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Chapt03

Accounting

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Back Forward Main Menu TOC Study Guide TOC Textbook Website MHHE Website
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C H A P T E R 3 Mass Relationships in Chemical Reactions I N T R O D U C T I O N I N THIS CHAPTER WE WILL CONSIDER THE MASSES OF ATOMS AND MOL - ECULES AND WHAT HAPPENS TO THEM WHEN CHEMICAL CHANGES OC - CUR . O UR GUIDE FOR THIS DISCUSSION WILL BE THE LAW OF CONSER - VATION OF MASS . 3.1 ATOMIC MASS 3.2 MOLAR MASS OF AN ELEMENT AND AVOGADRO’S NUMBER 3.3 MOLECULAR MASS 3.4 THE MASS SPECTROMETER 3.5 PERCENT COMPOSITION OF COMPOUNDS 3.6 EXPERIMENTAL DETERMINATION OF EMPIRICAL FORMULAS 3.7 CHEMICAL REACTIONS AND CHEMICAL EQUATIONS 3.8 AMOUNTS OF REACTANTS AND PRODUCTS 3.9 LIMITING REAGENTS 3.10 REACTION YIELD 69 Back Forward Main Menu TOC Study Guide TOC Textbook Website MHHE Website
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Section 3.4 describes a method for determining atomic mass. In this chapter we will use what we have learned about chemical structure and for- mulas in studying the mass relationships of atoms and molecules. These relationships in turn will help us to explain the composition of compounds and the ways in which composition changes. The mass of an atom depends on the number of electrons, protons, and neutrons it contains. Knowledge of an atom’s mass is important in laboratory work. But atoms are extremely small particles—even the smallest speck of dust that our unaided eyes can detect contains as many as 1 10 16 atoms! Clearly we cannot weigh a single atom, but it is possible to determine the mass of one atom relative to another experimentally. The first step is to assign a value to the mass of one atom of a given element so that it can be used as a standard. By international agreement, atomic mass (sometimes called atomic weight ) is the mass of the atom in atomic mass units (amu). One atomic mass unit is defined as a mass exactly equal to one-twelfth the mass of one carbon-12 atom. Carbon-12 is the carbon isotope that has six protons and six neutrons. Setting the atomic mass of car- bon-12 at 12 amu provides the standard for measuring the atomic mass of the other el- ements. For example, experiments have shown that, on average, a hydrogen atom is only 8.400 percent as massive as the carbon-12 atom. Thus, if the mass of one carbon- 12 atom is exactly 12 amu, the atomic mass of hydrogen must be 0.0084 12.00 amu or 1.008 amu. Similar calculations show that the atomic mass of oxygen is 16.00 amu and that of iron is 55.85 amu. Thus, although we do not know just how much an av- erage iron atom’s mass is, we know that it is approximately fifty-six times as massive as a hydrogen atom. AVERAGE ATOMIC MASS When you look up the atomic mass of carbon in a table such as the one on the inside front cover of this book, you will find that its value is not 12.00 amu but 12.01 amu. The reason for the difference is that most naturally occurring elements (including car- bon) have more than one isotope. This means that when we measure the atomic mass of an element, we must generally settle for the average mass of the naturally occur- ring mixture of isotopes. For example, the natural abundances of carbon-12 and car- bon-13 are 98.90 percent and 1.10 percent, respectively. The atomic mass of carbon-
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