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. 

Formula and molecular masses can be obtained just by adding up atomic masses (those on the inside front cover). Thus, for the molecular compound water, H2O, 

A The terms formula weight and molecular weight are often used in place of formula mass and molecular mass. This is similar to the situation described for atomic mass and atomic weight in the footnote on page 48. 

A The terms formula mass and molecular mass have essentially the same meaning, although when referring to ionic compounds, such as NaCl and MgCl2, formula mass is the proper term. 

Recall that in Chapter 2 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 10 ) 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. 

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The Mole Concept and Chemical Compounds—In this section, the concept of atomic mass is extended to molecular mass, the mass in atomic mass units of a molecule of a molecular compound, and formula mass, the mass in atomic mass units of a formula unit of an ionic compound. Likewise, the concept of the Avogadro constant and the mole is now applied to compounds, with emphasis on quantitative applications involving the mass of a mole of compound— the molar mass M. For several elements, we can distinguish between a mole of molecules (for example, P4) and a mole of atoms (that is, P). 

The empirical formula for a substance tells us the relative number of atoms of each element it contains. Thus, the formula H2O indicates that water contains two H atoms for each O atom. This ratio also applies on the molar level thus, 1 mol of H2O contains 2 mol of H atoms and 1 mol of O atoms. Conversely, the ratio of the number of moles of each element in a compoimd gives the subscripts in a compound s empirical formula. Thus, the mole concept provides a way of calculating the empirical formulas of chemical substances, as shown in the following examples. 

According to Section 2.1 and as demonstrated in Example 2.8, the formula for a compound is made up of the symbols for each element present. Subscripts following the elemental symbols indicate the number of each type of atom in the molecule represented. Thus, chemical formulas represent the numerical relationships that exist among the atoms in a compound. AppUcation of the mole concept to the atoms making up the formulas provides additional useful information. 

The mole is often referred to as a chemist s unit of quantity. Counting atoms is a difficult process and beyond the scope of most calculators, but measuring the mass of a sample is easy when we can relate the number of atoms in a sample to its mass. This is the unique purpose of the mole. A mole of any substance is its molecular formula weight expressed in grams. Avogadro s number s a universal constant that states the number of molecules in a mole Nq = 6.023 x 10 molecules/mole. One mole (abbreviated mol) of any element (chemical compound) has the same number of chemical particles as one mole of another element (chemical compound). In other words, 1 mole of any compound contains 6.02 x 10 molecules. Review the following problem using the mole concept. 

Methanol synthesis will be used many times as an example to explain some concepts, largely because the stoichiometry of methanol synthesis is simple. The physical properties of all compounds are well known, details of many competing technologies have been published and methanol is an important industrial chemical. In addition to its relative simplicity, methanol synthesis offers an opportunity to show how to handle reversible reactions, the change in mole numbers, removal of reaction heat, and other engineering problems. 

The stoichiometric calculations of Chapters 12 and 13 are based on the mole as the fundamental chemical unit in reactions. An alternative method of calculation utilizes the equivalent as a fundamental chemical unit. There are two kinds of equivalents, the type depending on the reaction in question we shall refer to them as acid-base equivalents (or simply as equivalents) and electron-transfer equivalents (or E-T equivalents). The concept of an equivalent is particularly useful when dealing with complex or unknown mixtures, or when working out the structure and properties of unknown compounds. In addition, it emphasizes a basic characteristic of all chemical reactions that is directly applicable to all types of titration analyses. 

I he previous chapters showed how the laws of conservation of mass and con--1- servation of atomic identity, together with the concept of the mole, determine quantitative mass relationships in chemical reactions. That discussion assumed prior knowledge of the chemical formulas of the reactants and products in each equation. The far more open-ended questions of which compounds are found in nature (or which can be made in the laboratory) and what types of reactions they undergo now arise. Why are some elements and compounds violently reactive and others inert Why are there compounds with chemical formulas H2O and NaCl, but never H3O or NaCli Why are helium and the other noble gases monatomic, but molecules of hydrogen and chlorine diatomic All of these questions can be answered by examining the formation of chemical bonds between atoms. 

Today, Avogadro s name is most often associated with his number. The mole and the modem definition of the number of particles in it, however, were not directly his invention. His fame hes in a single simple statement, made in 1811 .. . the number of integral molecules in gases is always the same for equal volumes." This concept opened the door to understanding atomic weight and the formulas for chemical compounds. (The modem version of this concept is one mole of any gas always occupies the same volume [22.4 L] at 0°C and I atm of pressure.)... 

Before the electron was discovered in 1897, oxidation and reduction reactions were already well-known concepts and so were chemical bonding and conductivity. It was suspected that a particle, the electron, was behind these phenomena. We now know that oxidation of a component of a system is the same as removal of electrons and that reduction is the same as addition of electrons. A chemical compound can only be oxidized if something else is reduced at the same time. In an electrochemical cell, ions from the electrolyte are reduced and deposited on the cathode. Corresponding oxidation reactions take place at the anode. The time-integrated current was originally measured in equivalents of charge. One equivalent corresponds to one mole of deposited metal, thus it is equal to Avogadro s number of electrons. 

The molar mass of a Key Concepts compound can be calculated from its chemical formula and can be used to convert from mass to moles of that compound. 

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