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Lesson Objectives

  • Be able to determine the number of valence electrons for any element and show with an electron dot diagram.
  • Use the octet rule to predict the charges of the most common ion of representative elements
  • Write electron configurations for ions.
  • Identify what atoms or other ions are isoelectronic with a particular ion.
  • Understand the particular stability that comes from transition metal ions that have either half-filled or completely filled d sublevels.

Lesson Vocabulary

  • electron dot diagrams
  • isoelectronic
  • octet rule

Check Your Understanding

Recalling Prior Knowledge

  • What is the Aufbau principle?
  • What energy is associated with the removal of an electron from a neutral atom?
  • What happens to this energy when a second or third electron is removed from a given atom?

As learned in previous chapters, ions are formed when atoms lose or gain electrons. Metal atoms have relatively few valence electrons, so when metals undergo chemical reactions, they tend to lose those valence electrons. Nonmetal atoms have more valence electrons than metals, so when nonmetals undergo reactions with metals, they tend to gain electrons. The chapter Chemical Nomenclature focused on the systems of naming and writing chemical formulas. In this lesson, we take a closer look at ions so that we can understand the physical and chemical properties of ionic compounds.

Electron Dot Diagrams

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

Element Electron Configuration Number of Valence Electrons
lithium 1s^22s^1 1
beryllium 1s^22s^2 2
nitrogen 1s^22s^22p^3 5
neon 1s^22s^22p^6 8

In each case, valence electrons are those in the second principal energy level. As one proceeds left to right across a period, the number of valence electrons increases by one. 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 has three valence electrons, Group 14 has four, up through Group 18 with eight. The eight valence electrons, a full outer s and p sublevel, give the noble gases their special stability.

When examining chemical bonding, it is necessary to keep track of the valence electrons of each atom. Electron dot diagrams are diagrams in which the valence electrons of an atom are shown as dots distributed around the element’s symbol. A beryllium atom, with 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 would be the same for each element in the representative element groups. Most transition elements have two valence electrons, though some that have 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 atoms tend to form compounds in ways that give them eight valence electrons and thus the electron configuration of a noble gas. An exception to an octet of electrons is in the case of the first noble gas, helium, which only has two valence electrons. This primarily affects the element hydrogen, which forms stable compounds by achieving 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 metals tend to lose all of their valence electrons, which leaves them with an octet from the next lowest principal energy level. Atoms of nonmetals 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 by the loss of one or more electrons. The most commonly formed cations of the representative elements are those that involve the loss of all of the 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.

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

The electron configuration of the sodium ion is now the same as that of the noble gas neon. The term isoelectronic refers to an atom and an ion of a different atom, or two different ions, that have the same electron configuration. The sodium ion is isoelectronic with the neon atom. Consider a similar process with magnesium and with aluminum:

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

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

In this case, the magnesium atom loses its two valence electrons in order to achieve the same noble-gas configuration. The aluminum atom loses its three valence electrons. The Mg2+ ion, the Al3+ ion, the Na+ ion, and the Ne atom are all isoelectronic. For representative elements under typical conditions, three electrons is that maximum number that will be lost.

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

\text{Na}\cdot \rightarrow \text{Na}^+ + e^-

Anions

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

&\text{F} + e^- && \rightarrow && \text{F}^- \\&1s^22s^22^p5 && && 1s^22s^22p^6 \ \text{(octet)}

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

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

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, three electrons is the maximum that will be gained in the formation of anions.

Outer electron configurations are constant within a group, so this pattern of ion formation repeats itself for Periods 3, 4, etc. (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 the nuclei of the two 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 is 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, 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. In addition, you have learned that the majority of transition metals are capable of adopting ions with different charges. Iron, which forms either the Fe2+ or Fe3+ ions, loses electrons as shown below.

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

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

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 in the process of forming transition metal cations. Because most transition metals have two valence electrons, the charge of 2+ is a very common one for their ions. This is the case for iron above. A half-filled d sublevel (d5) is particularly stable, which is the result of an iron atom losing a third electron.

(A) Rust is a complex combination of oxides of iron(III), among them iron(III) oxide, Fe2O3. (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 may attain a noble-gas electron configuration. Scandium is an example.

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

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

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

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

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. In most instances, that is eight electrons except for helium which has only two electrons.
  • Representative metals lose their valence electrons when forming ions, leaving them with a complete octet of the next-lowest energy level. Most nonmetals gain electrons when forming ions until their outer energy level as 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.
  • In 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. For the four elements in number 6, write the symbol of the most common ion and state which noble gas the ion is isoelectronic with.
  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. For the four elements in number 8, write the symbol of the most common ion and state which noble gas the ion is isoelectronic with.
  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 number 11, state whether each 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-.

Further Reading / Supplemental Links

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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Date Created:

Aug 02, 2012

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Aug 01, 2014
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