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# 8.2: Families and Periods of the Periodic Table

Difficulty Level: At Grade Created by: CK-12

## Lesson Objectives

The student will:

• identify groups in the periodic table.
• state the number of valence electrons for each A group in the periodic table.
• explain the relationship between the chemical behavior of families in the periodic table and their electron configuration.
• identify periods in the periodic table.
• describe the similarities among elements in the same period in the periodic table.

## Vocabulary

• actinide series
• alkali metals
• alkaline earth metals
• group (family)
• halogens
• lanthanide series
• noble gases
• period
• transition elements

## Introduction

When Mendeleev created his periodic table, he did not know why certain elements had similar chemistry. He placed the elements in their positions because they exhibited similar chemical behaviors. Thus, the vertical columns in Mendeleev’s table were composed of elements with similar chemistry. These vertical columns are called groups, or families. In this section, you are going to see that the elements in the same groups are related to each other by their electron configurations. Since the families of elements were organized by their chemical behavior, it is predictable that the individual members of each chemical family will have similar electron configurations.

If you examine a periodic table, you will often find a number written above each group (column). These numbers serve as labels, and groups are often referred to by their labels. Depending on the source or age of your periodic table, you may see two different numbering systems for referring to the families on the periodic table. In the older system, the numbers 1 – 8 and the letters A and B were used to label the groups. The newer convention is to label each group from 1 – 18 in sequential order. However, the older labeling scheme helps to provide more insight into the electron configurations of each group. As a result, in this text we will use the older labeling scheme to present each group. The periodic table below shows both numbering systems.

For an introduction to the electronic organization of the periodic table (1a, 1c, 1d, 1f), see http://www.youtube.com/watch?v=5MMWpeJ5dn4 (3:51).

## Group 1A

The electron configuration codes for the elements in Group 1A are:

lithium = $1s^22s^1 \ \ \ \ \ \text{or \ [He]}2s^1$
sodium = $1s^22s^22p^63s^1 \ \ \ \ \ \text{or \ [Ne]}3s^1$
potassium = $1s^22s^22p^63s^23p^64s^1 \ \ \ \ \ \text{or \ [Ar]}4s^1$
rubidium = $1s^22s^22p^63s^23p^64s^23d^{10}4p^65s^1 \ \ \ \ \ \ \text{or \ [Kr]}5s^1$
cesium = $1s^22s^22p^63s^23p^64s^23d^{10}4p^65s^24d^{10}5p^66s^1 \ \ \ \ \ \text{or \ [Xe]}6s^1$
francium = $1s^22s^22p^63s^23p^64s^23d^{10}4p^65s^24d^{10}5p^66s^24f^{14}5d^{10}6p^67s^1 \ \ \ \ \ \text{or \ [Rn]}7s^1$

The fact that all 1A elements participate in similar chemistry despite having vastly different nuclear sizes further illustrates the fact that electrons, particularly valence electrons, are the primary contributors to chemical reactivity. The electron configuration for the outermost energy levels of the 1A elements is the same, with the only difference being the energy level involved. Each larger member of the family has its single s electron in the next larger principal energy level. As the atomic sizes in this family increase, the valence electrons are located further from the nucleus and are therefore easier to lose. Lithium reacts readily with water, sodium reacts violently with water, potassium reacts so violently that the hydrogen formed begins to burn, and rubidium explodes instantly in water. The differences in these chemical reactions are just a matter of degree. Each larger atom reacts in the same way as the one before, but it reacts faster. The speed and ease of reaction of elements is referred to as reactivity. The relative reactivity of the elements in the 1A family (also called alkali metals) increases as the atoms become larger. Note that in the new labeling scheme, this group is labeled as Group 1.

All of the 1A elements have one valence electron and react in similar ways as the other members of the family – that’s why they are a part of the same family. The 1A metals have similar electron configurations, which causes them to have similar chemistry. Although Mendeleev did not know about these elements had similar electron configurations, he recognized the similar chemistry and organized the periodic table based on this similar chemistry.

## Group 2A

Here are the electron configuration codes for the first five elements in Group 2A:

beryllium = $1s^22s^2 \ \ \ \ \ \text{or \ [He]}2s^2$
magnesium = $1s^22s^22p^63s^2 \ \ \ \ \ \text{or \ [Ne]}3s^2$
calcium = $1s^22s^22p^63s^23p^64s^2 \ \ \ \ \ \text{or \ [Ar]}4s^2$
strontium = $1s^22s^22p^63s^23p^64s^23d^{10}4p^65s^2 \ \ \ \ \ \text{or \ [Kr]}5s^2$
barium = $1s^22s^22p^63s^23p^64s^23d^{10}4p^65s^24d^{10}5p^66s^2 \ \ \ \ \ \text{or \ [Xe]}6s^2$

All of the elements in this family have two valence electrons and have very similar chemistry. This group of metals is called the alkali earth metals. As with the 1A family, the elements in this family also increase in reactivity as the elements become larger and the valence electrons are held more loosely. In the new labeling conventions, this group is labeled as Group 2.

## Group 3A

The electron configurations of Group 3A (Group 3 according to new labeling conventions) show that all the members of this family have three valence electrons. The chemical behaviors of the elements in this family are similar to each other, but they are not as consistent from element to element as they are for other families of elements. Information in the following sections will explain why.

boron = $1s^22s^22p^1 \ \ \ \ \ \text{or \ [He]}2s^22p^1$
aluminum = $1s^22s^22p^63s^23p^1 \ \ \ \ \ \text{or \ [Ne]}3s^23p^1$
gallium = $1s^22s^22p^63s^23p^64s^23d^{10}4p^1 \ \ \ \ \ \text{or \ [Ar]}4s^23d^{10}4p^1$
indium = $1s^22s^22p^63s^23p^64s^23d^{10}4p^65s^24d^{10}5p^1 \ \ \ \ \ \text{or \ [Kr]}5s^24d^{10}5p^1$
thallium = $1s^22s^22p^63s^23p^64s^23d^{10}4p^65s^24d^{10}5p^66s^24f^{14}5d^{10}6p^1 \ \ \ \ \ \text{or \ [Xe]}6s^24f^{14}5d^{10}6p^1$

## Groups 4A through 8A

Group 4A (also known as Group 14) members have four valence electrons.

carbon = $\text{[He]}2s^22p^2$
silicon = $\text{[Ne]}3s^23p^2$
germanium = $\text{[Ar]}4s^23d^{10}4p^2$
tin = $\text{[Kr]}5s^24d^{10}5p^2$
lead = $\text{[Xe]}6s^24f^{14}5d^{10}6p^2$

Group 5A (or Group 15) members have five valence electrons.

nitrogen = $\text{[He]}2s^22p^3$
phosphorus = $\text{[Ne]}3s^23p^3$
arsenic = $\text{[Ar]}4s^23d^{10}4p^3$
antimony = $\text{[Kr]}5s^24d^{10}5p^3$
bismuth = $\text{[Xe]}6s^24f^{14}5d^{10}6p^3$

Instead of showing you the electron configurations for Groups 6A (or Group 16), 7A (or 17), and 8A (or 18), it would be good practice for you to practice writing the electron configurations for these families. When you finish writing them, check to make sure that the outer energy levels contain six valence electrons, seven valence electrons, and eight valence electrons for Groups 6A, 7A, and 8A, respectively. The group label provides a hint about the valence electrons: the number preceding the “A” is equal to the number of valence electrons for each atom in the group.

Groups 7A and 8A are also known by the names halogens and the noble gases, respectively. Groups 7A and 8A were two of the first families identified because the chemistry of their members are so similar to each other.

## Transition Elements

If you have been checking where the different groups that have been introduced so far are located on the periodic table, you may have noticed that there are a series of elements that are not part of Groups 1A – 8A (sometimes refer to as the main group). In some periodic tables, these groups of elements are in families called B groups, and in newer periodic tables, these groups are numbered 3 – 12.

Consider the fourth horizontal row in the periodic table. For potassium, the first element in this row, one electron resides in the 4s orbital. For calcium, a second electron is added to the 4s orbital. Beyond calcium, however, the pattern in which the electrons are added changes. Beginning with scandium, atomic number 21, the additional electrons do not enter the valence shell but instead enter the d sub-level of the third energy level (n = 3), as illustrated below. The electron configuration for scandium is $\text{[Ar]}4s^23d^1$.

The ten elements formed by filling in the 3d orbitals, as well as all other elements that have between 1 to 10 electrons in d orbitals, are called the transition elements. These elements, in general, differ from each other in the electron structure of the next-to-last energy level. For the most part, these elements have similar outer energy levels because they have two valence electrons. As a result, they have somewhat similar chemistry, even though they are not in the same vertical family.

## The Lanthanide and Actinide Series

There is still one more block of elements on the periodic table that has not been introduced. This block is usually placed below the periodic table and represents elements with electrons in the f orbitals. In the sixth horizontal row, the first electron goes into the 6s orbital and produces the element cesium. The second electron fills the 6s, orbital producing the element barium. The next 14 electrons that are added then enter the 4f orbitals (marked in red in the figure below).

This group of elements, atomic numbers 57 – 70, is called the lanthanide series. Elements with atomic numbers 89 – 102 are called the actinide series. As in the case of the transition elements, these elements have the electrons added to an inner energy level, rather than the valence shell. The number of valence electrons in these elements remains at 2 while the electrons enter and eventually fill the f orbitals. For some of the elements with more complicated electron configuration, this is a somewhat simplified way to generate the electron configurations. Due to the fact that outer orbitals, such as 5f and 6d, have increasingly similar energies, other factors can come into play and change the exact configuration of the lanthanides, actinides, and some transition metals.

## Sub-Level Filling Blocks on the Periodic Table

For Groups 1A and 2A, the last electron added enters an s orbital. For this reason, the 1A and 2A families make up an area on the periodic table known as the “s block.” In Groups 3A, 4A, 5A, 6A, 7A, and 8A, the last electron enters a p orbital. These six families of elements, therefore, make up an area called the “p block.” The various sub-level blocks, including the “d block” and “f block,” are indicated in the periodic table shown below.

Even though the periodic table was organized according to the chemical behavior of the elements, you can now see that the shape and design of the table is a perfect reflection of the electron configuration of the atoms. This is because the chemical behavior of the elements is also dictated by the electron configuration of the atoms.

## Periods of the Periodic Table

In addition to providing information about the electron configuration, an element’s position on the periodic table can also be used to figure out the energy level of the element’s valence electrons. Let’s try to figure out what we can learn from an element’s period in the periodic table. A period is a horizontal row of elements on the periodic table. The figure below shows how the different rows in the periodic table are numbered.

To understand what this means in terms of an element’s electron configuration, let’s consider the Group 1A metals. If we write the electron configuration for the Group 1A metal from each row of the Periodic Table, we have:

$\begin{array}{c|c|l} \mathbf{row} & \mathbf{element} & \mathbf{electron \ configuration} \\\hline\mathbf{2}&\text{Li}&1s^2\mathbf{2}s^1\\\mathbf{3}&\text{Na}&1s^22s^22p^6\mathbf{3}s^1\\\mathbf{4}&\text{K}&1s^22s^22p^63s^23p^6\mathbf{4}s^1\\\mathbf{5}&\text{Rb}&1s^22s^22p^63s^23p^64s^23d^{10}4p^6\mathbf{5}s^1\\\mathbf{6}&\text{Cs}&1s^22s^22p^63s^23p^64s^23d^{10}4p^65s^24d^{10}5p^6\mathbf{6}s^1\\\mathbf{7}&\text{Fr}&1s^22s^22p^63s^23p^64s^23d^{10}4p^65s^24d^{10}5p^66s^24f^{14}5d^{10}6p^6\mathbf{7}s^1\\\end{array}$

Do you see any patterns? For Group 1A metals, it seems that element’s row corresponds to the energy level of that element’s valence electron. Lithium (Li), for instance, is found in row 2 of the periodic table, and its valence electron is found in the n = 2 energy level. Cesium (Cs) is found in row 6 of the periodic table, and its valence electron is found in the n = 6 energy level. Now consider the Group 2A metals:

$\begin{array}{c|c|l} \mathbf{row} & \mathbf{element} & \mathbf{electron \ configuration} \\\hline\mathbf{2}&\text{Be}&1s^2\mathbf{2}s^2\\\mathbf{3}&\text{Mg}&1s^22s^22p^6\mathbf{3}s^2\\\mathbf{4}&\text{Ca}&1s^22s^22p^63s^23p^6\mathbf{4}s^2\\\mathbf{5}&\text{Sr}&1s^22s^22p^63s^23p^64s^23d^{10}4p^6\mathbf{5}s^2\\\mathbf{6}&\text{Ba}&1s^22s^22p^63s^23p^64s^23d^{10}4p^65s^24d^{10}5p^6\mathbf{6}s^2\\\mathbf{7}&\text{Ra}&1s^22s^22p^63s^23p^64s^23d^{10}4p^65s^24d^{10}5p^66s^24f^{14}5d^{10}6p^6\mathbf{7}s^2\\\end{array}$

For Group 2A metals, the same rule applies. Magnesium (Mg) is found in row 3 of the periodic table, and its valence electrons are found in the n = 3 energy level. Similarly, radium (Ra) is found in row 7 of the periodic table, and its valence electrons are found in the n = 7 energy level.

Consider the electron configurations for the Group 3A elements in each row:

$\begin{array}{c|c|l} \mathbf{row} & \mathbf{element} & \mathbf{electron \ configuration} \\\hline\mathbf{2}&\text{B}&1s^2\mathbf{2}s^2\mathbf{2}p^1\\\mathbf{3}&\text{Al}&1s^22s^22p^6\mathbf{3}s^2\mathbf{3}p^1\\\mathbf{4}&\text{Ga}&1s^22s^22p^63s^23p^6\mathbf{4}s^23d^{10}\mathbf{4}p^1\\\mathbf{5}&\text{In}&1s^22s^22p^63s^23p^64s^23d^{10}4p^6\mathbf{5}s^24d^{10}\mathbf{5}p^1\\\mathbf{6}&\text{Tl}&1s^22s^22p^63s^23p^64s^23d^{10}4p^65s^24d^{10}5p^6\mathbf{6}s^25d^{10}\mathbf{6}p^1\\\end{array}$

Even though the valence electrons in Group 3A elements are found in both s and p orbitals, it turns out that an element’s row still corresponds to the energy level of that element’s valence electrons. For example, gallium (Ga) is found in row 4 of the periodic table, and its valence electrons are found in the n = 4 energy level. Likewise, thallium (Tl) is found in row 6 of the periodic table, and its valence electrons are found in the n = 6 energy level.

Let's consider one last example by examining the electron configurations for the Group 7A elements:

$\begin{array}{c|c|l} \mathbf{row} & \mathbf{element} & \mathbf{electron \ configuration} \\\hline\mathbf{2}&\text{F}&1s^2\mathbf{2}s^2\mathbf{2}p^5\\\mathbf{3}&\text{Cl}&1s^22s^22p^6\mathbf{3}s^2\mathbf{3}p^5\\\mathbf{4}&\text{Br}&1s^22s^22p^63s^23p^6\mathbf{4}s^23d^{10}\mathbf{4}p^5\\\mathbf{5}&\text{I}&1s^22s^22p^63s^23p^64s^23d^{10}4p^6\mathbf{5}s^24d^{10}\mathbf{5}p^5\\\mathbf{6}&\text{At}&1s^22s^22p^63s^23p^64s^23d^{10}4p^65s^24d^{10}5p^6\mathbf{6}s^25d^{10}\mathbf{6}p^5\\\end{array}$

Once again, it turns out that an element’s period can be used to determine the energy level of that element’s valence electrons. Chlorine (Cl), for instance, is found in row 3 of the periodic table, and its valence electrons are found in the n = 3 energy level. Similarly, iodine (I) is found in row 5 of the periodic table, and its valence electrons are found in the n = 5 energy level.

Note, however, that determining the energy level of an element’s valence electrons by using the element’s row number works only for Groups 1A – 8A. This rule does not apply to Groups 1B – 8B (also known as Groups 3 – 12) or to lanthanides and actinides. The elements in these groups behave differently and do not obey the same rules as the Group 1A – 8A elements in terms of valence electrons and valence electron energy levels.

You may have noticed that the first row has not been mentioned at all. In all of the examples considered, only the elements starting from the second row were considered. The first row in the periodic table is a “special row” for several reasons. To begin with, the first row only contains hydrogen and helium. Can you figure out why there are only two elements in the first row?

Recall that an element’s period number corresponds to the energy level of that element’s valence electrons. Therefore, the first row must contain elements with valence electrons in the n = 1 energy level. In the n = 1 energy level, there is only one orbital: the 1s orbital. Hydrogen (electron configuration 1s1) has one valence electron in the 1s orbital, and helium (electron configuration 1s2) has two valence electrons in the 1s orbital. Since it is impossible to fit more than two electrons into the 1s orbital, atoms with a total of three or more electrons must have valence electrons in an energy level with n = 2 or greater. Atoms with a total of three or more electrons do not belong in the first row of the periodic table.

The first row is also special because its elements have special properties. Hydrogen, for example, is not a metal like the rest of the Group 1A elements. Instead, hydrogen atoms react with each other and form hydrogen gas, H2. Hydrogen is different from the other Group 1A elements, so the chemical properties of hydrogen cannot be compared with the chemical properties of the other elements in Group 1A.

The second element in the first row, helium, is also a special atom. Helium is noticeably placed away from hydrogen at the far right-hand side of the periodic table. Helium’s placement reflects its special chemical properties. While helium only has two valence electrons when the rest of the Group 8A elements have eight, helium is grouped with the 8A family because helium’s chemical behavior is similar to the chemical behaviors of the other noble gases. This is because helium and the other noble gases all have a completed outer energy shell, even though the number of valence electrons is different.

This video describes the organization and family properties within the periodic table (1b, 1c, 1f): http://www.youtube.com/watch?v=rD0maRGQhU8 (4:45).

## Lesson Summary

• The vertical columns on the periodic table are called groups or families because the elements in the column exhibit similar chemical behaviors.
• The groups on the periodic table also have similar electron configurations.
• The electron configuration in the outer energy level of the atoms in a family is the same, with the only difference in the electron configurations being the number of the outer energy level.
• All the members of a family of elements have the same number of valence electrons.
• The blocks of elements in the periodic table correspond to the sub-level in which the last electrons are being added.
• The horizontal rows on the periodic table are called periods.
• You can predict the energy level of an element’s valence electrons using the element’s period number in the periodic table.
• For the elements in Groups 1A, 2A, 3A, 4A, 5A, 6A, 7A, and 8A, the period in which an element is found corresponds to the energy level of that element’s valence electrons.

This website provides a dynamic periodic table that also provides information about the orbitals and electron configurations of the elements.

This website reviews the different groups in the periodic table.

This video is a ChemStudy film called “Transuranium Elements.” The film is somewhat dated but the information is accurate.

This video is a ChemStudy film called “Inert Gas Compounds.” The film is somewhat dated but the information is accurate.

## Review Questions

1. The horizontal rows in the periodic table are called ______________.
2. The vertical columns in the periodic table are called ______________.
3. Elements in a ______________ have similar chemical properties.
1. period
2. family
3. both a and b
4. neither a nor b
4. Which of the following elements is in the same family as fluorine?
1. silicon
2. antimony
3. iodine
4. arsenic
5. None of these.
5. Which of the following is a noble gas?
1. polonium
2. astatine
3. krypton
4. hydrogen
5. francium
6. Name the period (horizontal row) and the sub-level (s, p, d, f) of the valence electrons for the following elements. Then indicate the number of valence electrons for that element.
1. hydrogen
2. carbon
3. sulfur
4. barium
5. iodine
7. Of the s-block, p-block, d-block, and f-block of elements, which block contains most of the nonmetals?
8. Why are there no transition elements in periods 1 and 2?
9. Why do many of the transition elements have similar chemistry?
10. Use the periodic table to determine the energy level and sub-level of the highest energy electrons in each of the following elements:
1. N
2. Ca
3. Rb
4. P
5. In
11. Decide whether each of the following statements is true or false.
1. Li has valence electrons in the n = 1 energy level.
2. Si has valence electrons in the n = 3 energy level.
3. Ga has valence electrons in the n = 3 energy level.
4. Xe has valence electrons in the n = 5 energy level.
5. P has valence electrons in the n = 2 energy level.

All images, unless otherwise stated, are created by the CK-12 Foundation and are under the Creative Commons license CC-BY-NC-SA.

Feb 23, 2012

Jan 07, 2015