Total Orbitals
Total
Electrons
Total Electrons
per Level
n=1
s
1 (1s orbital)
2
2
n=2
s
p
1 (2s orbital)
3 (2p orbitals)
2
6
8
n=3
n=4
s
Complete
p
d
1 (3s orbital)
the3 (3p
chart
in your
orbitals)
5 (3d orbitals)
2
notes6 as
10
18
we discuss this.
The first level (n=1) has an s orbital. It has only 1.
There
are no1 other
orbitals in the
s
(4s orbital)
2 first energy
32 level.
p
3 (4p orbitals)
6
f
7 (4f orbitals)
14
Wedcall this 5orbital
the 1s orbital.
(4d orbitals)
10
“g orbitals”
http://jeries.rihani.com/index3.html
Another hypothesis by Glenn Seaborg is that element number 121
will start “the g block.”
The “g” block will be another grouping, similar to the Lanthanides
and Actinides, of 18 elements.
Since this is all science fiction, you obviously don’t have to know
what g orbitals look like.
A collection of Dr. Seaborg’s most important scientific publications
has been published in a book called “Modern Alchemist.”
To date, no elements have been discovered which have 8s electrons.
Element 119 is predicted to be the first element in the “8th period.”
Island of Stability
http://www.nytimes.com/1999/02/27/us/glenn-seaborg-leader-of-team-that-found-plutonium-dies-at-86.html
This is another hypothesis from Dr. Seaborg. His thought was
that element 114 would be an “island of stability,” especially if it
also had 184 neutrons.
Most synthesized elements only last for fractions of seconds.
However, in 1998 researchers synthesized element 114 and it
lasted for 30 seconds. Perhaps this is the “shore” of the Island of
Stability that Dr. Seaborg hypothesized.
The element 114 was made using some of the original Pu-244
that Dr. Seaborg himself made in the early 1940s. They
bombarded plutonium with Ca-48 atoms to form some of the
new element 114.
Element 114 is now know as Flerovium (symbol Fl)
Island of Stability
http://www.sciencecodex.com/files/Island%20of%20Stability%201.jpg
http://physicsworld.com/cws/article/print/19751
Famous picture of the “Island of Stability” showing the island off in the distance (top
right) with 114 protons and 184 neutrons. What would the atomic mass of the island be?
An element with Z = 184 is also predicted to be another “island of stability.”
Electron Configurations
What do I mean by “electron
configuration?”
The electron configuration is the
specific way in which the atomic
orbitals are filled.
Think of it as being similar to your
address. The electron configuration
tells me where all the electrons “live.”
Rules for Electon
Configurations
https://teach.lanecc.edu/gaudias/scheme.gif
In order to write an electron
configuration, we need to know the
RULES.
3 rules govern electron
configurations.
Aufbau Principle
Pauli Exclusion Principle
Hund’s Rule
Using the orbital filling diagram at
the right will help you figure out
HOW to write them
Start with the 1s orbital. Fill each
orbital completely and then go to
the next one, until all of the
elements have been acounted for.
Fill Lower Energy
Orbitals FIRST
Each line represents
an orbital.
1 (s), 3 (p), 5 (d), 7 (f)
High Energy
http://www.meta-synthesis.com/webbook/34_qn/qn3.jpg
The Aufbau Principle states
that electrons enter the
lowest energy orbitals first.
The lower the principal
quantum number (n) the
lower the energy.
Within an energy level, s
orbitals are the lowest
energy, followed by p, d and
then f. F orbitals are the
highest energy for that level.
Low Energy
No more than 2 Electrons in
Any Orbital…ever.
http://www.fnal.gov/pub/inquiring/timeline/images/pauli.jpg
The next rule is the Pauli Exclusion Principal.
The Pauli Exclusion Principle states that an atomic
orbital may have up to 2 electrons and then it is
full.
The spins have to be paired.
We usually represent this with an up arrow and a
down arrow.
Wolfgang Pauli, yet
another German Nobel
Prize winner
Since there is only 1 s orbital per energy level, only
2 electrons fill that orbital.
Quantum numbers describe an electrons position, and no 2 electrons
can have the exact same quantum numbers. Because of that,
electrons must have opposite spins from each other in order to
“share” the same orbital.
Hund’s Rule
http://intro.chem.okstate.edu/AP/2004Norman/Chapter7/Lec111000.html
Hunds Rule states that when you
get to degenerate orbitals, you fill
them all half way first, and then
you start pairing up the electrons.
What are degenerate orbitals?
Degenerate means they have the
same energy.
So, the 3 p orbitals on each level
are degenerate, because they all
have the same energy.
Don’t pair up the 2p electrons
until all 3 orbitals are half full.
Similarly, the d and f orbitals are
degenerate too.
Objective D
NOW that we know the rules, we can try to write
some electron configurations.
Remember to use your orbital filling guide to
determine WHICH orbital comes next.
Lets write some electron configurations for the first
few elements, and let’s start with hydrogen.
Electron Configurations
Element
Configuration
Element
Configuration
H Z=1
1s1
He Z=2
1s2
Li Z=3
1s22s1
Be Z=4
1s22s2
B
Z=5
1s22s22p1
C
Z=6
1s22s22p2
N Z=7
1s22s22p3
O
Z=8
1s22s22p4
F
1s22s22p5
Ne Z=10
1s22s22p6
(2p is now full)
Na Z=11
1s22s22p63s1
Cl Z=17
1s22s22p63s23p5
K Z=19
1s22s22p63s23p64s1
Sc Z=21
1s22s22p63s23p64s23d1
Fe Z=26
1s22s22p63s23p64s23d6
Br Z=35
1s22s22p63s23p64s23d104p5
Z=9
Note that all the numbers in the electron configuration add up to the atomic
number for that element. Ex: for Ne (Z=10), 2+2+6 = 10
Electron Configurations
of Alkali Metals (and H)
Element
Configuration
H Z=1
1s1
Li Z=3
1s22s1
Na Z=11
1s22s22p63s1
K Z=19
1s22s22p63s23p64s1
This similar configuration causes them to behave the same
chemically.
It’s for that reason they are in the same family or group on
the periodic table.
Each group will have the same ending configuration, in this
case something that ends in s1.
Emission Spectra =
Fingerprint of the Elements
http://www.cbu.edu/~jvarrian/252/emspex.jpg
Atomic emission spectrum is sometimes
called a line spectrum, to distinguish it from
the continuous spectrum.
Emission Spectra =
Fingerprint of the Elements
Atomic emission spectra are “unique.” You can
use the spectrum to identify the element.
It’s like a “fingerprint” which identifies the light
as coming from hydrogen, and not from
something else.
Scientists can look at light from a distant star
and analyze it and determine what types of
elements make up that star.
Just by looking at the light! Pretty impressive.
The End