Clackamas Community College

CH 104: INTRODUCTORY CHEMISTRY

 

 

Contact instructor:

Eden Francis

Physical Science
19600 Molalla Avenue
Oregon City, OR 97045
(503) 594-3352
TDD (503) 650-6649

Lesson 6: Electron Configurations

In this section, we will look at the relationship between electron configurations and the shape of the periodic table. We will start with a short review of electron configurations, take a closer look at the patterns of electron configurations of various groups in the periodic table, and look at some short cuts for determining electron configurations from position on the periodic table.

Electron Configurations | Review | Electron Configurations of Groups | Shortcuts

Electron Configurations

In Lesson 5 electron configurations were introduced, but you did not need to learn how to write electron configurations.  In this section we will go over the basics of electron configurations.

An electron configuration is, basically, an "address" that shows where the electrons can be found in an atom.  We use the wave-mechanical model when creating electron configurations.  Remember that the wave-mechanical model has atomic energy levels that are subdivided into "sublevels", which are further divided into "orbitals". 

The farther the energy level is from the nucleus, the higher the energy and more electrons can fit into the larger space.  The first energy level, the level closest to the nucleus, has space for only one sublevel - called the 1s sublevel - which holds one orbital.  Each orbital, no matter which sublevel or energy level it is in, can hold a maximum of two electrons.  So the first energy level can hold two electrons. 

If an atom has more than two electrons, then they will be found in higher energy levels.  In the second energy level there is more space than in the first energy level and we can see that there are two sublevels, called the 2s and 2p sublevels.  The 2s sublevel has just one orbital, which can hold two electrons, but the 2p sublevel has three orbitals, which can hold a total of six electrons (two electrons per orbital).

Examine the diagram below and then look at example 3 in your workbook. Both diagrams are read from the bottom to the top.

The 4th energy level has 4 sublevels called:

4f, has 7orbitals called 4f and can have 14 e-
4d, has 5 orbitals called 4d and can have 10e-
4p, has 3 orbitals called 4p and can have 6e-
4s, has 1 orbital called 4s and can have 2e-

The 3rd energy level has 3 sublevels called:

3d, has 5 orbitals called 3d and can have 10e-
3p, has 3 orbitals called 3p and can have 6e-
3s, has 1 orbital called 3s and can have 2e-

The 2nd energy level has 2 sublevels called:

2p, has 3 orbitals called 2p and can have 6e-
2s, has 1 orbital called 2s and can have 2e-

The 1st energy level has 1 sublevel called:

1s, has 1 orbital called 1s and can have 2e-

 

For atoms with just a few electrons, the order in which the electrons fill up the energy levels is exactly the way we would predict: the first two electrons fill up the 1s sublevel, the next two electrons go into the 2s sublevel, the next six go into 2p, and so on.  However, once we get to the third energy level, things get a little strange.  It turns out that it takes a little less energy for an electron to go into a 4s orbital than into a 3d orbital, and so the 4s fills up before electrons fill the 3d.  Example 3 is a diagram of the energy levels and how they overlap. [Take a ruler or straightedge and lay it across the page so that it is perpendicular to the line labeled " Arbitrary Energy Scale."  Line it up so that it goes through the circle labeled 1s.  Now move the ruler up the page - keeping it perpendicular to the Energy Scale line - and look at the order in which it gets to each sublevel.  It should progress in the expected order (1s, then 2s, then 2p, etc) until just after you reach 3p.  The next sublevel you get to is 4s - before 3d!  Keep going until you get to the top of the diagram and notice any other "irregularities".]

Writing Electron Configurations

When you write an electron configuration you write down where each electron can be found.  Remember that for a neutral atom, the number of electrons must be equal to the atomic number (the number of protons).  Electron configurations show the sublevel (1s, 2s, 2p, etc) and the number of electrons in each sublevel as a superscript after the sublevel.

Let's try the electron configuration of helium.  The atomic number of helium is 2, so a neutral atom of helium must have 2 electrons.  The electron configuration would be 1s2.

Why don't you try your hand at a few others?  Remember to fill each sublevel completely before moving on to the next sublevel.  (You can leave a sublevel partially full if you "run out" of electrons before you fill it up.)

Write out the electron configurations for:    Be, O, Ar, and Ca.

You can use the diagram in Ex. 3 to help or you can use this "filling order" shortcut to help figure out the electron configurations.

 

Answers:

Be has 4 electrons, 1s22s2

O has 8 electrons, 1s22s22p4

Ar has 18 electrons, 1s22s22p63s23p6

Ca has 20 electrons, 1s22s22p63s23p64s2

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Review of Electron Configurations

Historically, the properties of elements and the way those elements combined with other elements were the basis for the development of the periodic table (particularly the short form in the latter half of the 19th century). When physicists developed an understanding of how atoms were constructed (particularly the configuration of the electrons), it was possible to relate that structure to the periodic table. Our next lessons (on atomic bonding) will emphasize how the electron configurations dictate the chemical properties and combining patterns of the elements.

For now let’s focus on how electron configurations are related to the shape of the periodic table. To begin, I would like you to take a moment to write out the complete electron configuration for these three elements: hydrogen, lithium and sodium. Hydrogen has one electron. Lithium has three electrons. Sodium has eleven electrons. So take a moment to do that before continuing on.

H 1
Li 3
Na 11

These are the electron configurations you should have.  Notice that these three elements are all in group IA of the periodic table. Notice also that each of those electron configurations ends in s1. It is a different s1 for each element--it is 1s1 for hydrogen, 2s1 for lithium, and 3s1 for sodium--but notice the similarity in that they all end in s1.

H 1 1s1
Li 3 1s22s1
Na 11 1s22s22p63s1

Electron Configuration and the Periodic Table

Now take a look at the periodic table shown below. (A similar table is shown in example 2 in your workbook.) It is another periodic table, but instead of having atomic weights it has the last part of the electron configuration for each of the elements.

Periodic Table with Partial Electron Configurations
H
1s1
  He
1s2
Li
2s1
Be
2s2
  B
2s22p1
C
2s22p2
N
2s22p3
O
2s22p4
F
2s22p5
Ne
2s22p6
Na
3s1
Mg
3s2
Al
3s23p1
Si
3s23p2
P
3s23p3
S
3s23p4
Cl
3s23p5
Ar
3s23p6
K
4s1
Ca
4s2
Sc
4s23d1
Ti
4s23d2
V
4s23d3
Cr
4s13d5
Mn
4s23d5
Fe
4s23d6
Co
4s23d7
Ni
4s23d8
Cu
4s13d10
Zn
4s23d10
Ga
4s24p1
Ge
4s24p2
As
4s24p3
Se
4s24p4
Br
4s24p5
Kr
4s24p6
Rb
5s1
Sr
5s2
Y
5s24d1
Zr
5s24d2
Nb
5s14d4
Mo
5s14d5
Tc
5s24d5
Ru
5s14d7
Rh
5s14d8
Pd
4d10
Ag
5s14d10
Cd
5s24d10
In
5s25p1
Sn
5s25p2
Sb
5s25p3
Te
5s25p4
I
5s25p5
Xe
5s25p6
Cs
6s1
Ba
6s2
La*
6s25d1
Hf
6s25d2
Ta
6s25d3
W
6s25d4
Re
6s25d5
Os
6s25d6
Ir
6s25d7
Pt
6s15d9
Au
6s15d10
Hg
6s25d10
Tl
6s26p1
Pb
6s26p2
Bi
6s26p3
Po
6s26p4
At
6s26p5
Rn
6s26p6
Fr
7s1
Ra
7s2
Ac§
7s26d1
                             
 
  * Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu  
§ Th Pa U Np Pu Am Cm Bk Cf Es Fm Md No Lr

In the remaining pages of this section, we will take a closer look at the electron configurations of various groups in the periodic table, look at some short cuts for determining electron configurations, and look at how atomic orbitals are related to the periodic table.

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Electron Configurations of Groups

In this subsection we will look at the patterns of electron configurations that exist in the periodic table--first for the representative elements, then the transition metals, and then the overall patterns.

Representative Elements

H, Li and Na all have their final electron put into an "s" sublevel; their electron configurations all end with s1. If you look at the rest of the elements in that group you will see that the similarity continues on down the periodic table. All of the elements in group IA end in s1 for their electron configuration. What is different about them is which level of s1 they end with.

Notice also that this periodic table only shows the last part of the electron configuration. The reason for this is that the chemical properties of an element are determined by the outermost electron structure. So this periodic table was prepared with just those outermost electron configurations written on it.

I A
H   1s1
Li   2s1
Na   3s1
K   4s1
Rb   5s1
Cs   6s1
Fr   7s1

 

Now look at group IIA, all of them end in s2 but a different s2 for each one: 2s2 for beryllium, 3s2 for magnesium, and so on. Next, what comes after the s sublevel is filled? If we go one electron past magnesium, we get aluminum. Its outermost electron configuration is 3s23p1. Notice that if you look above and below aluminum, all of the outermost electron configurations for this group of elements have the same pattern. What is different is which period they are in. Boron is 2s22p1. Aluminum is 3s23p1. Gallium is Ga and is an element you didn't have to memorize, but from its position on the periodic table, you can see that it is 4s24p1. And so on, on down.

IIA IIIA

Be
2s2

B
2s22p1

Mg
3s2
Al
3s23p1
Ca
4s2
Ga
4s24p1
Sr
5s2
In
5s25p1
Ba
6s2
Tl
6s26p1
Ra
7s2

 

If you keep going to the right on the periodic table, notice how with each column there is one more p electron in each of the energy levels, right over to the inert gases.

B
2s22p1
C
2s22p2
N
2s22p3
O
2s22p4
F
2s22p5
Ne
2s22p6
Al
3s23p1
Si
3s23p2
P
3s23p3
S
3s23p4
Cl
3s23p5
Ar
3s23p6
Ga
4s24p1
Ge
4s24p2
As
4s24p3
Se
4s24p4
Br
4s24p5
Kr
4s24p6
In
5s25p1
Sn
5s25p2
Sb
5s25p3
Te
5s25p4
I
5s25p5
Xe
5s25p6
Tl
6s26p1
Pb
6s26p2
Bi
6s26p3
Po
6s26p4
At
6s26p5
Rn
6s26p6

 

Transition Metals

Notice that when you deal with the transition metals, the pattern is not quite so distinct.

Sc
4s23d1
Ti
4s23d2
V
4s23d3
Cr
4s13d5
Mn
4s23d5
Fe
4s23d6
Co
4s23d7
Ni
4s23d8
Cu
4s13d10
Zn
4s23d10

There are patterns here but the patterns are not as reliable. Let's start with element number 21, scandium. From calcium to scandium, the additional electron goes into the 3d1. So scandium has 4s23d1, or if you prefer, 3d14s2, as shown in example 2 in your workbook. You can write it in either direction, but it is 4s2 and 3d1. Then the next electron (the 22nd one) also goes in the 3d sublevel . Thus titanium, Ti, has 3d2 as part of its electron configuration. Vanadium, then, is 3d3 and so on across except that you will notice that chromium does not have 3d4 like you might expect. We continue with manganese at 3d5 and 4s2 and continue across with 3d6, 3d7, 3d8 but copper is not 3d9. You will not be expected to know what the exact electron configuration is for the transition elements when they alter that configuration a little bit. So when you are asked to figure out the electron configuration for a transition element, it probably will be one that follows the pattern, rather than one that doesn't.

 

Overall Patterns

The point here is to emphasize some of the patterns that exist in the relationship between the electron configurations of the elements and the location of the elements on the periodic table and even the shape of the periodic table. (Refer to ex. 4 in your workbook.)

Looking at this example note how the periodic table can be broken up into s, p, d and f blocks. The first two columns of the periodic table (groups IA and IIA) are in the s block because the elements in these two groups have their outermost electrons in s orbitals. Note that the s block has two groups because atoms can put two electrons in an s sublevel.

s1   s2
s1 s2   p1 p2 p3 p4 p5 p6
               
    d1 d2 d3 (d4) d5 d6 d7 d8 (d9) d10            
                                   
                                   
                                   
 
  f1 (f2) f3 f4 f5 f6 f7 (f8) f9 f10 f11 f12 f13 f14  
                           

 

The p block consists of groups IIIA through the inert gases. These are the elements which have their last electrons in p orbitals. Note that there are six groups in this block because atoms can put up to six electrons into a p sublevel.

The transition elements comprise the d block. The d block has 10 columns because up to 10 electrons will fit into a d sublevel.

The f block at the bottom of the periodic table has 14 columns because up to 14 electrons can fit into an f sublevel.

Remember the pattern we worked with that resulted in the electron configurations. 1s, then 2s, 2p, 3s, then 3p, 4s, then 3d, 4p, 5s. That resulted in electron configurations that looked like this. 1s2, 2s2, 2p6, 3s2, 3p6, 4s2, 3d10 and so on. The periodic table shows that same arrangement in a different way. The first period has 1s1 then 2. The second period has 2s1 then 2, then 2p1, 2, 3, 4, 5, and 6. Similarly, the third period has 3s1 then 2, then 3p1, 2, 3, 4, 5 and 6. The fourth period with the transition elements in the middle has 4s1 then 2, followed by 3d1, 2, 3, 4 is altered, then 5, 6, 7, 8, 9 is also altered, then 10, followed by 4p1, 2, 3, 4, 5 and 6. The periodic table continues on showing the pattern dictated by the electron configurations.

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Shortcuts to Electron Configurations

Next, let's use the periodic table to provide a short cut method of getting the last part of the electron configuration for nearly any element.  We will be able to determine both the outermost electron configurations and also the electron configuration of the last-occupied sublevel.

Let's take calcium, element number 20, as our first example. It is in the fourth period and in the s block, so its electron configuration will end with 4s. Calcium is in the second column of the s block, so it has two electrons in that sublevel and the outermost electron configuration is 4s2.  (Its last-occupied sublevel configuration is also 4s2.)

Fourth 
Period 



4s1
20
Ca
4s2

 

Iron, element number 26, is also in the fourth period but it is in the d block. You need to remember that 3d gets electrons after 4s. Because iron is in the sixth spot of the d block, the last electron in the configuration is 3d6. The s electrons are also important to the chemistry of the transition metals so they are usually included as outermost electrons. Therefore the electron configuration ends with 4s23d6 (or alternately 3d64s2).  The last-occupied sublevel configuration is just the 3d6 (because that is the place where the last electron is located).

Fourth 
Period 



3d1


3d2


3d3


3d4


3d5
26
Fe
4s23d6

 

For elements in the p block, count down the periods to get the energy level of the last electrons and count over from the edge of the p block to determine how many electrons are in the last sublevel. For example, nitrogen is in the second period and is the third element over in the p block, so its electron configuration ends with 2p3. If you are figuring the electron configuration for the purpose of determining the number of valence electrons, be sure to include electrons in the s sublevel. For nitrogen that would be 2s22p3 for a total of 5 valence electrons.  What would the last-occupied sublevel configuration be?  Right - 2p3!

Second 
Period 



2p1


2p2
7
N
2s22p3

Another example is iodine, element number 53. It is in the fifth period and it is the fifth element over in the p block, so its electron configuration ends with 5p5. Again, if our concern is valence electrons we would include the s sublevel, 5s25p5.

Fifth 
Period 



5p1


5p2


5p3


5p4
53
I
5s25p5

As you can see, using the periodic table to determine either the outermost electron configuration or the last-occupied sublevel configuration is much easier than writing out the entire electron configuration!

Practice

For practice, try figuring out the last-occupied sublevel electron configuration for the elements shown below (they are also listed in example 5 in your workbook). Try to use the periodic table in workbook example 1, without the electron configurations, for reference. Then use the periodic table in workbook example 2 to check your answers. That table includes the s electrons for the p block and d block elements. Since we're looking for just the last-occupied sublevel, don't include them in your answers for this question.

Na W Br B Ni Ba

 

Shortcuts and Complete Configurations

This short-cut method can also be used to determine the complete electron configurations. Use the location of the element on the periodic table to figure out the last entry in the configuration, as you have just done. Use the shape of the periodic table to remind you of the order in which the sublevels fill. Do this by reading the blocks in each period in order. First 1s. Then 2s, 2p. Then 3s, 3p. Then 4s, 3d, 4p. Next 5s, 4d, 5p, and so on. Keep adding electrons to that pattern until you have reached what you know to be the end of the electron configuration for that element based on its location on the periodic table.

1s   1s
2s   2p
3s 3p
4s 3d 4p
5s 4d 5p
6s * 5d 6p
7s § 6d 7p
 

*
§

4f  
5f

 

In the lab, we have a poster that illustrates not only the repeating nature of the electron configuration but also has an emphasis on the shapes of the orbitals. For future reference, note that the atomic size trends (discussed later in this lesson) are also apparent in this diagram. Next time you are in lab, take a moment to look at the illustrations to get a better idea of how the "layering" of the sublevels and orbitals happens.

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