Download Electrons in the Atom

Chapter 13: Orbitals and
Electron Configurations
Atomic Orbitals and Electron Configurations (Chap 13)
Quantum Mechanics
(Heisenburg and Schrodinger)
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 more as waves
(like light waves) 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.
In this case it is losing the energy and dropping a level.
Atomic Orbitals
http://milesmathis.com/bohr2.jpg
 Much like the Bohr model, the energy
levels in quantum mechanics describe
locations where you are likely to find
an electron.
 Orbitals are “geometric shapes”
around the nucleus where electrons
are found.
 Quantum mechanics calculates the
probabilities where you are “likely”
to find electrons.
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.
Sub-levels = Specific
Atomic Orbitals
 There are 4 types of atomic orbitals:
 s, p, d and f
 Each of these sub-levels represent the
blocks on the periodic table.
Blue = s block
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.
 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.




s sublevel begins at energy level 1
p sublevel begins at energy level 2
d sublevel begins at energy level 3
f sublevel begins at energy level 4
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-orbital = rosette
VERY COMPLICATED
SHAPE
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.”
No more than 2 Electrons
in Any Orbital…ever.
http://www.fnal.gov/pub/inquiring/timeline/images/pauli.jpg
Wolfgang Pauli, yet
another German
Nobel Prize winner
 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
opposite.
 We usually represent this
with an up arrow and a
down arrow.
Hund’s Rule
http://intro.chem.okstate.edu/AP/2004Norman/Chapter7/Lec111000.html
 Hund’s Rule states that when you
get to degenerate orbitals, you fill
them all half way first, and then
you start pairing up the electrons.
 Each orbital will get one electron
before any gets two.
Think of this rule as…..
“SHARE BEFORE YOU PAIR”
Don’t pair up the 2p electrons
until all 3 orbitals are half full.
PRACTICE!
 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.
 One last thing. Look at the previous slide and look
at just hydrogen, lithium, sodium and potassium.
 Notice their electron configurations. Do you see
any similarities?
 Since H and Li and Na and K are all in Group 1A,
they all have a similar ending. (s1)
Electron Configurations
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.
 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
opposite. We usually represent this with an
up arrow and a down arrow.
 Hund’s Rule states that each orbital will get
one electron before any gets two. Think of
this rule as…..“SHARE BEFORE YOU
PAIR”
The End