Janet L. Kaczmarek
Janet.Niebling@sunysb.edu

All atoms are electrically neutral even though they are comprised of charged, subatomic particles. The terms, oxidation state and oxidation number, have been developed to describe this "electrical state" of the atom. The oxidation state and oxidation number of an atom or ion is simply defined as the sum of the negative and positive charges in an atom or ion. Since every atom contains equal numbers of positive and negative charges, the oxidation state or oxidation number of any atom is always zero. This is illustrated by simply totaling the opposite charges of the atoms as shown by the following examples.

Ionic Compounds
The simplest type of reactions, oxidation-reduction (coupled), are those which occur between metals and nonmetals of the Representative Elements. The transfer of electrons between the atoms of these elements results in drastic changes to the elements involved. This is due to the formation of ionic compounds. The reaction between sodium and chlorine serves as a typical example. The element sodium is a rather "soft" metal solid, with a silver-grey color. Chlorine is greenish colored gas. When a single electron is transferred from the sodium atom to the chlorine atom, their atoms are transformed via a violent reaction into a totally different substance called, sodium chloride, commonly called table salt -- a white, crystalline, and brittle solid.

Sodium chloride exhibits properties quite distinct and different from sodium and chlorine. The changes in physical as well as chemical properties are due to the formation of cations and anions via the oxidation-reduction process, and the resultant, powerful attractive force that develops between these oppositely charge ions. This force of attraction is called the ionic, electrostatic or electrovalent bond, and serves to keep the sodium and chloride ions tightly bound in a highly organized network or lattice of alternating positive and negative charges. This entire complex of ions is called an ionic compound, and is illustrated below in two dimensions. Note how the oppositely charged ions are arranged.

Ionic Formulas
As indicated in the previous section, all ionic compounds are comprised of a definite ratio of cations and anions. This ratio of ions within the ionic compound is determined by the oxidation state of the cation and anion. In every ionic compound, the total positive charge of the cations must always equal the total negative charge of the anions, so that the net charge of the complex is always zero. Every ionic compound can be described by an ionic formula unit (or Empirical Formula) which lists the simplest whole number ratio of the ions in the ionic crystal lattice formed. The simplest whole number ratio of the sodium and chloride ions the network of ions shown above is:

Hence, the chemical formula for the ionic compound sodium chloride, " NaCl ".

Ionic Formula Unit (Empirical Formula)
The nature of any ionic compound can be accurately described by writing its chemical formula, called the ionic formula unit or Empirical formula, which lists the simplest whole number ratio of metal to nonmetal within the lattice structure of the compound. This notation uses a combination of atomic symbols and simple whole numbers, which are written as subscripts, to indicate the simplest whole number ratio of metal ions to nonmetal ions in the ionic compound formed.
Example:
atomic symbols identifying element
                            X Y₂
subscripts, identifying simple whole number ratio of the atoms in the compound. By convention the subscript,"1", is never written.

The following steps provides a simple guideline, based on the concepts presented previously, which allows one to quickly write correct formulas of ionic compounds formed by the oxidation-reduction reaction involving metals and nonmetals. (1) for any metal/nonmetal combination, list their symbols along with the proper oxidation state of the ions formed. Examples:    { note: by convention, the subscript "1" is not written }                             Na+1 O-2            Al+3 F-1           Ca+2 O-2 (2) Test to see if the ionic compound is neutral (the number of positive charges equals the number of negative charges) by adding the products of: the cation oxidation number and subscript, and anion oxidation number and subscript. [(Cation oxid. number x cation subscript) + (Anion oxid. number x Anion subscript)] = 0 If the sum equals zero, the compound is neutral.   If not, increase (or decrease) the subscript of one or both elements until the formula above equals 0 (compound is neutral). For example, the metal/nonmetal combinations used above are listed here as neutral compounds:                               Na2O                 AlF3                  Ca2O2 {note: because the compounds are neutral they do not have a charge associated with them.  It is understood that their overall charge is equal to 0}. For example, take this compound   Na2O2-2   it has an overall charge associated with it, therefore its overall charge is obviously not neutral or 0.  This is due to incorrect subscripts! (3) and if needed, reduce to the simplest whole number ratio, Empirical Formula.                                Example:

Transition Metals:
The behavior of the Transition metals is similar to that of the Representative metals. They are also oxidized by nonmetals, losing their electrons to the nonmetal and forming ionic compounds. However, many Transition metals exhibit multiple oxidation states, forming cations with different positive charges. This is due to the fact that many Transition Metals are characterized by a partially filled inner electron level, inside the valence shell. Electrons within this inner shell may sometimes behave as valence electrons and are lost along with the outermost electrons during oxidation. The number of electrons lost depends on the conditions under which the chemical reactions occur. Hence, many of these metals can exhibit "multiple oxidation" states, forming cations of different charges.

A typical example is iron. Depending on the conditions of the reaction, iron may form a cation with a "+2" or "+3" charge, by losing two or three electrons, respectively.  Roman numerals are commonly used to indicate the valence state of the transition element.  Ions of iron (Fe⁺² and Fe⁺³) may be indicated by writing Iron (II) and Iron (III) respectively.  Manganese, another Transition metal and an extreme example, may exist in the following oxidation states: "+2, +3, +4, +6, and +7, by losing 2, 3, 4, 6, or 7 electrons, respectively. Because the number of electrons lost by the metal depends on so many variables (temperature, amount and nature of nonmetal, etc.) the exact chemical formula of ionic compounds formed by the Transition Metals must be determined experimentally. The simple whole number ratio of the atoms in the derived formula can then be used to determine the oxidation state of the Transition Metal.