Oct 19, 2017

The Mole Concept and Chemical Compounds

Once we know the chemical formula of a compound, we can determine its formula mass. Formula mass is the mass of a formula unit in atomic mass units. It is always appropriate to use the term formula mass, but, for a molecular compound, the formula unit is an actual molecule, so we can speak of molecular mass. Molecular mass is the mass of a molecule in atomic mass units.
Weighted-average formula and molecular masses can be obtained just by adding up weighted-average atomic masses (those on the inside front cover). Thus, for the molecular compound water, H2O,

Mole of a Compound
Recall that in previous blog posts, a mole was defined as an amount of substance having the same number of elementary entities as there are atoms in exactly 12 g of pure carbon-12. This definition carefully avoids saying that the entities to be counted are always atoms. As a result, we can apply the concept of a mole to any quantity that we can represent by a symbol or formula—atoms, ions, formula units, or molecules. Specifically, a mole of compound is an amount of compound containing Avogadro's number (6.02214 x 1023) of formula units or molecules. The molar mass is the mass of one mole of compound—one mole of molecules of a molecular compound and one mole of formula units of an ionic compound.
The weighted-average molecular mass of H2O is 18.0153 u, compared with a mass of exactly 12 u for a carbon-12 atom. If we compare samples of water molecules and carbon atoms by using Avogadro's number of each, we get a mass of 18.0153 g H2O, compared with exactly 12 g for carbon-12. The molar mass of H2O is 18.0153 g H2O/mol H2O. If we know the formula of a compound, we can equate the following terms, as illustrated for H2O, MgCl2, and Mg(NO3)2.

Such expressions as these provide several different types of conversion factors that can be applied in a variety of problem-solving situations. The strategy that works best for a particular problem will depend, in part, on how the necessary conversions are visualized. As we learned in previous posts, the most direct link to an amount in moles is through a mass in grams, so generally the central focus of a problem is the conversion of a mass in grams to an amount in moles, or vice versa. This conversion must often be preceded or followed by other conversions involving volumes, densities, percentages, and so on. As we saw in previous posts, one helpful tool in problem solving is to establish a conversion pathway. In Table 3.1, we summarize the roles that density, molar mass, and the Avogadro constant play in a conversion pathway.

Mole of an Element - A Second Look
In previous blog posts, we took one mole of an element to be 6.02214 X 1023 atoms of the element. This is the only definition possible for such elements as iron, magnesium, sodium, and copper, in which enormous numbers of individual spherical atoms are clustered together, much like marbles in a can. But the atoms of some elements are joined together to form molecules. Bulk samples of these elements are composed of collections of molecules. The molecules of P4 and Ss are represented in Figure 3-5. The molecular formulas of elements that you should become familiar with are

H2 O2 N2 F2 CI2 Br2 I2 P4 Ss

For these elements, we speak of an atomic mass or a molecular mass, and molar mass can be expressed in two ways. Hydrogen, for example, has an atomic mass of 1.00794 u and a molecular mass of 2.01588 u; its molar mass can be expressed as 1.00794 g H/mol H or 2.01588 g H2/mol H2.
Another phenomenon occasionally encountered is the existence of an element in more than one molecular form, a situation referred to as allotropy. Thus, oxygen exists in two allotropic forms, the predominantly abundant diatomic oxygen, O2, and the much less abundant allotrope ozone, O3. The molar mass of ordinary dioxygen is 31.9988 g O2/molO2, and that of ozone is 47.9982 g O3/mol O3.