This is "Unit 3", section 3.5 from the book General Chemistry (v. 1.0).
3.5 Electron Configurations
Learning Objective
- To write the electron configuration of any element and relate its electron configuration to its position in the periodic table.
Now you can use the information you learned in the previous section to determine the electronic structure of every element in the periodic table. The process of describing each atom’s electronic structure consists, essentially, of beginning with hydrogen and adding one proton and one electron at a time to create the next heavier element in the table. All stable nuclei other than hydrogen also contain one or more neutrons. Because neutrons have no electrical charge, however, they can be ignored in the following discussion.
Electron Configurations of the Elements
The electron configurationThe arrangement of an element’s electrons in its atomic orbitals. of an element is the arrangement of its electrons in its atomic orbitals. By knowing the electron configuration of an element, we can predict and explain a great deal of its chemistry.
In Section 1.3 "Elements and the Periodic Table", we introduced the periodic table as a tool for organizing the known chemical elements. A periodic table is shown in Figure 3.5(a) "The Periodic Table". The elements are listed by atomic number (the number of protons in the nucleus), and elements with similar chemical properties are grouped together in columns.
Figure 3.5(a) The Periodic Table
Why does the periodic table have the structure it does? The answer is rather simple, if you understand electron configurations: the shape of the periodic table mimics the filling of the subshells with electrons.
Let us start with H and He. Their electron configurations are 1s1 and 1s2, respectively; with He, the n = 1 shell is filled. These two elements make up the first row of the periodic table (see Figure 3.5(b)).
Figure 3.5(b) The 1s Subshell
H and He represent the filling of the 1s subshell.
The next two electrons, for Li and Be, would go into the 2s subshell. Figure 3.5(c) shows that these two elements are adjacent on the periodic table.
Figure 3.5(c) The 2s Subshell
In Li and Be, the 2s subshell is being filled.
For the next six elements, the 2p subshell is being occupied with electrons. On the right side of the periodic table, these six elements (B through Ne) are grouped together (Figure 3.5(d)).
Figure 3.5(d) The 2p Subshell
For B through Ne, the 2p subshell is being occupied.
The next subshell to be filled is the 3s subshell. The elements when this subshell is being filled, Na and Mg, are back on the left side of the periodic table (Figure 3.5(e)).
Figure 3.5(e) The 3s Subshell
Now the 3s subshell is being occupied.
Next, the 3p subshell is filled with the next six elements (Figure 3.5(f)).
Figure 3.5(f) The 3p Subshell
Next, the 3p subshell is filled with electrons.
Instead of filling the 3d subshell next, electrons go into the 4s subshell (Figure 3.5(g)).
Figure 3.5(g) The 4s Subshell
The 4s subshell is filled before the 3d subshell. This is reflected in the structure of the periodic table.
After the 4s subshell is filled, the 3d subshell is filled with up to 10 electrons. This explains the section of 10 elements in the middle of the periodic table (Figure 3.5(h)).
Figure 3.5(h) The 3d Subshell
The 3d subshell is filled in the middle section of the periodic table.
And so forth. As we go across the rows of the periodic table, the overall shape of the table outlines how the electrons are occupying the shells and subshells.
Blocks in the Periodic Table
As you have learned, the electron configurations of the elements explain the otherwise peculiar shape of the periodic table. Although the table was originally organized on the basis of physical and chemical similarities between the elements within groups, these similarities are ultimately attributable to orbital energy levels and the way the individual subshells are filled in a particular order. As a result, the periodic table can be divided into “blocks” corresponding to the type of subshell that is being filled, as illustrated in Figure 3.5(l). For example, the two columns on the left, known as the s blockThe elements in the left two columns of the periodic table in which the ns orbital is being filled., consist of elements in which the ns orbitals are being filled. The six columns on the right, elements in which the np orbitals are being filled, constitute the p blockThe elements in the six columns on the right of the periodic table in which the np orbitals are being filled.. In between are the 10 columns of the d blockThe elements in the periodic table in which the (n − 1)d orbitals are being filled., elements in which the (n − 1)d orbitals are filled. At the bottom lie the 14 columns of the f blockThe elements in the periodic table in which the (n − 2)f orbitals are being filled., elements in which the (n − 2)f orbitals are filled. Because two electrons can be accommodated per orbital, the number of columns in each block is the same as the maximum electron capacity of the subshell: 2 for ns, 6 for np, 10 for (n − 1)d, and 14 for (n − 2)f. The f block could be part of the main body, but then the periodic table would be rather long and cumbersome. Figure 3.5(i) shows the blocks of the periodic table. Within each column, each element has the same valence electron configuration—for example, ns1 (group 1) or ns2np1 (group 13). As you will see, this is reflected in important similarities in the chemical reactivity and the bonding for the elements in each column.
Note the Pattern
Because each orbital can have a maximum of 2 electrons, there are 2 columns in the s block, 6 columns in the p block, 10 columns in the d block, and 14 columns in the f block.
Figure 3.5(i) Blocks on the Periodic Table
The periodic table is separated into blocks depending on which subshell is being filled for the atoms that belong in that section.
The electrons in the highest-numbered shell, plus any electrons in the last unfilled subshell, are called valence electronsThe electrons in the highest-numbered shell, plus any electrons in the last unfilled subshell.; the highest-numbered shell is called the valence shellThe highest-numbered shell in an atom that contains electrons.. (The inner electrons are called core electrons.) The valence electrons largely control the chemistry of an atom. If we look at just the valence shell's electron configuration, we find that in each column, the valence shell's electron configuration is the same. For example, take the elements in the first column of the periodic table: H, Li, Na, K, Rb, and Cs. Their electron configurations (abbreviated for the larger atoms) are as follows, with the valence shell electron configuration highlighted:
| H: | 1s1 |
| Li: | [He]2s1 |
| Na: | [Ne]3s1 |
| K: | [Ar]4s1 |
| Rb: | [Kr]5s1 |
| Cs: | [Xe]6s1 |
They all have a similar electron configuration in their valence shells: a single s electron. Because much of the chemistry of an element is influenced by valence electrons, we would expect that these elements would have similar chemistry-and they do. The organization of electrons in atoms explains not only the shape of the periodic table but also the fact that elements in the same column of the periodic table have similar chemistry.
Video: Lithium metal is cut, tested for conductivity, and reacted with water. The hydrogen gas evolved is collected in a test tube and ignited with a candle. Video Credit: Part of NCSSM CORE collection: http://www.dlt.ncssm. Please attribute this work as being created by the North Carolina School of Science and Mathematics. This work is licensed under Creative Commons CC-BY https://creativecommons.org/licenses/by/3.0/ via YouTube
Video: Sodium metal is cut, tested for conductivity, and reacted with water. The hydrogen gas evolved is collected in a test tube and ignited with a candle. Video Credit: Part of NCSSM CORE collection: http://www.dlt.ncssm. Please attribute this work as being created by the North Carolina School of Science and Mathematics. This work is licensed under Creative Commons CC-BY https://creativecommons.org/licenses/by/3.0/ via YouTube
Video: Potassium metal is cut, tested for conductivity, and reacted with water. The hydrogen gas evolved ignites spontaneously. Video Credit: Part of NCSSM CORE collection: http://www.dlt.ncssm. Please attribute this work as being created by the North Carolina School of Science and Mathematics. This work is licensed under Creative Commons CC-BY https://creativecommons.org/licenses/by/3.0/ via YouTube
The same concept applies to the other columns of the periodic table. Elements in each column have the same valence shell electron configurations, and the elements have some similar chemical properties. This is strictly true for all elements in the s and p blocks. In the d and f blocks, because there are exceptions to the order of filling of subshells with electrons, similar valence shells are not absolute in these blocks. However, many similarities do exist in these blocks, so a similarity in chemical properties is expected.
Similarity of valence shell electron configuration implies that we can determine the electron configuration of an atom solely by its position on the periodic table. Consider Se, as shown in Figure 3.5(j). It is in the fourth column of the p block. This means that its electron configuration should end in a p4 electron configuration. Indeed, the electron configuration of Se is [Ar]4s23d104p4, as expected.
Figure 3.5(j) Selenium on the Periodic Table
Example 3.5-1
From the element's position on the periodic table, predict the valence shell electron configuration for each atom. See Figure 3.5(k).
- Ca
- Sn
Solution
- Ca is located in the second column of the s block. We would expect that its electron configuration should end with s2. Calcium's electron configuration is [Ar]4s2.
- Sn is located in the second column of the p block, so we expect that its electron configuration would end in p2. Tin's electron configuration is [Kr]5s24d105p2.
Test Yourself
From the element's position on the periodic table, predict the valence shell electron configuration for each atom. See Figure 3.5(k).
- Ti
- Cl
Answer
- [Ar]4s23d2
- [Ne]3s23p5
Figure 3.5(k) Various Elements on the Periodic Table
Example 3.5-2
Write the electron configuration of mercury (Z = 80), showing all the inner orbitals.
Given: atomic number
Asked for: complete electron configuration
Strategy:
Using Figure 3.5(i) and the periodic table as a guide, fill the orbitals until all 80 electrons have been placed.
Solution:
By placing the electrons in orbitals following the order shown in Figure 3.5(i) and using the periodic table as a guide, we obtain
| 1s2 | row 1 | 2 electrons |
| 2s22p6 | row 2 | 8 electrons |
| 3s23p6 | row 3 | 8 electrons |
| 4s23d104p6 | row 4 | 18 electrons |
| 5s24d105p6 | row 5 | 18 electrons |
| row 1–5 | 54 electrons |
After filling the first five rows, we still have 80 − 54 = 26 more electrons to accommodate. According to Figure 3.5(i) , we need to fill the 6s (2 electrons), 4f (14 electrons), and 5d (10 electrons) orbitals. The result is mercury’s electron configuration:
1s22s22p63s23p64s23d104p65s24d105p66s24f145d10 = Hg = [Xe]6s24f145d10with a filled 5d subshell, a 6s24f145d10 valence shell configuration, and a total of 80 electrons. (You should always check to be sure that the total number of electrons equals the atomic number.)
Exercise
Although element 114 is not stable enough to occur in nature, two isotopes of element 114 were created for the first time in a nuclear reactor in 1999 by a team of Russian and American scientists. Write the complete electron configuration for element 114.
Answer: 1s22s22p63s23p64s23d104p65s24d105p66s24f145d106p67s25f146d107p2
Figure 3.5(l) Electron Configurations of the Elements
The electron configurations of elements indicated in red are exceptions due to the added stability associated with half-filled and filled subshells. The electron configurations of the elements indicated in blue are also anomalous, but the reasons for the observed configurations are more complex. For elements after No, the electron configurations are tentative.
Example 3.5-3
Use the periodic table to predict the valence electron configuration of all the elements of group 2 (beryllium, magnesium, calcium, strontium, barium, and radium).
Given: series of elements
Asked for: valence electron configurations
Strategy:
Step 1 Identify the block in the periodic table to which the group 2 elements belong. Locate the nearest noble gas preceding each element and identify the principal quantum number of the valence shell of each element.
Step 2 Write the valence electron configuration of each element by first indicating the filled inner shells using the symbol for the nearest preceding noble gas and then listing the principal quantum number of its valence shell, its valence orbitals, and the number of valence electrons in each orbital as superscripts.
Solution:
Step 1 The group 2 elements are in the s block of the periodic table, and as group 2 elements, they all have two valence electrons. Beginning with beryllium, we see that its nearest preceding noble gas is helium and that the principal quantum number of its valence shell is n = 2.
Step 2 Thus beryllium has an [He]2s2 electron configuration. The next element down, magnesium, is expected to have exactly the same arrangement of electrons in the n = 3 principal shell: [Ne]3s2. By extrapolation, we expect all the group 2 elements to have a ns2 valence electron configuration.
Exercise
Use the periodic table to predict the characteristic valence electron configuration of the halogens in group 17.
Answer: All have an ns2np5 valence electron configuration, one electron short of a noble gas electron configuration. (Note that the heavier halogens also have filled (n − 1)d10 subshells, as well as an (n − 2)f14 subshell for Rn; these do not, however, affect their chemistry in any significant way.)
Summary
For chemical purposes, the most important electrons are those in the outermost principal shell, the valence electrons. The arrangement of atoms in the periodic table results in blocks corresponding to filling of the ns, np, nd, and nf orbitals to produce the distinctive chemical properties of the elements in the s block, p block, d block, and f block, respectively.
Key Takeaway
- The arrangement of atoms in the periodic table arises from the lowest energy arrangement of electrons in the valence shell.