High School Chemistry/Families on the Periodic Table

With the introduction of electron configurations, we began to get a deeper understanding of the Periodic Table. An understanding of these electron configurations will prove to be invaluable as we look at bonding and chemical reactions. The orbital representation method for representing electron configuration is shown below. The orbital representation was learned in an earlier chapter but like many of the skills you learn in chemistry, it will be used a great deal in this chapter and in several chapters later in the course.

In this lesson, we will focus on the connection between the electron configuration and the main group elements of the Periodic Table. We will need to remember the sub-level filling groups in the Periodic Table. Keep the following figure in mind. We will use it for the next two chapters.

Lesson Objectives

Describe the patterns that exist in the electron configurations for the main group elements.
Identify the columns in the Periodic Table that contain 1) the alkali metals, 2) the alkaline earth metals, 3) the halogens, and 4) the noble gases, and describe the differences between each family's electron configuration.
Given the outermost energy level electron configuration for an element, determine its family on the Periodic Table.

Alkali Metals Have One Electron in Their Outer Energy Level

Elements Ending with s¹ = Alkali

In the Periodic Table, the elements are arranged in order of increasing atomic number. In previous material we learned that the atomic number is the number of protons in the nucleus of an atom. For a neutral atom, the number of protons is equal to the number of electrons. Therefore, for neutral atoms, the Periodic Table is also arranged in order of increasing number of electrons. Take a look now at the first group or column in the Periodic Table. It is the one marked "1A" in the Period Table figure above. The groups or families are the vertical rows of elements. The first group has seven elements representing the seven periods of the Periodic Table. Remember that a period in the Periodic Table is a horizontal row. Group 1A is the only group with seven elements in it.

Table 9.1: Electron Configurations for Group 1A Metals
Element	Atomic Number	Electron Configuration
Lithium (Li)	3	$1s^{2}2s^{1}$
Sodium (Na)	11	$1s^{2}2s^{2}2p^{6}3s^{1}$
Potassium (K)	19	$1s^{2}2s^{2}2p^{6}3s^{2}3p^{6}4s^{1}$
Rubidium (Rb)	37	$1s^{2}2s^{2}2p^{6}3s^{2}3p^{6}4s^{2}3d^{10}4p^{6}5s^{1}$
Cesium (Cs)	55	$1s^{2}2s^{2}2p^{6}3s^{2}3p^{6}4s^{2}3d^{10}4p^{6}5s^{2}4d^{10}5p^{6}6s^{1}$
Francium (Fr)	87	$1s^{2}2s^{2}2p^{6}3s^{2}3p^{6}4s^{2}3d^{10}4p^{6}5s^{2}4d^{10}5p^{6}6s^{2}4f^{14}5d^{10}6p^{6}7s^{1}$

Table 9.1: Electron Configurations for Group 1A Metals
Element	Atomic Number	Electron Configuration
Lithium (Li)	3	$1s^{2}2s^{1}$
Sodium (Na)	11	$1s^{2}2s^{2}2p^{6}3s^{1}$
Potassium (K)	19	$1s^{2}2s^{2}2p^{6}3s^{2}3p^{6}4s^{1}$
Rubidium (Rb)	37	$1s^{2}2s^{2}2p^{6}3s^{2}3p^{6}4s^{2}3d^{10}4p^{6}5s^{1}$
Cesium (Cs)	55	$1s^{2}2s^{2}2p^{6}3s^{2}3p^{6}4s^{2}3d^{10}$ $4p^{6}5s^{2}4d^{10}5p^{6}6s^{1}$
Francium (Fr)	87	$1s^{2}2s^{2}2p^{6}3s^{2}3p^{6}4s^{2}3d^{10}$ $4p^{6}5s^{2}4d^{10}5p^{6}6s^{2}4f^{14}5d^{10}6p^{6}7s^{1}$

What do you notice about all of the elements in Group 1? They all have s¹ as the outermost energy level electron configuration. The whole number in front of the "s" tells you what period the element is in. For example sodium, Na, has the electron configuration 1s²2s²2p⁶3s¹, so it is in period 3. It is the first element of this period.

This group of elements is called the alkali metals. They get their name from ancient Arabic (al kali) because "scientists" of the time found that the ashes of the vegetation they were burning contained a large amount of sodium and potassium. In Arabic, al kali means ashes. We know today that all alkali metals have electronic configurations ending in s¹. You might want to note that while hydrogen is often placed in group 1, it is not considered an alkali metal. The reason for this will be discussed later.

Alkaline Earth Elements Have Two Electrons in Their Outer Energy Level

Elements Ending with s² = Alkaline Earth

Taking a look at Group 2A in Table 9.2, we can use the same analysis we used with group 1 to see if we can find a similar trend. It is the second vertical group in the Periodic Table and it contains only six elements.

Table 9.2: Electron Configurations for Group 2A Metals
Element	Atomic Number	Electron Configuration
Beryllium (Be)	4	$1s^{2}2s^{2}$
Magnesium (Mg)	12	$1s^{2}2s^{2}2p^{6}3s^{2}$
Calcium (Ca)	20	$1s^{2}2s^{2}2p^{6}3s^{2}3p^{6}4s^{2}$
Strontium (Sr)	38	$1s^{2}2s^{2}2p^{6}3s^{2}3p^{6}4s^{2}3d^{10}4p^{6}5s^{2}$
Barium (Ba)	56	$1s^{2}2s^{2}2p^{6}3s^{2}3p^{6}4s^{2}3d^{10}4p^{6}5s^{2}4d^{10}5p^{6}6s^{2}$
Radium (Ra)	88	$1s^{2}2s^{2}2p^{6}3s^{2}3p^{6}4s^{2}3d^{10}4p^{6}5s^{2}4d^{10}5p^{6}6s^{2}4f^{14}5d^{10}6p^{6}7s^{2}$

Table 9.2: Electron Configurations for Group 2A Metals
Element	Atomic Number	Electron Configuration
Beryllium (Be)	4	$1s^{2}2s^{2}$
Magnesium (Mg)	12	$1s^{2}2s^{2}2p^{6}3s^{2}$
Calcium (Ca)	20	$1s^{2}2s^{2}2p^{6}3s^{2}3p^{6}4s^{2}$
Strontium (Sr)	38	$1s^{2}2s^{2}2p^{6}3s^{2}3p^{6}4s^{2}3d^{10}4p^{6}5s^{2}$
Barium (Ba)	56	$1s^{2}2s^{2}2p^{6}3s^{2}3p^{6}4s^{2}3d^{10}$ $4p^{6}5s^{2}4d^{10}5p^{6}6s^{2}$
Radium (Ra)	88	$1s^{2}2s^{2}2p^{6}3s^{2}3p^{6}4s^{2}3d^{10}$ $4p^{6}5s^{2}4d^{10}5p^{6}6s^{2}4f^{14}5d^{10}6p^{6}7s^{2}$

What do you notice about all of the elements in group 2A? They all have an outermost energy level electron configuration of s². The whole number in front of the "s" tells you what period the element is in. For example, magnesium, Mg, has the electron configuration 1s²2s²2p⁶3s², so it is in period 3 and is the second element in that period. Remember that the s sublevel may hold two electrons, so in Group 2A, the s orbital has been filled.

Elements in this group are given the name alkaline earth metals. They get their name because early "scientists" found that all of the alkaline earths were found in the earth's crust. Alkaline earth metals, although not as reactive as the alkali metals, are still highly reactive. All alkaline earth metals have electron configurations ending in s².

Noble Gases Have 8 Electrons in Their Outer Energy Level

Elements Ending with s²p⁶ = Noble Gases

The first person to isolate a noble gas was Henry Cavendish, who isolated argon in the late 1700s. The noble gases were actually considered inert gases until the 1960s when a compound was formed between xenon and fluorine which changed the way chemists viewed the "inert" gases. In the English language, inert means to be lifeless or motionless; in the chemical world, inert means does not react. Later, the name "noble gas" replaced "inert gas" for the name of Group 8A.

When we write the electron configurations for these elements, we see the same general trend that was observed with groups 1A and 2A; that is, similar electron configurations within the group.

Table 9.3: Electron Configurations for Group 8A Gases
Element	Atomic Number	Electron Configuration
Helium (He)	2	$1s^{2}$
Neon (Ne)	10	$1s^{2}2s^{2}2p^{6}$
Argon (Ar)	18	$1s^{2}2s^{2}2p^{6}3s^{2}3p^{6}$
Krypton (Kr)	36	$1s^{2}2s^{2}2p^{6}3s^{2}3p^{6}4s^{2}3d^{10}4p^{6}$
Xenon (Xe)	54	$1s^{2}2s^{2}2p^{6}3s^{2}3p^{6}4s^{2}3d^{10}4p^{6}5s^{2}4d^{10}5p^{6}$
Radon (Rn)	86	$1s^{2}2s^{2}2p^{6}3s^{2}3p^{6}4s^{2}3d^{10}4p^{6}5s^{2}4d^{10}5p^{6}6s^{2}4f^{14}5d^{10}6p^{6}$

Table 9.3: Electron Configurations for Group 8A Gases
Element	Atomic Number	Electron Configuration
Helium (He)	2	$1s^{2}$
Neon (Ne)	10	$1s^{2}2s^{2}2p^{6}$
Argon (Ar)	28	$1s^{2}2s^{2}2p^{6}3s^{2}3p^{6}$
Krypton (Kr)	36	$1s^{2}2s^{2}2p^{6}3s^{2}3p^{6}4s^{2}3d^{10}4p^{6}$
Xenon (Xe)	54	$1s^{2}2s^{2}2p^{6}3s^{2}3p^{6}4s^{2}3d^{10}$ $4p^{6}5s^{2}4d^{10}5p^{6}$
Radon (Rn)	86	$1s^{2}2s^{2}2p^{6}3s^{2}3p^{6}4s^{2}3d^{10}$ $4p^{6}5s^{2}4d^{10}5p^{6}6s^{2}4f^{14}5d^{10}6p^{6}$

Aside from helium, He, all of the noble gases have outer energy level electron configurations that are the same, ns²np⁶, where n is the period number. So Argon, Ar, is in period 3, is a noble gas, and would therefore have an outer energy level electron configuration of 3s²3p⁶. Notice that both the s and p sublevels are filled. Helium has an electron configuration that might fit into Group 2A. However, the chemical reactivity of helium, because it has a full first energy level, is similar to that of the noble gases.

Halogens Have 7 Electrons in Their Outer Energy Level

Elements Ending with s²p⁵ = Halogens

The halogens are an interesting group. Halogens are members of Group 7A, which is also referred to as 17. It is the only group in the Periodic Table that contains all of the states of matter at room temperature. Fluorine, F₂, is a gas, as is chlorine, Cl₂. Bromine, Br₂, is a liquid and iodine, I₂, and astatine, At₂, are both solids. What else is neat about Group 7A is that it houses four (4) of the seven (7) diatomic compounds. Remember the diatomics are H₂, N₂, O₂, F₂, Cl₂, Br₂, and I₂. Notice that the latter four are Group 17 elements. The word halogen comes from the Greek meaning salt forming. French chemists discovered that the majority of halogen ions will form salts when combined with metals. We all know some of these already: LiF, NaCl, KBr, and NaI.

Taking a look at Group 7A in the figure, we can find the same pattern of similar electron configurations as found with group 1A, 2A, and 8A. It is the 17^th group in the Periodic Table and it contains only five elements.

Table 9.4: Electron Configurations for Group 7A Elements
Element	Atomic Number	Electron Configuration
Fluorine (F)	9	$1s^{2}2s^{2}2p^{5}$
Chlorine (Cl)	17	$1s^{2}2s^{2}2p^{6}3s^{2}3p^{5}$
Bromine (Br)	35	$1s^{2}2s^{2}2p^{6}3s^{2}3p^{6}4s^{2}3d^{10}4p^{5}$
Iodine (I)	53	$1s^{2}2s^{2}2p^{6}3s^{2}3p^{6}4s^{2}3d^{10}4p^{6}5s^{2}4d^{10}5p^{5}$
Astatine (At)	85	$1s^{2}2s^{2}2p^{6}3s^{2}3p^{6}4s^{2}3d^{10}4p^{6}5s^{2}4d^{10}5p^{6}6s^{2}4f^{14}5d^{10}6p^{5}$

Table 9.4: Electron Configurations for Group 7A Elements
Element	Atomic Number	Electron Configuration
Fluorine (F)	9	$1s^{2}2s^{2}2p^{5}$
Chlorine (Cl)	17	$1s^{2}2s^{2}2p^{6}3s^{2}3p^{5}$
Bromine (Br)	35	$1s^{2}2s^{2}2p^{6}3s^{2}3p^{6}4s^{2}3d^{10}4p^{5}$
Iodine (I)	53	$1s^{2}2s^{2}2p^{6}3s^{2}3p^{6}4s^{2}3d^{10}$ $4p^{6}5s^{2}4d^{10}5p^{5}$
Astatine (At)	85	$1s^{2}2s^{2}2p^{6}3s^{2}3p^{6}4s^{2}3d^{10}$ $4p^{6}5s^{2}4d^{10}5p^{6}6s^{2}4f^{14}5d^{10}6p^{5}$

What is the general trend for the elements in Group 7A? They all have, as the outermost energy level electron configuration, ns²np⁵, where n is the period number. You should also note that these elements are one group away from the noble gases (the ones that generally don't react!) and the outermost electron configuration of the halogens is one away from being filled. For example, chlorine (Cl) has the electron configuration [Ne] 3s²3p⁵ so it is in period 3, the seventh element in the main group elements. The main group elements, as you recall, are equivalent to the s + p blocks of the Periodic Table (or the pink and orange groups in the diagram above).

The Oxygen Family Has 6 Electrons in the Outer Energy Level

Elements Ending with s²p⁴ = the Oxygen Family

Oxygen and the other elements in Group 6A have a similar trend in their electron configurations. Oxygen is the only gas in the group; all others are in the solid state at room temperature. Oxygen was first named by Antoine Lavoisier in the late 1700s but really the planet has had oxygen around since plants were first on the earth.

Taking a look at Group 6A in the figure below, we find the same pattern in electron configurations that we found with the other groups. Oxygen and its family members are in the 16^th group in the Periodic Table. In Group 16, there are, again, only five elements.

Table 9.5: Electron Configurations for Group 6A Elements
Element	Atomic Number	Electron Configuration
Oxygen (O)	8	$1s^{2}2s^{2}2p^{4}$
Sulfur (S)	16	$1s^{2}2s^{2}2p^{6}3s^{2}3p^{4}$
Selenium (Se)	34	$1s^{2}2s^{2}2p^{6}3s^{2}3p^{6}4s^{2}3d^{10}4p^{4}$
Tellurium (Te)	52	$1s^{2}2s^{2}2p^{6}3s^{2}3p^{6}4s^{2}3d^{10}4p^{6}5s^{2}4d^{10}5p^{4}$
Polonium (Po)	84	$1s^{2}2s^{2}2p^{6}3s^{2}3p^{6}4s^{2}3d^{10}4p^{6}5s^{2}4d^{10}5p^{6}6s^{2}4f^{14}5d^{10}6p^{4}$

Table 9.5: Electron Configurations for Group 6A Elements
Element	Atomic Number	Electron Configuration
Oxygen (O)	8	$1s^{2}2s^{2}2p^{4}$
Sulfur (S)	16	$1s^{2}2s^{2}2p^{6}3s^{2}3p^{4}$
Selenium (Se)	34	$1s^{2}2s^{2}2p^{6}3s^{2}3p^{6}4s^{2}3d^{10}4p^{4}$
Tellurium (Te)	52	$1s^{2}2s^{2}2p^{6}3s^{2}3p^{6}4s^{2}3d^{10}$ $4p^{6}5s^{2}4d^{10}5p^{4}$
Polonium (Po)	84	$1s^{2}2s^{2}2p^{6}3s^{2}3p^{6}4s^{2}3d^{10}$ $4p^{6}5s^{2}4d^{10}5p^{6}6s^{2}4f^{14}5d^{10}6p^{4}$

When we examine the electron configurations of the Group 6A elements, we see that all of these elements have the outer energy level electron configuration of ns²np⁴. We will see that this similar electron configuration gives all elements in the group similar properties for bonding.

These elements are two groups away from the noble gases and the outermost electron configuration is two away from being filled. Sulfur, for example, has the electron configuration 1s²2s²2p⁶3s²3p⁴ so it is in period 3. Sulfur is the sixth element in the main group elements. We know it is the sixth element across the period of the main group elements because there are 6 electrons in the outermost energy level.

The Nitrogen Family Has 5 Electrons in the Outer Energy Level

Elements Ending with s²p³ = the Nitrogen Family

Just as we saw with Group 6A, Group 5A has a similar oddity in its group. Nitrogen is the only gas in the group with all other members in the solid state at room temperature. Nitrogen was first discovered by the Scottish chemist Rutherford in the late 1700s. The air is mostly made of nitrogen. Nitrogen has properties that are different in some ways from its group members. As we will learn in later lessons, the electron configuration for nitrogen provides the ability to form very strong triple bonds.

Nitrogen and its family members belong in the 15^th group in the periodic table. In Group 15, there are also only five elements.

Table 9.6: Electron Configurations for Group 5A Elements
Element	Atomic Number	Electron Configuration
Nitrogen (N)	7	$1s^{2}2s^{2}2p^{3}$
Phosphorous (P)	15	$1s^{2}2s^{2}2p^{6}3s^{2}3p^{3}$
Arsenic (As)	33	$1s^{2}2s^{2}2p^{6}3s^{2}3p^{6}4s^{2}3d^{10}4p^{3}$
Antimony (Sb)	51	$1s^{2}2s^{2}2p^{6}3s^{2}3p^{6}4s^{2}3d^{10}4p^{6}5s^{2}4d^{10}5p^{3}$
Bismuth (Bi)	83	$1s^{2}2s^{2}2p^{6}3s^{2}3p^{6}4s^{2}3d^{10}4p^{6}5s^{2}4d^{10}5p^{6}6s^{2}4f^{14}5d^{10}6p^{3}$

Table 9.6: Electron Configurations for Group 5A Elements
Element	Atomic Number	Electron Configuration
Nitrogen (N)	7	$1s^{2}2s^{2}2p^{3}$
Phosphorous (P)	15	$1s^{2}2s^{2}2p^{6}3s^{2}3p^{3}$
Arsenic (As)	33	$1s^{2}2s^{2}2p^{6}3s^{2}3p^{6}4s^{2}3d^{10}4p^{3}$
Antimony (Sb)	51	$1s^{2}2s^{2}2p^{6}3s^{2}3p^{6}4s^{2}3d^{10}$ $4p^{6}5s^{2}4d^{10}5p^{3}$
Bismuth (Bi)	83	$1s^{2}2s^{2}2p^{6}3s^{2}3p^{6}4s^{2}3d^{10}$ $4p^{6}5s^{2}4d^{10}5p^{6}6s^{2}4f^{14}5d^{10}6p^{3}$

What is the general trend for the elements in group 5A? They all have, as the electron configuration in the outermost energy level, ns²np³, where n is the period number. These elements are three groups away from the noble gases and the outermost energy level electron configuration is three away from having a completed outer energy level. In other words, the p sublevel in the Group 15 elements is half full. Arsenic, for example, has the electron configuration 1s²2s²2p⁶3s²3p⁶4s²3d¹⁰4p³ so it is in period 4, the fifth element in the main group elements. We know it is the fifth element across the period of the main group elements because there are 5 electrons in the outermost energy level.

Lesson Summary

Families in the periodic table are the vertical columns and are also referred to as groups.
Group 1A elements are the alkali metals and all have one electron in the outermost energy level because their electron configuration ends in s¹.
Group 2A elements are the alkaline earth metals and all have two electrons in the outermost energy level because their electron configuration ends in s².
Group 5A elements all have five electrons in the outermost energy level because their electron configuration ends in s²p³.
Group 6A elements all have six electrons in the outermost energy level because their electron configuration ends in s²p⁴.
Group 7A elements are the halogens and all have seven electrons in the outermost energy level because their electron configuration ends in s²p⁵.
Group 8A elements are the noble gases and all have eight electrons in the outermost energy level because their electron configuration ends in s²p⁶.
Elements in group 8A have the most stable electron configuration in the outermost shell because the sublevels are completely filled with electrons.

Review Questions

If an element is said to have an outermost electronic configuration of ns²np³, it is in what group in the periodic table?
(a) Group 3A

(b) Group 4A

(c) Group 5A

(d) Group 7A
What is the general electronic configuration for the Group 8A elements? (Note: when we wish to indicate an electron configuration without specifying the exact energy level, we use the letter "n" to represent any energy level number. That is, ns²np³ represents any of the following; 2s²2p³, 3s²3p³, 4s²4p³, and so on.)
(a) ns²np⁶

(b) ns²np⁵

(c) ns²np¹

(d) ns²
The group 2 elements are given what name?
(a) alkali metals

(b) alkaline earth metals

(c) halogens

(d) noble gases
Using the diagram below, identify:
(a) The alkali metal by giving the letter that indicates where the element would be located and write the outermost electronic configuration.

(b) The alkaline earth metal by giving the letter that indicates where the element would be located and write the outermost electronic configuration.

(c) The noble gas by giving the letter that indicates where the element would be located and write the outermost electronic configuration.

(d) The halogen by giving the letter that indicates where the element would be located and write the outermost electronic configuration.

(e) The element with an outermost electronic configuration of s²p³ by giving the letter that indicates where the element would be located.

(f) The element with an outermost electronic configuration of s²p¹ by giving the letter that indicates where the element would be located.
In the periodic table, name the element whose outermost electronic configuration is found below. Where possible, give the name of the group.
(a) 5s²

(b) 4s²3d¹⁰4p¹

(c) 3s²3p³

(d) 5s²4d¹⁰5p²

(e) 3s¹

(f) 1s²

(g) 6s²5d¹⁰6p⁵

(h) 4s²4p⁴

Vocabulary

alkali metals: Group 1 in the periodic table (Li, Na, K, Rb, Cs, Fr).
alkaline earth metals: Group 2 in the periodic table (Be, Mg, Ca, Sr, Ba, Ra).
group: Columns of the periodic table.
halogens: Group 17 in the periodic table (F, Cl, Br, I, At).
main group elements: Equivalent to the s + p blocks of the periodic table, also known as "representative elements".
noble gases: Group 18 in the periodic table (He, Ne, Ar, Kr, Xe, Rn).
period: Horizontal rows of the periodic table.

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This material was adapted from the original CK-12 book that can be found here. This work is licensed under the Creative Commons Attribution-Share Alike 3.0 United States License