Download Electron Configurations - Birmingham City Schools

Electron Configurations
The Bohr model had limitations...the math only worked for 1 electron atoms (basically,
only hydrogen). Since it couldn't predict the chemical properties (ionization energy,
for example) for anything other than hydrogen, it had to be modified.
Other scientists found discrepancies in the trend in ionization energy. Others found
details in the spectra (remember Bohr's model came from element spectra) that
weren't explained by Bohr's model.
On top of that, scientists working with light noticed that light acted as both a particle
and a wave. A scientists named deBroglie turned that around and suggested that
particles could be waves...so maybe an electron can behave as a wave!
Experiments with electrons showed that electrons can exhibit wave like properties (like
diffraction).
Erwin Schroedinger came up with an equation to describe an electron as a wave. He
won the Nobel prize for this. It's called the Schroedinger Wave Equation.
The equation requires 4 numbers to describe an electron and we call these quantum
numbers:
Four quantum numbers:
n, l, ml, ms
n = energy level
l = angular momentum (orbital shape: s, p, d, f)
ml = Magnetic (axis: x, y, z)
ms = spin (up or down arrow)
It's not really important for you to know much about quantum numbers. However,
each of the quantum numbers have physical meaning...and that's the part you need
to know.
Energy level is pretty straight forward, but now there are sublevels which we call
s
p
d
f
Each sublevel has a certain number of orbitals. An orbital is a 3-d shape that
represents where an electron has 90% probability of being found.
(So for those of you paying attention...the other 10 % is the rest of the universe...crazy, huh?)
Energy
Level
Number Sublevels:
of
sublevels
# of orbitals in
that sublevel
Total number of
electrons in
sublevel
1
1
s
1
2
2
2
s
p
1
3
2
6
3
3
s
p
d
1
3
5
2
6
10
4
4
1
3
5
7
2
6
10
14
s
p
d
f
NOTICE THE PATTERNS!! It will help a lot!
There are 4 kinds of sublevels: s, p, d, f
Here's what they look like:
s orbitals are just spheres:
p orbitals have two lobes and there are 3 different p orbitals (one for each axis: px, py,
pz
d orbitals are more complicated, one of them is just odd. 5 different d orbitals:
f orbitals are a lot more complicated. There are seven of them.
Okay, so what do you do with that information? You describe how electrons are
arranged in atoms...There are two ways:
Orbital Diagrams
and
Electron Configurations
1s2 2s2 2p2
Each arrow represents an electron.
The big number = energy level
Each line represents an orbital, the sublevel
and energy level are written below
Letter = sublevel
Exponent = # of electrons in that sublevel
The spaces separate sublevels
*notice that there are 3 lines for p since p has 3
different orbitals
This method does not give as much
information s orbital diagrams.
Both of the above describe an element with six electrons...which one is it?
Answer: Carbon!
Here are some rules for writing orbital diagrams and electron configurations:
Aufbau Principle: An electron occupies the lowest energy orbital that can receive it
Low Energy
1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6
6d10
1s
2s 2p
3s 3p 3d
4s 4p 4d 4f
5s 5p 5d 5f
6s 6p 6d
7s
High Energy
5s2 4d10 5p6 6s2 4f14 5d10 6p6 7s2 5f14
1s
2s 2p
3s 3p 3d
4s 4p 4d 4f
5s 5p 5d 5f
6s 6p 6d
7s
Notice the pattern above... if you use slanted arrows, starting with 1s as shown in the
next picture, you will get the same order as the yellow highlight.
The following orbital diagram would be wrong, because the 2s is not filled
completely. Since 2s is lower in energy, it needs to be filled first.
Pauli Exclusion Principle:
No two electrons in the same atom can have the same set of four quantum numbers
Hund’s Rule
1. Orbitals of the same energy are occupied by on electron before any orbital is
occupied by a second electron
Which means that the following orbital diagram is WRONG:
2. All electrons in singly occupied orbitals must have the same spin.
Which means that the following orbital diagram is WRONG:
Orbital Diagrams
1. Determine the number of electrons from the atomic number on the periodic
table
2. Determine the highest energy level
3. Draw lines for each of the orbitals available to each energy level
4. Fill in electrons using up and down arrows according to Aufbau, Hund, and Pauli
principles
1.
2.
3.
4.
5.
Electron Configurations
1s22s22p6
Determine the number of electrons using the atomic number from the periodic
table
Write the energy level number
Write the orbital number
Write the number of electrons in that orbital as a superscript
Repeat until all electrons have been used
Example: Ti has 22 electrons
Using the yellow strip above, add up the exponents until you get to 22. Notice
that I changed the last exponent so that all the exponents add up to 22.
1s2 2s2 2p6 3s2 3p6 4s2 3d2
Periodic Trends
This is a map! If I say tell me the
element that ends with 5p2 you
would
- count down 5 rows
- go to 2nd column in p section
Answer: Sn (Tin)
Patterns on the periodic table:
1. s orbitals can hold 2 electrons, all elements in the first 2 families on the PT have
electron configurations than end with s1 (group 1) or s2 (group 2)
2. The p sublevel can hold 6 electrons and there are 6 columns in the purple
section that correspond to ending electron configs of p1, p2, p3 and so on.
3. The d sublevel can hold 10 electrons and there are 10 columns in the green
section that correspond to ending electron configurations of d1, d2, d3 and so
on.
4. The f sublevel can hold 14 electrons and there are 14 columns in the green
section that correspond to ending electron configurations of f1, f2, f3 and so on.
5. The energy level is the same as the row/period number
a) except for d electrons which are one less than the row number
b) also excpet for f electrons, but this is easy to remember because the first time
f electrons show up is in the fourth energy level, so the first row of the f section
is 4f and the second row is 5f
6. Notice that each of the sections generally corresponds to elements that have
similar properties
a) S BLOCK: Alkali metals and alkaline earth: react strongly with water, are
metallic and form only one type of ion
b) P BLOCK: All the non-metals (although, it also includes some metals...there's
always exceptions in chemistry!)
c) D BLOCK: Transition metals: metallic, have multiple oxidation states, often
form colored compounds.
d) F BLOCK: Many are radioactive/unstable
Noble Gas Configurations
These are a shorter version of electron configuration.
1. Identify the noble gas that comes BEFORE the element you are writing the
electron configuration for – EVEN IF IT IS FARTHER AWAY.
2. Put the noble gas in square brackets
3. Write the rest of the electron config.
Ne:
Na:
1s2 2s2 2p6
1s2 2s2 2p6 3s1
Notice the similarity of Na and Ne?
Abbreviate Na's electron config: [Ne] 3s1
Electron Configurations for Ions!
Follow the same procedures as previous examples:
Na+ has 10 electrons instead of 11
Write the electron config so that exponents add up to 10
1s2 2s2 2p6 3s1
Na+: 1s2 2s2 2p6
Na:
Notice that the electron configuration for the sodium ion is now the same as Neons.
Remember, elements react to become stable, like their nearest noble gas. In Na's
case, it's neon.
O2- has 8 electrons instead of 6
1s2 2s2 2p4
O2-: 1s2 2s2 2p6
O:
Extra information: Quantum Numbers
**Quantum numbers are written for electrons.
**It is easiest to write quantum numbers if you have an orbital diagram
1. Determine the energy level of the electron
a. Assign 1 for energy level 1, 2 for level 2, etc.
2. Determine the orbital
Orbital
Angular
Momentum
Quantum
Number
s
0
p
1
d
2
f
3
3. Determine the axis (degenerate orbital) that the electron is on
Orbital
Magnetic Quantum number
s
0
p
____
____
____
-1
0
+1
d
____
-2
4. Determine the spin
a. Up = + ½
b. Down = - ½
____
-1
____
0
____
+1
____
+2