# 8.1: Ions

Difficulty Level: At Grade Created by: CK-12

## Lesson Objectives

• Be able to determine the number of valence electrons for any element and draw an electron dot diagram for any atom.
• Use the octet rule to predict the charges of the most common ions formed by the representative elements.
• Write electron configurations for ions.
• Identify other atoms or ions that are isoelectronic with a particular ion.
• Know that transition metal ions with either half-filled or completely filled d sublevels are particularly stable.

## Lesson Vocabulary

• electron dot diagram
• isoelectronic
• octet rule

### Recalling Prior Knowledge

• What is the Aufbau principle?
• What is the name of the energy required to remove an electron from a neutral atom?
• How does the energy required to remove the first electron from an atom compare to the energy necessary to remove subsequent electrons?

As you learned in previous chapters, ions are formed when atoms lose or gain electrons. Metal atoms have relatively few valence electrons, so when neutral metals undergo chemical reactions, they tend to lose those valence electrons. Nonmetals have more valence electrons than metals, so when nonmetals undergo reactions with metals, they tend to gain electrons. In this lesson, we take a closer look at ions and the physical and chemical properties of ionic compounds.

## Electron Dot Diagrams

Recall that the valence electrons of an atom are the electrons that reside in the highest occupied principal energy level. Valence electrons are primarily responsible for the chemical properties of various elements. The number of valence electrons can be easily determined from the electron configuration. Several examples from the second period are shown in Table below. Note that for each of the second period elements, the valence electrons are in the second principal energy level.

Element Electron Configuration Number of Valence Electrons
lithium \begin{align*}1s^22s^1\end{align*} 1
beryllium \begin{align*}1s^22s^2\end{align*} 2
nitrogen \begin{align*}1s^22s^22p^3\end{align*} 5
neon \begin{align*}1s^22s^22p^6\end{align*} 8

For elements in the s and p blocks, the number of valence electrons can easily be determined from the group number. In the s block, Group 1 elements have one valence electron, while Group 2 elements have two valence electrons. In the p block, the number of valence electrons is equal to the group number minus ten. Group 13 elements have three valence electrons, Group 14 elements have four, and so on. The noble gases in Group 18 have eight valence electrons, and the full outer s and p sublevels are what give these elements their special stability.

When examining chemical bonding, it is necessary to keep track of the valence electrons on each atom. An electron dot diagram shows the valence electrons of an atom as dots distributed around the element’s symbol. For example, a beryllium atom, which has two valence electrons, would have the electron dot diagram below.

Since electrons repel each other, the dots for a given atom are distributed evenly around the symbol before they are paired. Table below shows the electron dot diagrams for the entire second period.

Electron Dot Diagrams for the Second Period Elements
Group Number Electron Dot Diagram
1
2
13
14
15
16
17
18

Electron dot diagrams for all elements in a given group of representative elements are identical (except for the elemental symbol), because each element in one of those groups has the same number of valence electrons. Most transition metals have two valence electrons in their ground state, though some of the elements with unusual electron configurations have only one.

Go to http://hyperphysics.phy-astr.gsu.edu/hbase/pertab/perlewis.html to answer the following question:

From this periodic table, explain how valence electrons are added to the symbol from one column to the next.

## The Octet Rule

The noble gases are unreactive because of their electron configurations. American chemist Gilbert Lewis (1875-1946) used this observation to explain the types of ions and molecules that are formed by other elements. He called his explanation the octet rule. The octet rule states that elements tend to form compounds in ways that give each atom eight valence electrons. An exception to this rule is the elements in the first period, which are particularly stable when they have two valence electrons. A broader statement that encompasses both the octet rule and this exception is that atoms react in order to achieve the same valence electron configuration as that of a noble gas. Most noble gases have eight valence electrons, but because the first principal energy level can hold a maximum of two electrons, the first noble gas (helium) needs only two valence electrons to fill its outermost energy level. As a result, the nearby elements hydrogen, lithium, and beryllium tend to form stable compounds by achieving a total of two valence electrons.

There are two ways in which atoms can satisfy the octet rule. One way is by sharing their valence electrons with other atoms, which will be covered in the next chapter. The second way is by transferring valence electrons from one atom to another. Atoms of metallic elements tend to lose all of their valence electrons, which leaves them with an octet from the next lowest principal energy level. Atoms of nonmetallic elements tend to gain electrons in order to fill their outermost principal energy level with an octet.

### Cations

As you have seen before, cations are the positive ions formed when an atom loses one or more electrons. The cations most commonly formed by the representative elements are those that involve the loss of all valence electrons. Consider the alkali metal sodium (Na). It has one valence electron in the third principal energy level. Upon losing that electron, the sodium ion now has an octet of electrons from the second principal energy level. The equation below illustrates this process.

\begin{align*} &\text{Na} && \rightarrow &&\text{Na}^+ + e^- \\ &1s^22s^22p^63s^1 && && 1s^22s^22p^6 \ \text{(octet)} \end{align*}

The electron configuration of the sodium ion is now the same as that of the noble gas neon. The term isoelectronic is used to describe two atoms or ions that have the same electron configuration. The sodium ion is isoelectronic with the neon atom. Consider the analogous processes for magnesium and aluminum:

\begin{align*} &\text{Mg} && \rightarrow &&\text{Mg}^{2+} + 2e^- \\ & 1s^22s^22p^63s^2 && && 1s^22s^22p^6 \ \text{(octet)} \end{align*}

\begin{align*} &\text{Al} && \rightarrow &&\text{Al}^{3+} + 3e^- \\ & 1s^22s^22p^63s^23p^1 && && 1s^22s^22p^6 \ \text{(octet)} \end{align*}

Both of these atoms form ions by losing all of their valence electrons, two in the case of magnesium, and three in the case of aluminum. The same noble gas configuration is achieved by all of these ions. In other words, the Mg2+ ion, the Al3+ ion, the Na+ ion, and the Ne atom are all isoelectronic. Under typical conditions, the representative elements form cations by losing a maximum of three electrons.

We can also show the loss of valence electron(s) with an electron dot diagram.

\begin{align*}\text{Na}^{\bullet} \rightarrow \text{Na}^+ + e^-\end{align*}

### Anions

Anions are the negative ions formed by gaining one or more electrons. When nonmetal atoms gain electrons, they often do so until they reach an octet of valence electrons in their outermost principal energy level. This process is illustrated below for the elements fluorine, oxygen, and nitrogen.

\begin{align*} &\text{F} + e^- && \rightarrow && \text{F}^- \\ &1s^22s^22p^5 && && 1s^22s^22p^6 \ \text{(octet)} \end{align*}

\begin{align*} &\text{O} + 2e^- && \rightarrow && \text{O}^{2-} \\ &1s^22s^22p^4 && && 1s^22s^22p^6 \ \text{(octet)} \end{align*}

\begin{align*} &\text{N} + 3e^- && \rightarrow && \text{N}^{3-} \\ &1s^22s^22p^3 && && 1s^22s^22p^6 \ \text{(octet)} \end{align*}

All of these anions are isoelectronic with each other and with neon. They are also isoelectronic with the three cations from the previous section. Under typical conditions, a maximum of three electrons will be gained during the formation of anions.

Outer electron configurations are constant within a group, so this pattern of ion formation repeats itself for Periods 3, 4, and so on (Figure below).

It is important not to misinterpret the concept of being isoelectronic. A sodium ion is very different from a neon atom (Figure below), because their nuclei contain different numbers of protons. One is an essential ion that is a part of table salt, while the other is an unreactive gas that makes up a very small part of the atmosphere. Likewise, sodium ions are very different than magnesium ions, fluoride ions, and all the other members of this isoelectronic series (N3-, O2-, F-, Ne, Na+, Mg2+, Al3+).

Neon atoms and sodium ions are isoelectronic. Neon is a colorless and unreactive gas that glows a distinctive red-orange color in a gas discharge tube. Sodium ions are most commonly found in crystals of sodium chloride, which is the chemical name for ordinary table salt.

You can go to http://web.jjay.cuny.edu/~acarpi/NSC/3-atoms.htm to see animations of atoms and ions.

Learning the octet rule can be fun! Watch this music video about the octet rule: http://www.youtube.com/watch?v=WzWk-mx_14E (6:30)

1. How does this song compare an outer energy level with 8 electrons to emotions? 2. What are the two exceptions to the octet rule in this song?

### Transition Metal Ions

Transition metals belong to the d block, meaning that the d sublevel of electrons is in the process of being filled with up to ten electrons. Many transition metals cannot lose enough electrons to attain a noble gas electron configuration. Additionally, you have learned that the majority of transition metals are capable of adopting ions with different charges. Iron, which can form either Fe2+ or Fe3+ ions, loses electrons as shown below.

\begin{align*} &\text{Fe} && \rightarrow && \text{Fe}^{2+} + 2e^- \\ &\text{[Ar]}3d^64s^2 && && \text{[Ar]}3d^6 \end{align*}

\begin{align*} &\text{Fe} && \rightarrow && \text{Fe}^{3+} + 3e^- \\ &\text{[Ar]}3d^64s^2 && && \text{[Ar]}3d^5 \end{align*}

According to the Aufbau process, the electrons fill the 4s sublevel before beginning to fill the 3d sublevel. However, the outermost s electrons are always the first to be removed when forming transition metal cations. Because most transition metals have two valence electrons, a charge of 2+ is very common for transition metal ions, as we have already seen in the case of iron. A half-filled d sublevel (d5) is also particularly stable. This type of configuration is obtained when an iron atom loses a third electron.

(A) Iron(III) oxide, Fe2O3, is the primary component of rust. (B) Iron(II) sulfate, FeSO4, has been known since ancient times as green vitriol and was used for centuries in the manufacture of inks.

Some transition metals that have relatively few d electrons, such as scandium, may be able to attain a noble gas electron configuration.

\begin{align*} &\text{Sc} && \rightarrow && \text{Sc}^{3+} + 3e^- \\ &\text{[Ar]}3d^14s^2 && && \text{[Ar]} \end{align*}

Others may attain configurations that include a full d sublevel, such as zinc and copper.

\begin{align*} &\text{Zn} && \rightarrow && \text{Zn}^{2+} + 2e^- \\ &\text{[Ar]}3d^{10}4s^2 && && \text{[Ar]}3d^{10} \end{align*}

\begin{align*} &\text{Cu} && \rightarrow && \text{Cu}^+ + e^- \\ &\text{[Ar]}3d^{10}4s^1 && && \text{[Ar]}3d^{10} \end{align*}

The resulting configuration above, with 18 electrons in the outermost principal energy level, is referred to as a pseudo noble gas electron configuration. It gives particular stability to the Zn2+ and Cu+ ions.

## Lesson Summary

• An electron dot diagram shows the chemical symbol of an element with dots that represent valence electrons evenly distributed around the symbol.
• The octet rule states that elements form chemical compounds so that each atom will acquire the electron configuration of a noble gas. Most noble gases have eight valence electrons, except for helium, which has only two.
• Representative metals generally lose all of their valence electrons when forming ions, leaving them with a complete octet of electrons from the next-lowest energy level. Most nonmetals gain electrons when forming ions until their outer energy level has acquired an octet.
• Atoms and ions that have the same electron configuration are called isoelectronic. Common ions of representative elements are isoelectronic with a noble gas.
• When forming ions, transition metals lose their valence s-sublevel electrons before they lose their d-sublevel electrons. Half-filled or completely filled d sublevels give transition metal ions greater stability.

## Lesson Review Questions

### Reviewing Concepts

1. What is the maximum number of valence electrons that an atom can have?
2. State the number of protons and electrons in each of the following ions.
1. K+
2. F-
3. P3-
4. Ti4+
5. Cd2+
6. Cr3+
3. What is wrong with this statement? “When a chlorine atom gains an electron, it becomes an argon atom.”
4. Why can the majority of transition metals form 2+ ions?
5. What is a pseudo noble gas electron configuration?

### Problems

1. How many electrons must each of the atoms below lose to achieve a noble gas electron configuration?
1. Li
2. Sr
3. Al
4. Ba
2. Write the symbol of the most common ion formed by each element in problem 6, and name the noble gas with which each ion is isoelectronic.
3. How many electrons must each of the atoms below gain to achieve a noble gas electron configuration?
1. Br
2. S
3. N
4. I
4. Write the symbol of the most common ion formed by each element in problem 8, and name the noble gas with which each ion is isoelectronic.
5. Write electron configurations for each of the following atoms. Then write the symbol for the most common ion each would form and the electron configuration of that ion.
1. Be
2. Cl
3. Se
4. Rb
6. Write electron configurations for the following ions.
1. Cs+
2. Y3+
3. Ni2+
4. As3-
5. Te2-
6. Ag+
7. Pb4+
8. Mn2+
7. For each ion in problem 11, state whether it has (1) a noble gas configuration, (2) a pseudo noble gas configuration, or (3) neither.
8. Split the following ions into isoelectronic groups by noble gas: O2-, Sr2+, Ca2+, H-, V5+, I-, Ba2+, Na+, S2-, Al3+, La3+, Li+, As3-.

## Points to Consider

Ionic compounds adopt the structure of an extended, three-dimensional lattice of alternating positive and negative ions held together by electrostatic attractive forces.

• How strong is an ionic crystal?
• Is an ionic crystal malleable or brittle? Why?
• Will ionic compounds conduct an electric current?

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