- Understand the relationship between electron sublevels and the length of periods of 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.
- Identify which groups are identified by the common names: alkali metals, alkaline earth metals, halogens, and noble gases.
- Locate transition elements, lanthanides, and actinides.
- alkali metal
- alkaline earth metal
- inner transition element
- noble gas
- representative (main-group) elements
- transition element
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 of the periodic table, called periods. The length of each period is determined by the number of electrons that are capable of occupying the sublevels that fill during that period, as seen in Table below.
Period Length and Sublevels in the Periodic Table
Number of Elements in Period
Sublevels in Order of Fill
4s 3d 4p
5s 4d 5p
6s 4f 5d 6p
7s 5f 6d 7p
Recall that the four different sublevels (s, p, d, 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, filling the 3s and 3p sublevels. Notice that the 3d sublevel does not actually fill until after the 4s sublevel. This results in the fourth period containing 18 elements due to the additional 10 electrons that are contributed by the d sublevel. The fifth period is similar to the fourth. After the 6s sublevel fills, the 4f sublevel with its 14 electrons fills. This is followed by the 5d and the 6p. The total number of elements in the sixth period is 32. The later elements in the seventh period are still being created. So while there are a possible of 32 elements in the period, the current number is slightly less.
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, 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 below shows a version of the periodic table which 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 in Figure below.
A block diagram of the periodic table shows which sublevels are being filled at any point.
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 sublevels 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 are always filling during the period which corresponds to that element’s highest principal energy level. That is, the second period is where the 2s and 2p sublevels fill. 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. These groups consist 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 of the electron configuration of each of the alkali metals is ns1, where the n refers to the highest occupied principal energy level. As an 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. Figure below shows the element sodium.
Sodium, like all alkali metals, is very soft. A fresh surface, exposed from cutting, exhibits 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. Their high reactivity means that they must be stored under oil in order to prevent them from reacting with air or with water vapor. The reactions between alkali metals and water is particularly vigorous and spectacular. The reaction rapidly produces 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 “free” or uncombined states. For example, all naturally occurring sodium exists as one or many compounds of sodium 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 below). These elements have two valence electrons, both belonging to 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 are less reactive than the alkali metals. They are also harder, stronger, and denser than the alkali metals. They also make many compounds with nonmetals.
Alkaline earth metals: beryllium, magnesium, calcium, strontium, and barium. Strontium and barium react with air and must be stored in oil.
Watch video experiments of 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 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 which is not reasonably included with any other 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. This group, called the noble gases, are very unreactive because their outermost s and p sublevels are completely filled. Helium, being in Group 1, 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 except for helium. 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. This continues for the remainder of the groups. 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 owing to their completely filled outermost s and p sublevels. As noted above, helium is somewhat of an exception with a configuration of 1s2, while all of the other noble gases have configurations of ns2np6. 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 it was thought that it was impossible for any of them to react and form compounds. Beginning in the early 1960s, several compounds of xenon and highly reactive fluorine were synthesized. The name of the group was 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 and has 16 – 10 = 6 valence electrons. Since sulfur is located in period 3, its outer electron configuration is [Ne]3s23p4. In the older system of labeling groups with A and B designations, the representative elements are designated 1A through 8A. The number of valence electrons is equal to the group number preceding the A. Sulfur is a member of Group 6A.
The properties of the p block elements exhibit a wide variation. The stair-step line separating metals from nonmetals runs through the p block. As a result, there are 8 metals, all 7 metalloids, and all 15 nonmetals. 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. Lead (Figure below) is used in car batteries, bullets, and as radiation shielding.
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 the full outer s and p sublevel, which makes them very reactive. They undergo especially vigorous reactions with the reactive alkali metals. As elements, 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 video experiments of p block elements:
Transition elements are the elements that are found in Groups 3-12 on the periodic table. The term refers to the fact that 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 reversed from the fill order – the 4s filled before the 3d begins. Because they are all metals, the transition elements are often called the transition metals. As a group, they display typical metallic properties and are less reactive than the metals in Groups 1 and 2. Some of the more familiar ones are so unreactive that they can be found in nature in their free, or uncombined state (Figure below). These include platinum, gold, and silver.
Silver (left) and chromium (right) are two typical transition metals.
Compounds of many transition elements are distinctive for being widely and vividly colored. Electron transitions that occur within the d sublevel release energies that result in the emission of visible light of varied wavelengths (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 that can accommodate ten 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 is always in a principal energy level that is one lower than that of the s sublevel. 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 sublevel. Zirconium is in Period 5 and Group 4. Recall from an earlier chapter 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 in the group 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 to begin filling 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 accommodate 14 electrons, so the f block is 14 elements in length. It is pulled out of the main body of the period table and is shown at the very bottom. Because of that, the elements of the f block do not belong to a group, being wedged in between Groups 3 and 4. The lanthanides are the 14 elements from cerium (atomic number 58) to lutetium (atomic number 71). The 4f sublevel is in the process of being filled for the lanthanides. 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). The 5f sublevel is in the process of being filled. The actinides are all radioactive elements and only the first four have been found 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 of atoms of four different elements are shown below. Without consulting the periodic table, name the period, group, and block where this 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. The group is found by adding 10 + 2 + 3 from the configuration. The element is in Group 15. Since the p sublevel is not 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 incompletely filled, it is part of the f block and also 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. So 6 + 2 = 8 and the element is in Group 8. The incompletely 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 periodic table which contains the electron configurations to make sure it has 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 3 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.
- Chemical properties of elements can largely be explained by the outer electron configuration.
- The periodic table is divided into 4 blocks (s, p, d, f) based on what sublevel is in the process of being filled.
- Alkali metals, alkaline earth metals, halogens, and noble gases are common names given to several groups.
- 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 that each of these elements would be found.
- Give the common name of 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
- 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 all five d orbitals half-filled
- Which element has 5 occupied principal energy levels and chemical properties that are similar to 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?