Chemical Naming and Structure: Writing Names and Formulas

Oxidation Numbers

There are situations in which an atom does not have a clear and unmistakable charge, yet it would be helpful to determine whether or not it has lost or gained electrons. This is particularly true when dealing with oxidation-reduction reactions. In these situations, we need oxidation numbers. By convention in this lesson, oxidation numbers are written with the sign given first, followed by the number. Charges are given in the reverse order. Thus, “+2” is an oxidation number, while “2+” is a charge.

Assigning Oxidation Numbers

Oxidation numbers may be assigned to individual atoms by following these simple rules:

• 1. Elements have oxidation numbers of zero.
• 2. Hydrogen in a compound is usually +1. The only notable exception would be a metal hydride, such as NaH, in which hydrogen would be -1.
• 3. Oxygen in a compound is usually −2. Again, the most important exceptions would be peroxides, such as H₂O₂ or BaO₂, in which oxygen is −1.
• 4. Monatomic ions have an oxidation number equal to their charge.
• 5. All other oxidation numbers are assigned so that the sum of the oxidation numbers equals the charge on the species.

Examples

Rule 1 is straightforward. Fe, N₂, and As are all examples of species whose oxidation numbers equal zero because they are elemental. Please remember that species such as Fe²⁺ or N³⁻ are not elements but ions; thus, they do not fall under this rule.

Rules 2 and 3 are also rather simple. In the water molecule (H₂O), each hydrogen atom has an oxidation number of +1 and the oxygen atom is −2. In sulfurous acid (H₂SO₃), each hydrogen atom is still +1 and each oxygen atom is still −2.

Rule 4 is easy when ions are written separately. The Cr³⁺ ion has an oxidation number of +3 while the Cl¹⁻ ion is −1. The use of this rule is less obvious when the ions are in a compound. For example, potassium sulfide (K₂S) consists of a potassium ion, K+, and a sulfide ion, S²⁻. We do not see these charges in the compound’s formula, so we must infer them. Once we do, we can see that the oxidation numbers are +1 and −2, respectively.

Rule 5 requires the most thought. We invoke this rule when none of the others apply. Let’s reconsider H₂SO₃. The sulfur atom is not a simple sulfide ion here; rather, it is part of the sulfite ion, SO₃²⁻. As such, we cannot assign it a −2 oxidation number. Instead, we remember that each hydrogen is +1 and each oxygen is −2. In addition, bear in mind that the compound as a whole has no charge. Thus, all oxidation numbers must add up to zero. Representing sulfur’s unknown oxidation number with the variable “x”, we may find it with the following expression:

Molecular Names and Formulas

We recall that molecular compounds are formed when nonmetal atoms share valence electrons, frequently to achieve eight valence electrons for each atom. Often, the same two nonmetal atoms may be able to combine in more than one ratio, thus forming more than one potential compound. For example, sulfur is known to combine with oxygen in a 1:2 ratio (SO₂) and in a 1:3 ratio (SO₃). Since multiple combinations are possible, it is important that the name of a molecular compound indicates how many of each atom is present in that compound. This is done with prefixes. Below are listed the prefixes used in naming molecular compounds, along with the numerical value that each prefix represents.

The name of a binary molecular compound follows a straightforward pattern. The first element’s name is listed first, preceded by whichever prefix indicates how many of that element’s atoms are in the compound. The only exception to this is if there is only one of that atom. In that case, the prefix mono is omitted. The name of the second element follows, modified so that it ends in –ide. It is always preceded by the appropriate prefix. A few examples are listed below:

• CS₂ carbon disulfide
• PCl₅ phosphorus pentachloride
• N₂O₄ dinitrogen tetroxide
• O₂F₂ dioxygen difluoride

Overview of Ionic Compounds

In ionic compounds, positive ions combine with negative ions in a ratio such that their charges add up to zero. This means that a particular pair of ions will always combine in the same ratio, which makes the use of prefixes unnecessary. For this reason, when we name any ionic compound, we simply name the positive ion first and the negative ion second. When we are given that name, we must know the appropriate charge on each ion in order to put together a correct formula.

Binary Ionic Compounds

A binary ionic compound contains only monatomic ions. Positive monatomic ions have the same name as the element from which they come. Negative monatomic ions are named by modifying the name of the element so that it ends with –ide. Many of the elements in the s block and the p block of the Periodic Table follow the Octet Rule when bonding ionically, so their charges are fairly easy to predict. The table below summarizes these charges:

Armed with this information, we should be able to write names and formulas for many simple binary ionic compounds. For example, a compound made from calcium and bromine would contain Ca²⁺ ions and Br¹⁻ ions. We would name this calcium bromide. To make the net charge of the compound zero, we would need two bromide ions for every calcium ion, so the formula would be CaBr₂

Ions With More Than One Possible Charge

Many transition metals, as well as several metals near the bottom of the p block (most commonly Sn and Pb), are able to form at least two different ions. Iron, for example, can form Fe²⁺ and Fe³⁺. If we were to name a compound iron sulfide, we would have no idea which ion was in the compound in question. This would then prevent us from being able to write the formula.

We solve this dilemma by using Roman numerals in parentheses as part of the ion’s name. Thus, the name of Fe²⁺ is not just iron, but iron(II). Fe³⁺ is named iron(III). Iron(II) sulfide is made from Fe²⁺ and S²⁻ ions, so its formula is FeS. Iron(III) sulfide contains Fe³⁺ and S²⁻ ions, so its formula is Fe₂S₃.

Ionic Compounds Containing Polyatomic Ions

Many ionic compounds contain polyatomic ions. These compounds follow the same rules that we have seen before, with only a few minor differences. Most polyatomic ions are negative (ammonium, NH₄⁺, is the most common exception), but their names do not usually end with –ide. Most polyatomic ions have names that end with –ite or –ate. Sulfide is simply a sulfur atom with a charge—S²⁻. Sulfate, on the other hand, is a sulfur atom and several other atoms together that have a charge—SO₄²⁻.

The most common polyatomic ions contain an atom along with one or more oxygen atoms. The ion whose name is the element name modified so that it ends in –ate will be considered our “base” ion. For example, if we modify “nitrogen” so that it ends with –ate, we get “nitrate.” So the nitrate ion, NO₃¹⁻, will be considered our base ion. Unfortunately, there is no hard and fast rule regarding how many oxygen atoms are in the base ion, although it is usually three or four. Once the base ion is known, however, there is a pattern to the names of the other ions made with that element. The following table shows this pattern and uses chlorine (base ion of chlorate, ClO₃¹⁻) as an example.

By following this pattern, we should be able to see that since NO₃¹⁻ is the nitrate ion, then NO₂¹⁻ is the nitrite ion.

The only other difference when using polyatomic ions concerns the use of subscripts in the compound’s formula. When a formula requires more than one of a polyatomic ion, we must put the ion formula in parentheses and place the subscript outside of the closing parenthesis. For example, calcium nitrate requires one Ca²⁺ ion and two NO₃¹⁻ ions; therefore, its formula is Ca(NO₃)₂.