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ELECTRON CONFIGURATION Tutorial and Quiz

Last update (7/29/10)

In the last tutorial, I already introduced you to quantum numbers which define the size, shape, orientation, and spin of electrons. Since that tutorial was already information overload, I didn't go into how this progression is represented.
Like I said before, learning chemistry is like being a symbologist. Symbols is your specialty. Just take pride in it and the learning doesn't hurt so much.
chemistry in new light

As you know, all areas of chemistry can be broken down into these 3 areas of focus. In learning about electron configuration, the building blocks are simply electrons with some mention of protons. Force and energy are also considered in understanding how the electrons choose their configuration. Mathematics in this topic is on two extremes. First graders can do most of it because it's like counting to 14. Then there's the theory behind it that's way behind this course.

Let's do a quick review.

This is the electron of hydrogen. We often show it as a particle, a wave, or electron cloud. Roll cursor over image to see the transformation.

 

Bohrs Atom DeBroglie Atom

Bohr’s atom showed electrons as particles similar to planets going around the sun. de Broglie’s atom treats electrons more as waves with wave patterns that fit symmetrically within the atom. In both cases, the energy levels of the electrons must go up by an incremental quantity (quantum).

Both of these played a role in the modern theory of the atom and the structure of the electrons. As elements were built by adding one proton and one electron at a time, the electrons would find a position and shape that maximized their distance from each other (from repulsion) yet kept them as close as possible to the positive nucleus (attraction). The way they positioned themselves followed a fairly basic pattern known as its quantum number.

Electron spin animation
Each electron has their own unique 4 quantum numbers represented by n, l, m, and s. This is known as the Pauli Exclusion Principle named after Wolfgang Pauli and his statement that electrons are excluded from having the same quantum numbers. In other words, they can't be in the same place at the same time with one exception. If their spins are opposite, they can. A spinning charge produces a magnetic field. If they spin in opposite directions, then the north and south poles will align and they will attract each other (opposites poles attract). At least this is one way the spin of electrons is visualized.
Bohrs Atom DeBroglie Atom

Sometimes these quantum numbers are represented as n, l, ml, and ms, but here we will just use n, l, m, and s. The n is called the principle quantum number. It's often referred to as the shell number because in Bohr's atom, each increasing orbit forms a shell around the inner ones. In the past these shells were called K, L, M, N, and O. Those correspond to n=1, 2, 3, 4, & 5. You may still see reference to these letters in the literature.
    "n" defines the basic size of the the electron. In other words, the size of the region where it spends most of its time.

Also, "n" corresponds to the row (or period) on the Periodic Table.

By the way, the first elements in the left two columns have outer electrons in the s orbital.

 

 

Bohr's orbits
This is the hydrogen atom with its one electron in its ground state (lowest energy). The "1" in front of "1s1" means this is n=1. The "1" exponent means there's just "1" electron in that orbital. The "↑" indicates a +½ spin. Whether electrons actually spin is debated but it's one property of electron's configuration. When n=1, then l and m quantum numbers can only be zero. If the electron was boosted to the second shell (n=2), then the arrow below would be in the 2s box.

n=1, l=0, m=0
       
1s
 
2s
 
2p
 
 
s=+½
           
helium atom
This is the helium atom. It has 2 electrons. These are able to share the same s orbital because of opposite spins. The "2" in "1s2" means there are 2 electrons in that orbital. "s" orbitals can only hold 2 electrons, so that's as high as you will ever see it. The 1 in "1s2" means this is the s orbital belonging to the n=1 shell.

n=1, l=0, m=0
       
1s2
 
2s
 
2p
↑↓
 
 
s=+½,-½   
           
lithium
This is lithium. The third electron is in shell n=2.

n=1, l=0, m=0
 
n=2, l=0, m=0
   
1s2
 
2s1
 
2p
↑↓
 
 
s=+½,-½    
 
s=+½
       
Beryllium
This is beryllium. The fourth electron completes the 2s orbital. Again, it has to have opposite spin.

n=1, l=0, m=0
 
n=2, l=0, m=0
   
1s2
 
2s2
 
2p
↑↓
 
↑↓
 
s=+½,-½   
 
s=+½,-½   
       

boron

orbitals

This is boron. The 2s orbital is full, so the expansion begins in the 2p orbital. "2p" means this orbital belongs to the n=2 shell. The fifth electron is the first p orbital electron. It is drawn as a particle in the upper image but actually has the dumb bell shape of a p orbital shown in the lower image. Notice the first p orbital occurs when the l quantum number is 1. So when you hear l=1, think p orbital. Also, note, that l starts at 0, and then counts up to n-1. Since this n=2 row, l started as l=0 for the s orbital and can only go up to l=1 (one less than the n value).

n=1, l=0, m=0
 
n=2, l=0, m=0
 
n=2, l=1, m=0
1s2
 
2s2
 
2p1
↑↓
 
↑↓
 
s=+½,-½
 
s=+½,-½
 
s=+½
   

azimuth
Note: The l quantum number is a lower case "L" that is italicized. The letter "l" and number "1" look alike but remember they are different. There are 3 names for the l quantum number: azimuthal, angular, and orbital. The word "azimuthal" comes from "azimuth" which is from Arabic "as sumut" meaning "the way" or "the direction". It is used in astronomy and gunnery to indicate the clockwise angle from North. In the image the azimuth angle is about 200°. So we see why it is also called the angular quantum number. It indicates a rotational motion of the wave. It's also called the orbital quantum number because it is this l quantum number that determines the kind of orbital (shape of the electron). This is important because the different types of orbitals (different l numbers) have different strengths in chemical bonding.
carbon
This is carbon. It is adds another electron to the p orbital. This time I'm showing the electron cloud of the p orbitals. Notice the x, y, and z axes. The first 3 p orbital electrons will align with each axis. Carbon only has 2 p orbital electrons. I put the first one (pink) on the x axis and the second one (yellow) on the y axis. They don't share the same space by having opposite spins because that takes more energy than filling up an empty p orbital. Changing orientation keeps the electrons farther apart. The different orientations are indicated with the m quantum number. This is also called the magnetic quantum number. Notice the m quantum number starts with the negative of the l quantum number. Since we are at the l=1 quantum number (p orbital), the m quantum numbers start out as -1, then go up to 0 on the next electron, and +1 on the third electron. Each time it changes its orientation along a different axis.

n=1, l=0, m=0
 
n=2, l=0, m=0
 
n=2, l=1
1s2
 
2s2
 
2p2
↑↓
 
↑↓
 
s=+½,-½    
 
s=+½,-½    
 
m=-1
s=
m=0
s=+½
m=+1
oxygen
This is oxygen. We skipped nitrogen, but nitrogen added one electron to the z axis of the p orbital. Oxygen has 8 protons and 8 electrons. The first 4 are in the first 2s orbitals. The last electron was forced to share one of the p orbitals with an existing electron. I put it with the one on the z axis, but they are all the same. Notice the m quantum number increased to +1 for the 3rd and 4th electrons aligning them with the z axis. However, they must have opposite spins to share that same space. By the way, the 2 unpaired electrons in m=-1 and m=0 orientations are responsible for the magnetic properties of oxygen.

n=1, l=0, m=0
 
n=2, l=0, m=0
 
n=2, l=1
1s2
 
2s2
 
2p4
↑↓
 
↑↓
 
↑↓
s=+½,-½    
 
s=+½,-½    
 
m=-1
s=
x
m=0
s=+½
y
m=+1
s=+½,-½
z  
neon
This is neon. We skipped fluorine, which added one more electron to fill 2p2y. At neon all of the p orbitals are filled. It has 3 orientations (x, y, and z) that are filled with 2 electrons each. This is a very stable configuration for the 8 electrons in the 2s and 2p orbitals. You may know it as the famous octet that many elements try to reach. So we see all of the p orbital electrons have the same n=2 principal quantum number and the same l=1 angular quantum number. Pairs of electrons of opposite spins are aligned along x, y, and z axes indicated by the m=-1, m=0, and m=-1 magnetic quantum numbers.

n=1
 
n=2
l=0
 
l=0
 
l=1
1s2
 
2s2
 
2p6
↑↓
 
↑↓
 
↑↓
↑↓
↑↓
m=0
s=+½,-½    
 
m=0
s=+½,-½    
 
m=-1
s=+½,-½
x
m=0
s=+½,-½
m=+1
s=+½,-½ 
z
After all of these symbols, test your knowledge. Match the left symbol with the one best match on the right.
Problem 1: s orbital a) dumbbell b) n=1 c) l=0 d) l=1
Problem 2: p orbital a) dumbbell b) n=1 c) l=0 d) +½
Problem 3: n a) number b) l=0 c) shell d) +½
Problem 4: angular quantum number a) spherical b) n=1 c) shell d) l=0
Problem 5: a) +½ b) x c) s d) l=0
Problem 6: Pauli exclusion principle a) spherical b) l=0 c) l=1 d)↑↓
Problem 7: magnetic quantum number a) spherical b) x,y,z c) l=0 d)↑↓

Problem 8: How many electrons are in this element that has this electron configuration?
1s22s22p3

Problem 9: How many electrons are needed for this element to complete the octet in shell 2 (n=2)?
1s22s22p3

Problem 10: What is this element?
1s22s22p3


After the 2p electron orbital is filled, the elements start filling the 3s orbital. The filling is just like it did in row 2. So if you look at rows 2 and 3 on the Periodic Table, you will see they look the same. In other words, 2 electrons fill the s orbital and 6 electrons fill the p orbital. The only difference in the quantum numbers is the n=3, which is reflected in the electron configuration notation as 3s23p6. Note: I left off the "s=" below to save space. The spins of +½ and -½ are obvious.
n=1
 
n=2
 
n=3
l=0
 
l=0
 
l=1
 
l=0
 
l=1
1s2
 
2s2
 
2p6
 
3s2
 
3p6
↑↓
 
↑↓
↑↓
↑↓
↑↓
 
↑↓
 
↑↓
↑↓
↑↓
m=0
+½,-½  
 
m=0
+½,-½ 
 
m=-1
+½,-½
x
m=0
+½,-½
m=+1
+½,-½ 
z
 
m=0
+½,-½ 
 
m=-1
+½,-½
x
m=0
+½,-½
m=+1
+½,-½ 
z
Electron Energy levels
According to quantum theory, the angular (orbital) number represented by "l" counts up to one less than the principle quantum number of n. For example, if n=3, then "l" would go 0, 1, 2. Therefore, for n=3, 0=3s orbital, 1=3p orbital, and 2=3d orbital. In the image on the left, we can see the energy levels of this progression goes up. However, because the energy levels of 4s orbital electrons are lower than the 3d electrons, they get filled before the 3d electrons. Look at the Periodic Table above. The first two columns for row 4 are the elements where the 4s electrons are being filled. The yellow block of 10 elements that follow are the ten 3d electrons that are filled after the 4s electrons. So the order of electron filling is from lowest energy to higher energy. This is called the Aufbau principle named after the German word that means "building up."
In other words, the general pattern of electron filling is guided by the counting up of the quantum numbers.
(n=1: l=0, m=0)
(n=2: l=0, m=0 > l=1, m=-1, 0, 1)
(n=3: l=0, m=0 > l=1, m=-1 ,0, 1 > l=2, m=-2, -1, 0, 1, 2)
(n=4: l=0, m=0 > l=1, m=-1 ,0, 1 > l=2, m=-2, -1, 0, 1, 2 > l=3, m=-3, -2, -1, 0, 1, 2, 3 )
The above progress is modified to let electrons orbitals that have the lowest energy fill first. The first instance of that (as mentioned) is when the 4s electrons fill before the 3d electrons. This also happens when 5s electrons fill before 4d electrons (Period 5). On row 6 (period 6) the two 6s electrons are filled followed by the first 5d electron for element 57 (Lanthanum). But then the 4f electrons start filling. That's the first green row you see at the bottom. That starts with element 58. On some Periodic Tables element 57 (Lanthanum) is included on the bottom row. They call this the lanthanum series. That the series of elements where the f orbital electrons start to get filled.
Below is progression up through the transition metals that is the d block on the Periodic Table (two panels above). The last element that has all of these electrons is zinc. It's electron configuration would be written as 1s22s22p63s23p64s23d10. As you can see it gets a bit long. To shorten it they usually use the element symbol of the noble gas that came before the element. Argon has the electron configuration of 1s22s22p63s23p6, so for zinc, it's OK to write [Ar]4s23d10. The [Ar] means the inner electrons have the same electron configuration as argon. Then on top of the inner electrons are two 4s electrons and ten 3d electrons.
n=1
 
n=2
 
n=3
 
n=4
 
n=3
l=0
 
l=0
 
l=1
 
l=0
 
l=1
 
l=0
 
l=2
1s2
 
2s2
 
2p6
 
3s2
 
3p6
 
4s2
 
3d10
↑↓
 
↑↓
↑↓
↑↓
↑↓
 
↑↓
 
↑↓
↑↓
↑↓
 
↑↓
 
↑↓
↑↓
↑↓
↑↓
↑↓
m=0
+½,-½  
 
m=0
+½,-½ 
 
m=-1
+½,-½
x
m=0
+½,-½
m=+1
+½,-½ 
z
 
m=0
+½,-½ 
 
m=-1
+½,-½
x
m=0
+½,-½
m=+1
+½,-½ 
z
 
m=0
+½,-½ 
 
m=-2
+½,-½
yz
m=-1
+½,-½
xz 
m=0
+½,-½ 
xy
m=+1
+½,-½
x2-y2
m=+2
+½,-½
z2 
d orbitals Recall the magnetic quantum number (m) dictates an orientation of the electron. In the above chart I used yz, xz, xy, etc. for the d orbital electrons. The image on the left shows these orientations. Again, the whole purpose is to give electrons more space between them because of the repulsion. The wave shape of the five d electrons on the left is bit more complex than those of the s and p orbitals.
Below is the electron configuration up to krypton. Again, the filling of the 3d orbital interrupts the filling of the 4p orbital. The pattern below repeats itself for the n=5 shell. Then on the n=6 shell (row 6), there's an interruption of the filling of the d orbitals by f orbitals. The reason, as usual, is they have lower energy levels.
n=1
 
n=2
 
n=3
 
n=4
 
n=3
 
n=4
l=0
 
l=0
 
l=1
 
l=0
 
l=1
 
l=0
 
l=2
 
l=1
1s2
 
2s2
 
2p6
 
3s2
 
3p6
 
4s2
 
3d10
 
4p6
↑↓
 
↑↓
↑↓
↑↓
↑↓
 
↑↓
 
↑↓
↑↓
↑↓
 
↑↓
 
↑↓
↑↓
↑↓
↑↓
↑↓
 
↑↓
↑↓
↑↓
m=0
+½,-½  
 
m=0
+½,-½ 
 
m=-1
+½,-½
x
m=0
+½,-½
m=+1
+½,-½ 
z
 
m=0
+½,-½ 
 
m=-1
+½,-½
x
m=0
+½,-½
m=+1
+½,-½ 
z
 
m=0
+½,-½ 
 
m=-2
+½,-½
yz
m=-1
+½,-½
xz 
m=0
+½,-½ 
xy
m=+1
+½,-½
x2-y2
m=+2
+½,-½
z2 
 
m=-1
+½,-½
x
m=0
+½,-½
m=+1
+½,-½ 
z
Below is the entire Periodic Table which shows the electronic configuration for all elements. This was tough to create, but I got it done. The progression is pretty consistent except for a few exceptions. For example, look at chromium (group 6, period 4) and copper (group 11, period 4). Chromium shows 4s13d5. Instead of having 2 electrons in the 4s2 and 4 electrons filling the first 4 electrons in 3d orbital (3d4) as predicted, it took an electron from 4s2 and leaving it as 4s1 and filled the first 5 electrons in 3d orbital (3d5). There still are just 6 electrons (group 6) but they are not split 2 in the 4s orbital and 4 in the 3d orbital. The 4s13d5 means 1 is in the 4s orbital and 5 are in the 3d orbital. So every once in awhile, you will see these anomalies.  It's not fully understood why this happens, but it's probably because that arrangement is more stable (has lower energy).
   
G R O U P   #
 
1
2
3
4
5
6
7
8
9
10
11
12
13
14
15
16
17
18
P
E
R
I
O
D

#

(n)
1
1
H
|
|
|
|
|
|
|
|
|
|
|
|
|
|
|
|
1
He
2
2s1
Li
2s2
Be
|
|
|
|
|
|
|
|
|
|
2p1
B
2p2
C
2p3
N
2p4
O
2p5
F
2p6
Ne
3
3s1
Na
3s2
Mg
|
|
|
|
|
|
|
|
|
|
3p1
Al
3p2
Si
3p3
P
3p4
S
3p5
Cl
3p6
Ar
4
4s1
K
4s2
Ca
3d1
Sc
3d2
Ti
3d3
V
4s13d5
Cr
3d5
Mn
3d6
Fe
3d7
Co
3d8
Ni
4s13d10
Cu
3d10
Zn
4p1
Ga
4p2
Ge
4p3
As
4p4
Se
4p5
Br
4p6
Kr
5
5s1
Rb
5s2
Sr
4d1
Y
4d2
Zr
4d3
Nb
5s14d5
Mo
4d5
Tc
4d6
Ru
4d7
Rh
4d8
Pd
5s14d10
Ag
4d10
Cd
5p1
In
5p2
Sn
5p3
Sb
5p4
Te
5p5
I
5p6
Xe
6
6s1
Cs
6s2
Ba
5d1
*La 
5d2
Hf
5d3
Ta
6s15d5
W
5d5
Re
5d6
Os
5d7
Ir
5d8
Pt
6s15d10
Au
5d10
Hg
6p1
Tl
6p2
Pb
6p3
Bi
6p4
Po
6p5
At
6p6
Rn
7
7s1
Fr
7s2
Ra
6d1
+Ac 
6d2
Rf
6d3
Db
7s16d5
Sg
6d5
Bh
6d6
Hs
6d7
Mt
6d8
Ds
7s16d10
Rg
6d10
Uub
7p1
Uut
7p2
Uuq
7p3
Uup
7p4
Uuh
7p5
Uus
7p6
Uuo
8
8s1
Uue
8s2
Ubn
 
*5d1
La
5d14f1
Ce
4f3
Pr
4f4
Nd
4f5
Pm
4f6
Sm
4f7
Eu
5d14f7
Gd
4f9
Tb
4f10
Dy
4f11
Ho
4f12
Er
4f13
Tm
4f14
Yb
4f145d1
Lu
 
6d1
+Ac 
6d2
Th
6d15f2
Pa
5f4
U
5f5
Np
5f6
Pu
5f7
Am
6d15f7
Cm
5f9
Bk
5f10
Cf
5f11
Es
5f12
Fm
5f13
Md
5f14
No
5f147p1
Lr

The purpose of learning these electron configurations is to know what electrons are available for bonding. This gives you insight in how elements will bond. In the next chapter when you learn more about bonding, you will see the usefulness of these electron configurations.

Problem 11: Which orbital is being filled in the alkali metal column?
Problem 12: Find the element in period 3 and group 17. What is its full electron configuration starting with 1s2?
Problem 13: The electron configuration of indium is [Kr]4d105s25p1. Indium forms compounds with chlorine. One is indium (I) chloride, formula InCl. What electron do you think indium gave up?
Problem 14: Another compound of indium and chlorine is indium (III) chloride, formula InCl3. What electrons do you think indium gave up?
Problem 15: When zinc combines with non-metals, it always has an oxidation number of +2. The electronic configuration of zinc is [Ar]3d104s2, which is also written [Ar]4s23d10. What 2 electrons does zinc lose when it forms an ionic bond?
Problem 16: When aluminum combines with non-metals, it always has an oxidation number of +3. What 3 electrons does aluminum lose when it forms an ionic bond? (See table above).

Thanks for doing this tutorial and quiz. Send your answers to Ken Costello at chm151@chemistryland.com

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