- 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 on the periodic table: alkali metals, alkaline earth metals, halogens, and noble gases.
- Locate the transition elements, lanthanides, and actinides on the periodic table.
- alkali metal
- alkaline earth metal
- inner transition element
- noble gas
- representative (main-group) elements
- transition element
Check Your Understanding
Recalling Prior Knowledge
- What is the difference between a principal energy level and a sublevel?
- How many orbitals are always present in each of the types of sublevels: s, p, d, f?
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 of the periodic table is related to electron configurations, which in turn influences chemical reactivity.
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 below (Table below).
Period Length and Sublevels in the Periodic Table
Number of Elements in Period
Sublevels in Order of Filling
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]3d84s2. 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. The Figure below shows a version of the periodic table that includes abbreviated electron configurations.
This periodic table shows the outer electron configurations of the elements.
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 (Figure below).
A block diagram of the periodic table shows which sublevels are being filled at any point.
The Figure above also illustrates how the d sublevel is always one principal level behind the period in which that sublevel occurs. In other words, the 3d sublevel fills during the fourth period. The f sublevel is always two levels behind. The 4f sublevel belongs to the sixth period.
We will now examine each of these blocks 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 the second period. The s-block elements and the p-block elements are together called the representative or main-group elements.
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 outermost principal energy. Recall that such electrons are called valence electrons. 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 1s22s1. This single valence electron is what gives the alkali metals their extreme reactivity. Pictured below (Figure below) is the element, sodium.
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 (Figure below). 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.
The alkaline earth metals include beryllium, magnesium, calcium, strontium, and barium. Strontium and barium react with air and must be stored in oil.
Watch videos of experiments involving the s-block elements:
Hydrogen and Helium
Looking at the block diagram (Figure above), 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.
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 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 ns2np1-6. For example, the electron configuration of elements in Group 13 is ns2np1, the configuration of elements in Group 15 is ns2np3, 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 ns2np6 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 3s23p4. In the older system of labeling groups, the representative elements are designated 1A through 8A. 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 6A, so it has 6 valence electrons.
The properties of the p-block elements vary widely. The line separating metals from nonmetals runs through the p-block. As a result, this block includes 8 metals, all 7 metalloids, and all nonmetals except for hydrogen. Note that there is some variation among different periodic tables over how to classify the rare elements polonium and astatine. The metals of the p-block are much more stable than the s-block metals. Aluminum and tin are frequently used in packaging, and lead (Figure below) is used in car batteries, bullets, and radiation shields.
Lead blocks are used in radiation shielding.
The elements of Group 17 (fluorine, chlorine, bromine, iodine, and astatine) are called the halogens. The halogens all have the general electron configuration ns2np5, 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, bromine is a dark orange liquid, and iodine is a dark purple-gray solid. Astatine is so rare that its properties are mostly unknown.
Watch videos of experiments involving the p-block elements:
Transition elements are the elements that are found in Groups 3-12 on the periodic table. The d sublevel, which is in the process of being filled, 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]3d14s2. 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 below). 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.
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 below).
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 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. Since there are five d orbitals, each of which can accommodate 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)d1-10ns2. 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]4d25s2. 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. Recall that there are several deviations from the expected order of filling the d sublevel that cannot always be easily understood. The element cobalt (Co) is in Period 4 and Group 9. It has the expected electron configuration of [Ar]3d74s2. Directly below cobalt is the element rhodium (Rh). However, its configuration is [Kr]4d85s1, meaning that one of its 5s electrons has moved to the 4d sublevel. The total of nine electrons still allows you to predict that rhodium is a member of Group 9.
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. Rhodium is an example of a transition metal with only one valence electron, because its configuration deviates from the expected filling order.
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)f1-14ns2. 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. The lanthanides and actinides together are sometimes called the inner transition elements.
Sample Problem 6.1: Electron Configurations and the Periodic Table
The electron configurations for atoms of four different elements are shown below. Without consulting the periodic table, name the period, group, and block in which each element is located. Determine the number of valence electrons for each. Then, using a periodic table, name the element and identify it as a metal, nonmetal, or metalloid.
Step 1: Plan the problem.
The period is the highest occupied principal energy level. The group is the vertical column. The block depends on which sublevel is in the process of being filled. The valence electrons are those in the outermost principal energy level. For name and type of element, use a periodic table.
Step 2: Solutions.
- The highest occupied principal energy level is the fifth, so this element is in Period 5. By adding 10 + 2 + 3 from the given configuration, we can see that the element is in Group 15. Since the p sublevel is partially filled, it is in the p-block. There are five electrons in the outermost energy level, so it has 5 valence electrons. The element is antimony, a metalloid.
- The element is in Period 7. Since the f sublevel is partially filled, it is part of the f-block, which means it does not belong to a group. It has 2 valence electrons. The element is americium, a metal from the actinides.
- The element is in Period 4 and Group 2. Even though the 4s sublevel is filled, the last electron went into that sublevel, making it a member of the s-block. It has 2 valence electrons. The element is calcium, a metal.
- The element is in Period 6. In determining the group, the f electrons can be ignored since they do not affect groups. By adding 6 + 2 = 8, we can place the element in Group 8. The partially filled d sublevel makes it a member of the d-block. It has 2 valence electrons. The element is osmium, a metal (specifically, a transition metal).
Step 3: Think about your result.
Once you have identified each element, you can use the electron configurations from a periodic table to make sure they have been identified correctly.
- For each of the following, identify the period, group, and block. Determine the number of valence electrons. Identify the element by name and type (metal, nonmetal, or metalloid).
- For the two elements in problem 1 above, how many unpaired electrons are in the atom?
- Which two elements have three unpaired electrons in the 3d sublevel?
You can watch video lectures on this topic from Kahn Academy. These videos combine orbital configurations (from the chapter Electrons in Atoms) with locations on the periodic table:
- An element’s placement in the periodic table is determined by its electron configuration.
- The chemical properties of various elements are intimately related to their valence electron configurations.
- The periodic table is divided into four blocks (s, p, d, 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.
Lesson Review Problems
- How many elements are in the second period? The fourth? The sixth?
- Use a periodic table to identify the block in which each of the following elements would be found.
- Give the common name for each of the following groups.
- Group 17
- Group 2
- Group 1
- Group 18
- Where on the periodic table are the most reactive metals?
- Where on the periodic table are the most reactive nonmetals?
- Why are the noble gases almost completely unreactive?
- Which groups are called the representative elements?
- What common name is given to the set of elements located in the d-block?
- Without referring to the periodic table, write the full electron configuration of the following elements.
- the element in Period 2 and Group 15
- the alkaline earth metal in Period 3
- the noble gas in Period 3
- Using a periodic table, identify the following elements by name.
- the halogen in Period 5
- the alkali metal in Period 4
- the element in Period 4 that has two electrons in the p sublevel
- the second period element with six valence electrons
- the lanthanide with three f sublevel electrons
- the element in Period 6 with five half-filled d orbitals
- Which element has five occupied principal energy levels and chemical properties that are similar to those of gallium?
- Without looking at the periodic table, identify the period, group, and block for elements that have the electron configurations below. Determine the number of valence electrons. Using the periodic table, identify the element by name and type (metal, nonmetal, or metalloid).
- Which two elements have one unpaired electron in the 4p sublevel?
Points to Consider
So far you have learned that the periodic table is based on both electronic structure and similarities in chemical reactivity. Next we will consider how some important physical properties of elements show distinct trends within periods and groups.
- What determines the size of an atom?
- If a valence electron were removed from an atom, what charge would the resulting particle have?