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CH104: Chapter 3 – Ions and Ionic Compounds

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3.1 Introduction to the Octet Rule

3.2 Ions and the Periodic Table

Common CationsCommon AnionsIons of Transition Metals

3.3 Ionic Bonding

3.4 Practice Writing Correct Ionic Formulas

3.5 Naming Ions and Ionic Compounds

3.6 Polyatomic Ions

3.7 Naming Polyatomic Ions

3.8 Properties and Types of Ionic Compounds

3.9 Arrhenius Acids and Bases

3.10 Focus on the Environment – Acid Rain

3.11 Chapter Summary

3.12 References

3.1 Introduction to the Octet Rule

Up until now we have been discussing only the elemental forms of atoms which are neutrally charged. This is because the number of electrons (negative in charge) is equal to the number of protons (positive in charge). The overall charge on the atom is zero, because the magnitude of the negative charge is the same as the magnitude of the positive charge. This one-to-one ratio of charges is not, however, the most common state for many elements. Deviations from this ratio result in charged particles called ions.

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Throughout nature, things that are high in energy tend to move toward lower energy states. Lower energy configurations are more stable, so things are naturally drawn toward them. For atoms, these lower energy states are represented by the noble gas elements. These elements have electron configurations characterized by full s and p subshells. This makes them stable and unreactive. They are already at a low energy state, so they tend to stay as they are.

The elements in the other groups have subshells that are not full, so they are unstable when compared to the noble gases. This instability drives them toward the lower energy states represented by the noble gases that are nearby in the periodic table. In these lower energy states, the outermost energy level has eight electrons (an “octet”). The tendency of an atom toward a configuration in which it possesses eight valence electrons is referred to as the “Octet Rule.

There are two ways for an atom that does not have an octet of valence electrons to obtain an octet in its outer shell. One way is the transfer of electrons between two atoms until both atoms have octets. Because some atoms will lose electrons and some atoms will gain electrons, there is no overall change in the number of electrons, but with the transfer of electrons the individual atoms acquire a nonzero electric charge. Those that lose electrons become positively charged, and those that gain electrons become negatively charged. Recall that atoms carrying positive or negative charges are called ions. If an atom has gained one or more electrons, it is negatively charged and is called an anion. If an atom has lost one or more electrons, it is positively charged and is called a cation. Because opposite charges attract (while like charges repel), these oppositely charged ions attract each other, forming ionic bonds. The resulting compounds are called ionic compounds.

The second way for an atom to obtain an octet of electrons is by sharing electrons with another atom. These shared electrons simultaneously occupy the outermost shell of both atoms. The bond made by electron sharing is called a covalent bond. Covalent bonding and covalent compounds will be discussed in Chapter 4 “Covalent Bonding and Simple Molecular Compounds”.

At the end of chapter 2, we learned how to draw the electron dot symbols to represent the valence electrons for each of the elemental families. This skill will be instrumental in learning about ions and ionic bonding. Looking at Figure 3.1, observe the Noble Gas family of elements. The electron dot symbol for the Nobel Gas family clearly indicates that the valence electron shell is completely full with an octet of electrons. If you look at the other families, you can see how many electrons they will need to gain or lose to reach the octet state. Above, we noted that elements are the most stable when they can reach the octet state. However, it should also be noted that housing excessively high negative or positive charge is unfavorable. Thus, elements will reach the octet state and also maintain the lowest charge possible. You will note that for the IA, IIA, IIIA and transition metals groups, it is more economical to lose electrons (1-3 electrons) from their valence shells to reach the octet state, rather than to gain 5-7 electrons. Similarly main group columns VA, VIA, and VIIA tend to gain electrons (1-3) to complete their octet, rather than losing 5-7 electrons. Some atoms, like carbon, are directly in the middle. These atoms don’t like to gain or lose electrons, but tend to favor the sharing model of chemical bonding. The remaining sections of this chapter will focus on the formation of ions and the resulting ionic compounds.


Figure 3.1 Periodic Table with Electron Dot Symbols.


Figure 3.2 Ionization Within and Electric Field. (A) Depiction of St. Elmo’s Fire at the tips of a ship’s masts. (B) In many high voltage applications plasma ionization is an unwanted side effect. Shown is a long exposure photograph of corona discharge on an insulator string of a 500 kV overhead power line. This type of plasma discharge represent a significant power loss for electric utilities.

Photograph depicted in a (A) by: Unknown Author

Photograph depicted in a (B) by: Nitromethane

3.2 Ions and the Periodic Table

The elements on the right side of the periodic table, nonmetals, gain the electrons necessary to reach the stable electron configuration of the nearest noble gas. Elements on the left side of the periodic table, metals, lose the electrons necessary to reach the electron configuration of the nearest noble gas. Transition elements can vary in how they move toward lower energy configurations.

Common Cations

Group IA elements form ions with a +1 charge. They lose one electron upon ionization, moving into the electron configuration of the previous noble gas. For example as shown in Figure 3.3, when a sodium (Na) atom is ionized, it loses one of its 11 electrons, becoming a sodium ion (Na+) with the electron configuration that looks like the previous noble gas, neon. The sodium ion has one fewer electron than it has protons, so it has a single positive charge and is called a cation.


Figure 3.3 The Formation of a Sodium Ion.  Sodium tends to lose it’s valence shell electron in the third shell during ionic bond formation. It is left with a full octet in the second shell and now has the electron configuration of neon. Note that it still has the same number of protons (11) as the original sodium atom and retains the identity of sodium. However, there are now only 10 electrons within the electron cloud, resulting in a net positive (+1) charge.

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.

Na → Na+ + e− 1s22s22p63s1 1s22s22p6(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:

  Mg → Mg2+ + 2e−1s22s22p63s2 1s22s22p6(octet) Al → Al3+ + 3e−1s22s22p63s23p1 1s22s22p6(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 elemental Ne atom are all isoelectronic. For most elements under typical conditions, three electrons is the maximum number that will be lost or gained. Only larger atoms, such as lead and uranium, can typically carry larger charge states.

Overall, Group IIA elements lose two valence electrons to reach the electron configuration of the noble gas preceding them in the periodic table and Group IIIA elements lose three electrons to form ions with a +3 charge. This gives them the electron configuration of the noble gas that comes before them in the periodic table.

While hydrogen is in the first column, it is not considered to be an alkali metal, and so it does not fall under the same classification as the elements below it in the periodic table. This is because hydrogen only has an s-subshell and can only house a total of 2 electrons to become filled and obtain the electron configuration of helium. Thus, instead of following the octet rule, it reaches greater stability by gaining a “duet” of electrons through bonding with other atoms. Thus, hydrogen can form both covalent bonds and ionic bonds, depending on the element that it is interacting with. When it participates in ionic bonds, it most often will lose its electron forming a +1 cation. Note, that hydrogen only has one electron to begin with, so when it loses an electron in the ionized state, there is only a single proton left in the nucleus of the atom. Thus, when hydrogen is ionized to H+ it is often referred to as a proton. It can also be ionized, forming a -1 anion. In this case, the H– anion is named using standard convention forming the hydride ion. During the ionization of hydrogen, the H+ state is more common than the H– state. In addition, the H+ ion is very important in the chemistry of acids. Acids are defined as compounds that donate H+ ions in aqueous solutions, and will be discussed in more detail in Chapter 9.


Common Anions

Elements on the other side of the periodic table, the nonmetals, tend to gain electrons in order to reach the stable electron configurations of the noble gases that come after them in the periodic table.

Group VIIA elements gain one electron when ionized, obtaining a -1 charge. For example as shown in Figure 3.4, chlorine (Cl), when ionized, gains an electron to reach the electron configuration of the noble gas that follows it in the periodic table, argon. This gives it a single negative charge, and it is now a chloride ion (Cl–); note the slight change in the suffix (-ide instead of -ine) to create the name of this anion.



Fig 3.4 The Formation of a Chloride Ion. On the left, a chlorine atom has 17 electrons. On the right, the chloride ion has gained an extra electron for a total of 18 electrons and a 1– charge. Note that the chloride ion has now filled its outer shell and contains eight electrons, satisfying the octet rule.

Group VIA elements gain two electrons upon ionization, obtaining -2 charges and reaching the electron configurations of the noble gases that follow them in the periodic table. Whereas, Group VA elements gain three electrons, obtaining -3 charges and also reaching the electron configurations of the noble gases that follow in the periodic table.

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.

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F + e− → F− 1s22s22p5 1s22s22p6(octet) O + 2e− → O2− 1s22s22p4 1s22s22p6(octet) N + 3e− → N3− 1s22s22p3 1s22s22p6(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.

It is important not to misinterpret the concept of being isoelectronic. A sodium ion is very different from a neon atom 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+)