Electron Configuration and the
Periodic Table
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C HAPTER
Chapter 1. Electron Configuration and the Periodic Table
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Electron Configuration and
the Periodic Table
Lesson Objectives
• Distinguish between core and valence electrons
• Understand the relationship between the number of orbitals in various energy sublevels and the length of the
periods in the periodic table.
• Identify each block of the periodic table and be able to determine which block each element belongs to based
on its electron configuration.
• Describe the relationship between outer electron configuration and group number. Be able to determine the
number of valence electrons for any element.
• Locate the following groups of elements on the periodic table: alkali metals, alkaline earth metals, halogens,
noble gases, transition elements, lanthanides, and actinides.
Lesson Vocabulary
• valence electrons: The electrons that are in the highest occupied principal energy level (n).
• core electrons: The electrons that are closer to the nucleus and less available for interaction with other atoms.
• representative (main-group) elements: elements that have the s and p sublevels for a given principal energy
level.
• alkali metals: The elements in Group 1 (lithium, sodium, potassium, rubidium, cesium, and francium).
• alkaline earth metals: The elements in Group 2 (beryllium, magnesium, calcium, strontium, barium, and
radium).
• noble gases: The elements of Group 18 (helium, neon, argon, krypton, xenon, and radon).
• halogens: The elements of Group 17 (fluorine, chlorine, bromine, iodine, and astatine).
• transition elements: The elements that are found in Groups 3-12 on the periodic table.
• lanthanides: The 14 elements from cerium (atomic number 58) to lutetium (atomic number 71).
• actinides: The 14 elements from thorium (atomic number 90) to lawrencium (atomic number 103).
Check Your Understanding
• How to atoms form chemical bonds with one another? Are some elements more chemically reactive than
others?
Introduction
The development of the periodic table was largely based on elements that display similar chemical behavior. In the
modern table, these elements are found in vertical columns called groups. In this lesson, you will see how the form
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of the periodic table is related to electron configurations, which in turn influences chemical reactivity. We will first
start with the following introductory video: http://www.youtube.com/watch?v=5MMWpeJ5dn4 (3:51).
MEDIA
Click image to the left for more content.
Periods and Blocks
There are seven horizontal rows, or periods, on the periodic table. The length of each period is determined by
electron capacity of the sublevels that fill during that period, as seen in Table 1.1.
TABLE 1.1: Period Length and Sublevels in the Periodic Table
Period
1
2
3
4
5
6
7
Number of Elements in Period
2
8
8
18
18
32
32
Sublevels in Order of Filling
1s
2s 2p
3s 3p
4s 3d 4p
5s 4d 5p
6s 4f 5d 6p
7s 5f 6d 7p
Recall that the four different sublevels (s, p, d, and f) each consist of a different number of orbitals. The s sublevel
has one orbital, the p sublevel has three orbitals, the d sublevel has five orbitals, and the f sublevel has seven orbitals.
In the first period, only the 1s sublevel is being filled. Since all orbitals can hold two electrons, the entire first period
consists of just two elements. In the second period, the 2s sublevel, with two electrons, and the 2p sublevel, with
six electrons, are being filled. Consequently, the second period contains eight elements. The third period is similar
to the second, except the 3s and 3p sublevels are being filled. Because the 3d sublevel does not fill until after the 4s
sublevel, the fourth period contains 18 elements, due to the 10 additional electrons that can be accommodated by the
3d orbitals. The fifth period is similar to the fourth. After the 6s sublevel fills, the 4f sublevel is populated with up
to 14 electrons. This is followed by the 5d and the 6p sublevels. The total number of elements in the sixth period is
32. The seventh period also contains 32 elements, most of which are too unstable to be found in nature. All 32 have
been detected or synthesized, although for some of the later elements in this period, only a handful of atoms have
ever been made.
The period to which a given element belongs can easily be determined from its electron configuration. As an
example, consider the element nickel (Ni). Its electron configuration is [Ar]3d8 4s2 . The highest occupied principal
energy level is the fourth, as indicated by the 4 in the 4s2 portion of the configuration. Therefore, nickel can be found
in the fourth period of the periodic table. Figure 1.1 shows a version of the periodic table that includes abbreviated
electron configurations.
Based on electron configurations, the periodic table can be divided into blocks denoting which sublevel is in the
process of being filled. The s, p, d, and f blocks are illustrated below in Figure 1.2.
Figure 1.2 also illustrates how the d sublevel is always one principal level behind the period in which that sublevel
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Chapter 1. Electron Configuration and the Periodic Table
FIGURE 1.1
This periodic table shows the outer electron configurations of the elements.
occurs. In other words, the 3d sublevels fills during the fourth period. The f sublevel is always two levels behind.
The 4f sublevel belongs to the sixth period.
Numbering the Periodic Groups
The vertical columns, or groups, of the periodic table contain elements that exhibit similar properties. Two different
ways of numbering the groups are commonly in use. The currently preferred convention is to number each column
of the periodic table from 1-18. Group 1 includes hydrogen, lithium and sodium, and group 18 includes helium,
neon, argon, and krypton. An older method is to skip the d and f blocks and utilize Roman numerals from IA to
VIIIA. The letter A differentiates these groups from the d block groups, which are numbered using the letter B (from
IB to VIIIB). For example, the element carbon could be described as being part of group 14 or group IVA, while
scandium is in group 3 or group IIIB.
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FIGURE 1.2
A block diagram of the periodic table shows which sublevels are being filled at any point.
Valence Electrons
Because they are held more loosely to the nucleus than the inner electrons, it is the outermost electrons that dictate
the chemical behavior of a given element. Specifically, much can be predicted about the chemical reactivity of a
given element based solely on the number of electrons in its highest occupied principal energy level (n). These
electrons are referred to as valence electrons. The remaining electrons, which are closer to the nucleus and less
available for interaction with other atoms, are referred to as core electrons.
Consider the element magnesium, which has 12 electrons in a configuration of 1s2 2s2 2p6 3s2 . The highest occupied
principal energy level is 3, so all electrons with a quantum number of n = 3 are valence electrons. Thus, magnesium
has two valence electrons. The other 10 electrons (in the n = 1 and n = 2 levels) are its core electrons.
Representative Elements
We will now examine each block of the periodic table in more detail. The s and p sublevels for a given principal
energy level are filled during the correspondingly numbered period. For example, the 2s and 2p sublevels fill during
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Chapter 1. Electron Configuration and the Periodic Table
the second period. The s-block elements and the p-block elements are together called the representative or maingroup elements.
The s-block
The s-block consists of the elements in Group 1 and Group 2, which are primarily composed of highly reactive
metals. The elements in Group 1 (lithium, sodium, potassium, rubidium, cesium, and francium) are called the alkali
metals. All of the alkali metals have a single s electron in their valence energy level. The general form for the
electron configuration of each alkali metal is ns1 , where the n refers to the highest occupied principal energy level.
For example, the electron configuration of lithium (Li), the alkali metal of Period 2, is 1s2 2s1 . This single valence
electron is what gives the alkali metals their extreme reactivity. Figure 1.3 shows the alkali metal element sodium.
FIGURE 1.3
Like all alkali metals, sodium is very soft.
A fresh surface, which can be exposed by
cutting the sample, exhibits a luster that is
quickly lost as the sodium reacts with air.
All alkali metals are very soft and can be cut easily with a knife. Due to their high reactivity, they must be stored
under oil to prevent them from reacting with oxygen or water vapor in the air. The reactions between alkali metals
and water are particularly vigorous and include the rapid production of large quantities of hydrogen gas. Alkali
metals also react easily with most nonmetals. All of the alkali metals are far too reactive to be found in nature
in their pure elemental form. For example, all naturally occurring sodium exists as a compound, such as sodium
chloride (table salt).
The elements in Group 2 (beryllium, magnesium, calcium, strontium, barium, and radium) are called the alkaline
earth metals (see Figure 1.4). These elements have two valence electrons, both of which reside in the outermost s
sublevel. The general electron configuration of all alkaline earth metals is ns2 . The alkaline earth metals are still too
reactive to exist in nature as free elements, but they are less reactive than the alkali metals. They tend to be harder,
stronger, and denser than the alkali metals, and they also form numerous compounds with nonmetals.
Hydrogen and Helium
Looking at the block diagram ( Figure 1.2), you may be wondering why hydrogen and helium were not included
in our discussion of the alkali metal and alkaline earth metal groups. Though hydrogen, with its 1s1 configuration,
appears as though it should be similar to the rest of Group 1, it does not share the properties of that group. Hydrogen
is a unique element that cannot be reasonably included in any single group of the periodic table. Some periodic
tables even separate hydrogen’s square from the rest of Group 1 to indicate its solitary status.
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FIGURE 1.4
The alkaline earth metals include beryllium, magnesium, calcium, strontium, and
barium. Strontium and barium react with
air and must be stored in oil.
Helium has a configuration of 1s2 , which would seem to place it with the alkaline earth metals. However, it is instead
placed in Group 18 at the far right of the periodic table. The elements in this group, called the noble gases, are very
unreactive because their outermost s and p sublevels are completely filled. Since it is part of the first period, helium
does not have a p sublevel. Its filled 1s sublevel makes it very similar to the other members of Group 18.
The p-block
The p-block consists of the elements in groups 13-18. The p sublevel always fills after the s sublevel of a given
principal energy level. Therefore, the general electron configuration for an element in the p-block is ns2 np1−6 . For
example, the electron configuration of elements in Group 13 is ns2 np1 , the configuration of elements in Group 15 is
ns2 np3 , and so on. The elements of Group 18 (helium, neon, argon, krypton, xenon, and radon) are called the noble
gases. They are an especially important group of the periodic table because they are almost completely unreactive,
due to their completely filled outermost s and p sublevels. As noted above, helium might at first seem to be out of
place, because it has a configuration of 1s2 instead of the ns2 np6 configuration that is characteristic of the other noble
gases. However, because there are no 1p orbitals, helium also has a completely filled outermost energy level, which
leads to the various chemical properties exhibited by the other noble gases.
Note that the noble gases were not a part of Mendeleev’s periodic table because they had not yet been discovered.
In 1894, English physicist Lord Rayleigh and Scottish chemist Sir William Ramsay detected argon as a small
percentage of the atmosphere. Discovery of the other noble gases soon followed. The group was originally called
the inert gases because they were believed to be completely unreactive and unable form compounds. However,
beginning in the early 1960s, several compounds of xenon were synthesized by treating it with highly reactive
fluorine gas. The name of the group was later changed to noble gases.
The number of valence electrons in elements of the p-block is equal to the group number minus 10. As an example,
sulfur is located in Group 16, so it has 16 –10 = 6 valence electrons. Since sulfur is located in period 3, its outer
electron configuration is 3s2 3p4 . In the older system of labeling groups, the representative elements are designated
IA through VIIIA. Using this system, the number of valence electrons is equal to the number preceding the letter A.
Using the same example, sulfur is a member of Group VIA, so it has 6 valence electrons.
The elements of Group 17 (fluorine, chlorine, bromine, iodine, and astatine) are called the halogens. The halogens
all have the general electron configuration ns2 np5 , giving them seven valence electrons. They are one electron short
of having full outer s and p sublevels, which makes them very reactive. They undergo especially vigorous reactions
with the reactive alkali metals. In their pure elemental forms, chlorine and fluorine are gases at room temperature,
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Chapter 1. Electron Configuration and the Periodic Table
bromine is a dark orange liquid, and iodine is a dark purple-gray solid. Astatine is so rare that its properties are
mostly unknown.
Transition Elements
Transition elements are the elements that are found in Groups 3-12 on the periodic table. The d sublevel, which
becomes increasingly filled from left to right across the period, is in a lower principal energy level than the s sublevel
filled before it. For example, the electron configuration of scandium, the first transition element, is [Ar]3d1 4s2 .
Remember that the configuration is not written in the same order as the sublevels are filled; the 4s sublevel gets filled
before electrons are placed into 3d orbitals. Because they are all metals, the transition elements are often called the
transition metals ( Figure 1.5). As a group, they display typical metallic properties but are less reactive than the
metals in Groups 1 and 2. Some of the more familiar transition metals are unreactive enough to be found in nature
as pure elements, such as platinum, gold, and silver.
FIGURE 1.5
Silver (left) and chromium (right) are two
typical transition metals.
Many transition elements make compounds that are distinctive for being vividly colored. Electron transitions that
occur within the d sublevel absorb some of the wavelengths present in white light, and the wavelengths that are not
absorbed are perceived by observers as the color of the compound ( Figure 1.6).
FIGURE 1.6
Transition metal compounds dissolved in
water exhibit a wide variety of bright colors. From left to right are shown solutions
of cobalt(II) nitrate, potassium dichromate, potassium chromate, nickel(II) chloride, copper(II) sulfate, and potassium
permanganate.
The d-block
The transition elements found in Groups 3-12 are also referred to as the d-block, since the d sublevel is in the process
of being filled across the d-block from left to right. Since there are five d orbitals, each of which can accommodate
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two electrons, there are ten elements in each period of the d-block. The general electron configuration for elements
in the d-block is (n - 1)d1−10 ns2 . The d sublevel being filled belongs to a principal energy level that is one lower than
the s sublevel that has just been filled. For example, the configuration of zirconium (Zr) is [Kr]4d2 5s2 . The group
number can easily be determined from the combined number of electrons in the s and d sublevels. Zirconium is in
Period 5 and Group 4.
Because electrons in the d sublevel do not belong to the outermost principal energy level, they are not valence
electrons. Most d-block elements have two valence electrons, which are the two electrons from the outermost s
sublevel.
The f-block
The first of the f sublevels is the 4f sublevel. It fills after the 6s sublevel, meaning that f sublevels are two principal
energy levels behind. The general electron configuration for elements in the f-block is (n - 2)f1−14 ns2 . The seven
orbitals of the f sublevel can each accommodate two electrons, so the f-block is 14 elements in length. It is usually
shown pulled out of the main body of the periodic table and is placed at the very bottom. Because of that, the
elements of the f-block do not belong to any of the numbered groups; they are wedged in between Groups 3 and
4. The lanthanides are the 14 elements from cerium (atomic number 58) to lutetium (atomic number 71). Most
lanthanides have a partially filled 4f sublevel. They are all metals and are similar in reactivity to the Group 2 alkaline
earth metals.
The actinides are the 14 elements from thorium (atomic number 90) to lawrencium (atomic number 103). Most
actinides have a partially filled 5f sublevel. The actinides are all radioactive elements, and only the first four have
been found to occur naturally on Earth. All of the others have only been artificially made in the laboratory.
Lesson Summary
• An element’s placement in the periodic table is determined by its electron configuration.
• Valence electrons (those in the outermost principal energy level) dictate the chemical behavior of each element.
Their relatively large distance from the nucleus makes them more available to interact with other atoms.
• Core electrons are the electrons that are closer to the nucleus and therefore do not participate in bonding.
• The periodic table is divided into 4 blocks (s, p, d, and f) based on which sublevel is in the process of being
filled.
• Alkali metals, alkaline earth metals, halogens, and noble gases are the common names of groups 1, 2, 17, and
18.
• Transition elements are members of the d-block, while the f-block consists of the lanthanides and the actinides.
Review Questions
1.
2.
3.
4.
5.
6.
7.
8.
9.
8
Sketch a periodic table, labeling the s, p, d and f blocks.
What can be said about the elements within a given group of the periodic table?
How do valence electrons and core electrons differ?
What blocks of the periodic table make up the representative elements?
Describe the relationship between the electron configuration of the alkali earth metals and their reactivity.
How do alkaline earth metals differ from the alkali metals?
Describe the properties of hydrogen and helium?
Why are the noble gases almost completely unreactive?
What are some unique properties of transition metals?
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Chapter 1. Electron Configuration and the Periodic Table
10. What block to the lanthanides and actinides belong to?
11. Use a periodic table to identify the block in which each of the following elements would be found.
a.
b.
c.
d.
rubidium
holmium
palladium
tellurium
12. Write the electron configurations for the following elements:
a. Na
b. Cl
c. Zr
Further Reading / Supplemental Links
• Winter, M. (1993-2011). WebElements: the periodic table on the WWW, from http://www.webelements.com/
Points to Consider
• Archaeological evidence suggests that people have been using tin for at least 5500 years. Tin is used to form
many useful alloys (mixtures of two or more metals). Bronze is an alloy of tin and copper, while solder is an
alloy of tin and lead.
• Gallium melts near room temperature and has one of the largest liquid ranges of any metal, so it has found use
in high temperature thermometers.
• Lead is a soft, malleable, and corrosion resistant material. The ancient Romans used lead to make water pipes,
some of which are still in use today.
• Can you think of other elements which have similar uses to those listed here?
References
1. Christopher Auyeung. CK-12 Foundation . CC BY-NC 3.0
2. Christopher Auyeung and Joy Sheng. CK-12 Foundation . CC BY-NC 3.0
3. Courtesy of the US Department of Energy. http://www.etec.energy.gov/Operations/Sodium/Sodium_Index.ht
ml . Public Domain
4. Hi-Res Images of Chemical Elements. Be: http://images-of-elements.com/beryllium.php; Mg: http://imag
es-of-elements.com/magnesium.php; Ca: http://images-of-elements.com/calcium.php; Sr: http://images-of-el
ements.com/strontium.php; Ba: http://images-of-elements.com/barium.php . CC BY 3.0
5. Courtesy of Hi-Res Images of Chemical Elements. Silver: http://images-of-elements.com/silver.php; Chromi
um: http://images-of-elements.com/chromium.php . CC BY 3.0
6. Ben Mills (User:Benjah-bmm27/Wikimedia Commons). http://commons.wikimedia.org/wiki/File:Colouredtransition-metal-solutions.jpg . Public Domain
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