The Classification of Matter. Chemistry is an experimental physical science, concerned with the properties of and relationships between matter and energy. We may define these two fundamental manifestations of nature in the traditional way. Matter occupies space and has mass. Energy is usually defined as the the ability to do work; however, this definition is not particularly useful in elucidating the nature of energy. A feel for energy is better obtained by trying to think of examples of things that are NOT matter. For instance, light is not matter, nor is heat. Gravity is not matter, nor is it energy. It is instead a force. Forces and energy are intimately connected, so we may view gravity as a disguised form of energy.

Scientists have found it useful to classify matter in several different ways, two of which are discussed here. First, three general categories of matter are recognized: elements, compounds, and mixtures. Elements are the simplest forms of matter; they cannot be broken down (decomposed) into simpler substances by physical or chemical methods. There are 109 known elements, 92 of which occur in nature. Elements are often referred to as the building blocks. Each element is assigned a name and a symbol, which usually consists of the first one or two letters of the element name (symbols that violate this rule are derived from older names for the elements that are seldom used). The first letter of an element symbol is capitalized. Names and symbols for the known elements are presented on the inside front cover of the book. Compounds are combinations of two or more elements in definite proportions by mass. For example, sodium chloride (table salt) is a compound that contains 1.54 g of chlorine per gram of sodium. Any pure sample of sodium chloride, no matter where or how obtained, contains these two elements in this mass ratio. Sodium chloride is only one of more than six million known compounds. Other familiar examples are water, methane (natural gas), ammonia, and sugar. Together, elements and compounds are called pure substances. This implies that they are homogeneous (uniform throughout), have a constant invariable composition, and possess characteristic physical and chemical properties. For example, water contains 11% hydrogen and 89% oxygen by mass; is a clear, colorless liquid with a freezing point of 0^oC and boiling point of 100^oC (all physical properties); and reacts violently when exposed to metallic sodium (a chemical property). A physical property of a pure substance is a characteristic of that substance alone. It can be determined and described without changing the substance into other substances. A chemical property of a substance can be observed and described only in terms of at least one other substance. For example, water and sodium react violently.

Mixtures occur more commonly in nature than do pure substances. A mixture is composed of two or more pure substances, and is characterized by variable composition. For example, a mixture of salt and sand (a compound of the elements, silicon and oxygen) can contain any amount of salt or sand. Mixtures may be homogeneous (having the same composition throughout), such as a solution of sugar in water; or heterogeneous (having nonuniform composition), such as the salt/sand mixture above. A mixture may be separated into components using physical methods, which do not alter the chemical composition of any of the pure substances composing it. Thus we may separate salt from sand by treating the mixture with water to dissolve the salt; filtering the resulting mixture to remove the sand; and finally evaporating the water from the salt solution, leaving the salt behind. A solution , though homogeneous, may nonetheless have variable composition. Any amount of salt, up to a maximum limit, can be dissolved in a given amount of water. Solutions are called homogeneous mixtures. A simple classification scheme for matter is shown in Figure B-1.

Second, matter is classified by physical state; that is, is it solid, liquid, or gas? A solid has a definite volume, a definite shape, and no tendency to flow. A liquid has a definite volume, but flows readily to take the shape of its container. A gas has neither definite volume nor shape; it completely fills any container that it occupies. A gas is easily compressed- forced to occupy a smaller volume-whereas liquids and solids are nearly incompressible. Air is gas; water is liquid; sand is solid.

Chemists are interested in the properties of substances. For example, s/he may be interested in the mass of gas in a bulb; in what volume of water is contained in a flask; or in the temperature of a block of dry ice. Mass, volume, and termperature are examples of properties, to which we can assign numerical values via measurement. Mass and volume are examples of extensive properties, which are defined as properties the values of which depend upon the amount (extent) of substance present. Doubling the amount of gas in a flask doubles the mass of gas. Temperature, in contrast, is an intensive property, which is a property the value of which is independent of the amount of substance present. The temperature of a block of dry ice is the same whether the block is the size of a pea or the size of a house.

Intensive properties are generally considered to be more fundamental than extensive properties, because they are characteristic of substances under specified conditions, independent of the amount of substance available. An example of a very useful intensive property is density, defined as the mass of substance per unit volume of substance. Density is expressend in units of g/mL for solids and liquids, and in g/L for gases. For example, if a 12.0-mL sample of water is found to weight 11.98 g at 21^oC, the density, r, is calculated as 11.98 g/12.0 mL = 0.998 g/mL at 21^oC. Water at 21^oC, no matter how much of it we have, has a density of 0.998 g/mL. This number is characteristic of water at the specified temperature, and serves to identify it. Significantly, density, an intensive property, is expressed as the ratio of two extensive properties, mass and volume. This is always true of intensive properties. Other examples of intensive properties that are ratios of extensive properties are pressure (force/area), heat capacity (heat/unit amount), and cell voltage (energy/charge). All of these quantities are discussed at some point in the text.

In addition to the distinction between extensive and intensive properties, chemists distinguishe physical and chemical properties, as briefly mentioned earlier. A physical property of a substance can be measured without converting the substance into other substances. For example, chlorine is a pale green gas with density 2.94 g/L at 25^oC. It can be liquefied at -34.6^oC and solidified at -101.0^oC, under 1 atm of pressure. Color, density, boiling and freezing point are examples of physical properties because their determination does not involve converting a substance into other substances. In contrast, chemical properties are statements of the behavior of a substance when in the presence of other substances. For example, chlorine reacts vigorously with sodium metal, evolving light energy and heat energy, to produce sodium chloride. In the process, chlorine is destroyed. Baking soda reacts vigorously with vinegar to produce bubbles of carbon dioxide and a solution of sodium acetate. In the process, the baking soda disappears, to be replaced by other substances.

Chemical Nomenclature. There are in excess of 4 million chemical compounds known. A haphazard approach to naming them rapidly becomes unworkable. For this reason, chemists have devised systematic rules for naming compounds. There are two sets of rules -- one for organic compounds (compounds of carbon), the other for inorganic compounds (compounds not containing carbon). The rules for naming organic compounds will not concern us in this book. The system of rules for naming inorganic compounds was developed by the International Union of Pure and Applied Chemistry (IUPAC); it provides meaningful, unambiguous names for most inorganic compounds, and is the required system for authors of modern chemical publications.

For purposes of naming, most compounds can be considered to consist of two parts, one "positive", the other "negative." In ionic (non-molecular) compounds, this is actually true. In molecular compounds, it is possible to associate a positive nature with one portion of the molecule, a negative nature with the other. With very few exceptions, the positive part of the molecule is named first and listed first in the formula. The negative part is named and written last.

a. Ionic compounds. Monatomic cations are commonly formed from metallic elements. They take the name of the element. Thus:

If an element can form more than one positive ion, the positive charge is indicated by a Roman numeral in parentheses following the name of the metal.

An older method still used to distinguish two differently charged ions of a metal is to use the endings -ous and -ic for the lower and higher charged ions, respectively. These suffixes are used with the Latin root for the element name:

The two common polyatomic cations are ammonium ion, NH₄⁺, and mercury(I) ion, Hg₂²⁺.

Monatomic anions are formed from atoms of the nonmetallic elements, and are named by replacing the end of the element name with -ide.

Only a few polyatomic anions have names ending in -ide:

Table B-1 contains names for the most common cations and anions. Note that there are many polyatomic anions containing oxygen. These are called oxoanions. A given element, for example, sulfur, may form more than one oxoanion. When this occurs, there are rules for indicating the relative numbers of oxygen atoms in the anion. When an element forms two oxoanions, they are given the suffixes -ate and -ite, respectively, to indicate more and less oxygen:

For series of 4 members (oxoanions of the halogens, group 17), prefixes are also used. The prefix hypo-means less oxygen; per- means more oxygen:

If you memorize these rules, the name of one oxoanion in a series is sufficient to deduce the rest.

Finally, polyatomic anions with high charges readily add one or more hydrogen ions (H⁺) to give anions of lower charge. These are named by prefixing the word hydrogen or dihydrogen, as appropriate, to the name of the hydrogen-free oxoanion. An older method, unfortunately still used, is to use the prefix bi-:

With these ion-naming rules, we can proceed to name compounds. First, write the name of the cation, with Roman numeral charge designation if necessary; then follow this by the name of the anion:

The formulas are written to give electrically neutral species.

Hydrates. Ionic compounds that crystallize from water (aqueous) solution often contain one or more molecules of water of hydration, and are called hydrates. The number of molecules of water per formula unit of compound is designated using the numerical prefixes below.

The formulas are written as usual, with the water of hydration written following the formula and separated from it by a dot:

Finally, ionic compounds containing a cation other than H⁺ and an anion other than OH^- are referred to as salts.

b. Acids. There is an important class of compounds, the acids, named by special rules. They are formed from hydrogen ions (the positive part) and an anion (the negative part). When the anion is monatomic, the acid name has the prefix hydro-, followed by the anion name with -ide changed to -ic:

Many important acids are derived from oxoanions. When the anion name ends in -ate, the corresponding acid name ends in -ic. Similarly, anion -ite goes with acid -ous. Prefixes in the anion name are retained in the acid name:

Table B-2 has a listing of many common acids.

c. Molecular compounds (binary compounds composed of two nonmetals). When 2 or more nonmetals combine, the result is usually a covalent compound. To name such compounds, state first the name of the "positive" element, followed by the name of the "negative" element with its last syllable changed to -ide. The positive element is always given first in both name and formula. The number of atoms of each element is indicated by the prefixes in Table B-3.

The prefix mono- is often dropped. Examples follow.

Why are different systems used for ionic and covalent compounds? Covalent compounds consist of discrete molecules, and it is important for the name to indicate the actual numbers of atoms of each type of element in a molecule. The name phosphorus(V) oxide for P₄O₁₀ gives no indication of the numbers of P and O atoms, so is not a good name. In contrast to covalent compounds, ionic compounds do not consist of discrete molecules. Therefore the formula for an ionic compound is written as an empirical formula. The roman numeral system suffices to indicate relative numbers of ions of each type.

To avoid uncertainty as to which system to use, we adopt the arbitrary convention that elements to the left of the zigzag line will be considered metals, those to the right nonmetals. The only exception is hydrogen, which is a nonmetal. Thus

d. Common names. A few compounds are so common that we use their common names rather than systematic ones: