Chemical Nomenclature Kevin Pyatt, Ph.D. Donald Calbreath, Ph.D. Say Thanks to the Authors Click http://www.ck12.org/saythanks (No sign in required) AUTHORS Kevin Pyatt, Ph.D. To access a customizable version of this book, as well as other Donald Calbreath, Ph.D. interactive content, visit www.ck12.org EDITORS Donald Calbreath, Ph.D. Max Helix CK-12 Foundation is a non-profit organization with a mission to reduce the cost of textbook materials for the K-12 market both in the U.S. and worldwide. Using an open-content, web-based collaborative model termed the FlexBook®, CK-12 intends to pioneer the generation and distribution of high-quality educational content that will serve both as core text as well as provide an adaptive environment for learning, powered through the FlexBook Platform®. 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Printed: August 20, 2014 www.ck12.org Chapter 1. Chemical Nomenclature CHAPTER 1 Chemical Nomenclature CHAPTER OUTLINE 1.1 Ionic Compounds 1.2 Molecular Compounds 1.3 Acids and Bases 1.4 References The opening image shows crystals of the mineral cinnabar, which is the most common mercury-containing ore. It is primarily composed of the compound mercuric sulfide, HgS. Cinnabar has been mined since the Stone Age for its uses as a pigment and as a source of pure mercury. Mercury can be easily harvested from cinnabar by roasting cinnabar powder (vermillion), which actually vaporizes the mercury metal. The hot mercury vapor can then be recondensed to yield pure liquid mercury. Cinnabar is one of thousands of minerals that occur naturally on earth. Some minerals are common (like cinnabar) while others are rarer (like ores of gold and silver). Another less common mercury-containing mineral is mon- troydite, which is composed primarily of mercuric oxide (HgO). The composition of mercuric oxide was first determined in 1774 by the English chemist Joseph Priestley. Priestley showed that heating montroydite powder produced mercury metal and a gas, which he called phlogiston-free air. The gas was later determined to be oxygen. The revised name, based on information gained from experiments, allowed chemists to have a better understanding of chemical composition and chemical reactions. Clear unambiguous naming makes communication much easier and more reliable. Parent Gry: commons:wikimedia:org=wiki=File:Cinabre_macl%C3%A9_%28Chine%29_: jpg: Public Domain: 1 1.1. Ionic Compounds www.ck12.org 1.1 Ionic Compounds Lesson Objectives • Define and give examples of chemical formulas for ionic compounds. • Be able to name and write the formulas for both monatomic and polyatomic ions. • Explain and use the Stock system for naming ionic compounds, when necessary. • Memorize the list of polyatomic ions, including both the formulas and the charges. • When given the chemical formula for an ionic compound, be able to write its name. • When given the name of an ionic compound, be able to write its chemical formula. Lesson Vocabulary • empirical formula: The lowest whole-number ratio between two ions. • binary ionic compound: A compound made up of a cation and an anion. • ternary ionic compound: An ionic compound that is composed of more than two elements. • monatomic ion: Form when a single atom gains or loses electrons. • polyatomic ion: An ion composed of more than one atom. • oxoanion: Anions in which one or more oxygen atoms are all bonded to a central atom of some other element. Introduction As we saw in the previous chapter, ions are formed when atoms gain or lose electrons. If an atom loses one or more electrons, the resulting ion has a positive charge (more protons are present than electrons). If the atom gains one or more electrons, the resulting ion has a negative charge (more electrons are present than protons). Positive ions are called cations, and negative ions are called anions. Because opposite charges attract one another, cations and anions are held together by strong electromagnetic forces. An ionic compound consists of a large three-dimensional array of alternating cations and anions. For example, sodium chloride (NaCl) is composed of Na+ and Cl− ions arranged into a structure like the one shown in Figure 1.1. The most straightforward way to describe this structure with a chemical formula is to give the lowest whole-number ratio between the two ions, which is known as an empirical formula. In the case of NaCl, there are equal numbers of sodium ions and chloride ions in the salt crystal. In contrast, a crystal of magnesium chloride has twice as many chloride ions as magnesium ions, so it has a formula of MgCl2. Naming Ionic Compounds Ionic compounds are composed of one type of cation and one type of anion. The name of an ionic compound can be formed by writing the name of the cation followed by the name of the anion. For example, NaCl is composed 2 www.ck12.org Chapter 1. Chemical Nomenclature FIGURE 1.1 A crystal of table salt, sodium chloride, is a large array of alternating positive and negative ions. The purple spheres represent the Na+ ions, while the green spheres represent the Cl− ions. of sodium ions (Na+) and chloride ions (Cl−), so its name is sodium chloride. Ionic compounds like NaCl that are composed of only two elements are referred to as binary ionic compounds. Similarly, KOH is composed of potassium ions (K+) and hydroxide ions (OH−), so its name is potassium hydroxide. Ionic compounds like KOH that are composed of more than two elements are referred to as ternary ionic compounds. To learn how to name various ionic compounds, we simply need to learn the names of individual ions. Monatomic Ions Monatomic ions form when a single atom gains or loses electrons. For the main group elements, cations are generally formed by removing all of the valence electrons from the atom. Since the numbers of valence electrons for the representative elements are constant within a particular group, all we need is the group number of a given element to know its charge when it becomes a cation. Group 1 elements form ions with a 1+ charge, Group 2 metal ions have a 2+ charge, and the ions of Group 13 elements tend to have a 3+ charge. Heavier p-block metals such as tin and lead are special cases and will be discussed with the transition metal ions. The name of a monatomic cation is the same as the name of the neutral element. For example, the sodium atom (Na) loses a single electron to form the sodium ion (Na+), while Al3+ is an aluminum ion. Anions form when an atom gains electrons. Nonmetallic atoms typically gain enough electrons to obtain the same electron configuration as the nearest noble gas. All the elements in Group 17 have seven valence electrons, which are arranged into a outer configuration of ns2np5. To achieve a noble gas configuration (ns2np6), each of these elements needs to gain just one electron, resulting in an anion with a 1− charge. Similarly, Group 16 elements can obtain an ns2np6 valence configuration by forming ions with a 2− charge, and the Group 15 nonmetals will form ions with a 3− charge. Naming anions is slightly different than naming cations. The end of the element’s name is dropped and replaced with the –ide suffix. For example, when the chlorine atom (Cl) gains one electron, it becomes the chloride ion (Cl−). This structure has the same electron configuration as the noble gas argon. Similarly, sulfur can gain two electrons to become the sulfide ion (S2−), which also has a noble gas configuration. Most main group elements, particularly those in groups 1, 2, 16, and 17, gain or lose enough electrons to form ions that have the same electron configuration as that of the nearest noble gas. Table 1.1 shows the names and charges for common monatomic ions of the representative elements: TABLE 1.1: Common Monatomic Ions 1+ 2+ 3+ 3- 2- 1- lithium, Li+ beryllium, Be2+ aluminum, Al3+ nitride, N3− oxide, O2− fluoride, F− sodium, Na+ magnesium, gallium, Ga3+ phosphide, P3− sulfide, S2− chloride, Cl− Mg2+ potassium, K+ calcium, Ca2+ arsenide, As3− selenide, Se2− bromide, Br− 3 1.1. Ionic Compounds www.ck12.org TABLE 1.1: (continued) 1+ 2+ 3+ 3- 2- 1- rubidium, Rb+ strontium, Sr2+ telluride, Te2− iodide, I− cesium, Cs+ barium, Ba2+ Transition Metal Ions Most transition metals differ from the metals of Groups 1, 2, and 13 in that they are capable of forming more than one type of stable cation. For example, iron sometimes loses two electrons to form the Fe2+ ion, but it is also common for iron to lose three electrons to form the Fe3+ ion. Although they are members of the p block and not the d block, tin and lead also form more than one type of ion.