Atomic mass unit table

Our modern model describes the atom as an electrically neutral sphere with a tiny nucleus in the center containing positively charged protons and neutral neutrons. The negatively charged electrons are moving in complex paths outside the nucleus in energy levels at different distances from the nucleus. These subatomic particles have very little mass expressed in grams so we often use the unit of an atomic mass unit (amu or simply u). An amu is 1/12 the mass of a carbon atom that contains six protons and six neutrons. Table 2.1 summarizes the properties of the three subatomic particles. 

A factor Avogadro s constant multiplied by 10 enters these expressions on condition that atomic and electronic masses be expressed, as is customary in spectral analyses, in unified atomic mass unit both Uifi and f/o,i contain mass in their units, despite their values being formally independent of atomic mass. The standard errors associated with values of k and in Table 1 include contributions from errors of pertinent fundamental physical constants [94]. 

The mass number gives the total number of protons and neutrons in an atom of an element, but it does not convey the absolute mass of the atom. To work with the masses of elements, we use comparative masses. Initially, Dalton and the other pioneers of the atomic theory used the lightest element hydrogen and compared masses of other elements to hydrogen. The modern system uses C-12 as the standard and defines one atomic mass unit (amu) as 1/12 the mass of one C-12 atom. One amu is approximately 1.66 X 10 g. This standard means the masses of individual protons and neutrons are slightly more than 1 amu as shown in Table 4.6. 

The total mass of an atom is called its atomic mass. This is the sum of the masses of all the atom s components (electrons, protons, and neutrons). Because electrons are so much less massive than protons and neutrons, their contribution to atomic mass is negligible. As we explore further in Section 9.2, a special unit has been developed for atomic masses. This is the atomic mass unit, amu, where 1 atomic mass unit is equal to 1.661 X 10-24 gram, which is slightly less than the mass of a single proton. As shown in Figure 3.21, the atomic masses listed in the periodic table are in atomic mass units. As is explored in the Calculation Corner on page 95, the atomic mass of an element as presented in the periodic table is actually the average atomic mass of its various isotopes. 

Carbon-13 has a mass of 13.0034 atomic mass units and makes up 1.11 percent of naturally occurring carbon. Use this information to show that the atomic mass of carbor shown in the periodic table, 12.011 atomic mass units, is correct. 

The element bromine, Br (atomic number 35), has two major isotopes of similar abundance, both around 50 percent. The atomic mass of bromine is reported in the periodic table as 79.904 atomic mass units. Choose the most likely set of mass numbers for these two bromine isotopes (a) 80Br, 81 Br (b) 79Br, 80Br (c) 79Br, 81 Br. 

Because the mass of an atom s electrons is negligible compared with the mass of its protons and neutrons, defining 1 amu as 1/12 the mass of a atom means that protons and neutrons each have a mass of almost exactly 1 amu (Table 2.1). Thus, the mass of an atom in atomic mass units—called the atom s isotopic mass—is numerically close to the atom s mass number. A jH atom, for instance, has a mass of 1.007 825 amu a 292U atom has a mass of 235.043 924 amu and so forth. 

About 1837 electrons are equal in mass to the mass of one proton or one neutron. A summary of each type of particle, its mass and relative charge is shown in Table 3.1. You will notice that the masses of all these particles are measured in atomic mass units (amu). This is because they are so light that their masses cannot be measured usefully in grams. 

Arsenic is a group 15 element on the periodic table along with nitrogen, phosphorus, antimony, and bismuth. The atomic mass of arsenic is 74.921 60 atomic mass units (amu) and its atomic number (Z)... 

Table 2-1. Some Nuclidic Masses in Atomic Mass Units...

The masses of individual atoms are very small. Even the heaviest atom discovered has a mass less than 5 x 10-25 kg. Since 1 kg is 2.21b, the mass referred to is less than 1.10 x 10-24 lb. It is convenient to define a special unit in which the masses of the atoms are expressed without having to use exponents. This unit is called the atomic mass unit, referred to by the symbol u in the literature. It is defined as exactly the mass of a 12C atom. The mass of the 12C atom is taken to be exactly 12u the mass of the 23Na atom is 22.9898 u. Table 2-1 lists the masses of some nuclides to which reference will be made in this chapter, as well as others. 

Although there are particles not used here, the basic particles listed in Table 21-1 can be used to define and illustrate the concepts presented. Note that the proton and neutron are referred to as nucleons. The masses in Table 21-1 are presented in atomic mass units (u, Chapter 2) and their charges are expressed in multiples of the elementary charge (1.6022 x 10 19 C, Chapter 19). Note that the neutron has slightly more mass than the proton. Also, the mass of the electron is considered to be 1/1836 that of a proton or, if you prefer, the mass of 1 proton is 1836 times that of an electron. 

Although mass spectra usually show the particle masses rounded to the nearest whole number, the masses are not really integral. The 12C nucleus is defined to have a mass of exactly 12 atomic mass units (amu), and all other nuclei have masses based on this standard. For example, a proton has a mass of about 1, but not exactly Its mass is 1.007825 amu. Table 12-3 shows the atomic masses for the most common isotopes found in organic compounds. 

Examine Figure 5.3. Since the atomic mass unit is based on carbon-12, why does the periodic table show a value of 12.01 u, instead of exactly 12 u Carbon is made up of several isotopes, not just carbon-12. Naturally occurring carbon contains carbon-12, carbon-13, and carbon-14. If all these isotopes were present in equal amounts, you could simply find the average of the masses of the isotopes. This average mass would be about 13 u, since the masses of carbon-13 and carbon-14 are about 13 u and 14 u respectively. 

How can you use this relationship to relate mass and moles The periodic table tells us the average mass of a single atom in atomic mass units (u). For example, zinc has an average atomic mass of 65.39 u. One mole of an element has a mass expressed in grams numerically equivalent to the element s average atomic mass expressed in atomic mass units. One mole of zinc atoms has a mass of 65.39 g. This relationship allows chemists to use a balance to count atoms. You can use the periodic table to determine the mass of one mole of an element. 

The second number, at the bottom of each listed element in the Periodic Table, is called the atomic mass. Atomic masses are not nice, round numbers but rather numbers such as 28.086 and 65.37. If you round off the atomic mass, the number you end up with tells you the total number of protons and neutrons in the most commonly occurring form of the atom. The atomic mass is measured in what are called atomic mass units. In this system, the mass of a proton is equal to exactly 1 and the mass of a neutron is equal to exacdy l. The mass of an electron is really small compared to the masses of protons and neutrons, so we can ignore the electron s mass for most purposes in chemistry. So the electrons are still there in our atoms—we just don t count their mass. 

Table 1), possibly also containing molecules such as CHjOH. In addition, very massive positive ions (upt to about 320 atomic mass units) were detected but could as yet not be identified. 

To convert between moles and grams, chemists use the molar mass of a substance. The molar mass of an element is the mass in grams of one mole of the element. Molar mass has the unit grams per mol (g/mol). The mass in grams of 1 mol of an element is numerically equal to the element s atomic mass from the periodic table in atomic mass units. For example, the atomic mass of copper to two decimal places is 63.55 amu. Therefore, the molar mass of copper is 63.55 g/mol. Skills Toolkit i shows how to convert between moles and mass in grams using molar mass. 

Recall from Table 4-1 that the masses of both protons and neutrons are approximately 1.67 x 10 g. While this is a very small mass, the mass of an electron is even smaller—only about that of a proton or neutron. Because these extremely small masses expressed in scientific notation are difficult to work with, chemists have developed a method of measuring the mass of an atom relative to the mass of a specifically chosen atomic standard. That standard is the carbon-12 atom. Scientists assigned the carbon-12 atom a mass of exactly 12 atomic mass units. Thus, one atomic mass unit (amu) is defined as the mass of a carbon-12 atom. Although a mass of 1 amu is very nearly equal to the mass of a single proton or a single neutron, it is important to realize that the values are slightly different. As a result, the mass of silicon-30, for example, is 29.974 amu, and not 30 amu. Table 4-2 gives the masses of the subatomic particles in terms of amu. 

The atomic mass unit (amu), which is represented with the symbol u, is based on a particular isotope of carbon, called carbon-12. Carbon-12 is considered to have a mass of exactly 12 u, and all of the other elemental isotopes are measured relative to that isotope. The atomic masses, shown on the periodic table, represent a weighted average of the masses of the naturally occurring isotopes of each element. For example, some periodic tables show an atomic mass of 1.00794 u for hydrogen, despite the fact that no particular isotope of hydrogen has a mass number equal to that value. 

The periodic table can also be used to calculate the molar mass of molecules and formula units as well. If you can add up the mass of all of the atoms in a molecule to find the molecular mass in atomic mass units, the molar mass of the same molecular compound will have the same value with the unit, grams (g). Following are some examples ... 

Find the molar mass of the solute, CaCl2. Remember The process of finding the molar mass of a substance is the same as finding the molecular or formula mass of the substance. We look up the masses listed on the Periodic Table of Elements for each of the elements involved and multiply by the appropriate subscripts. The only difference is that you use the unit symbol g for grams, instead of u for atomic mass units. 

The masses for the elements listed in the table inside the back cover of this text are relative masses in terms of atomic mass units (amu) or daltons. The atomic mass unit is based on a relative scale in which the reference is the C carbon isotope, which is assigned a mass of exactly 12 amu. Thus, the amu is by definition 1/12 of the mass of one neutral c atom. The molar mass of is then... 

Atoms vary in the number of protons and neutrons tliey contain, resulting in different atomic weights, these different weights make up the 92 chemical elements in nature and in the periodic table (Fig. 1.3). The atomic number of an element is the number of protons in the nucleus. The atomic mass is the total mass of protons, neutrons and electrons in a single atom, often expressed in unified atomic mass units (AMU). 

You know from Chapter 2 that average atomic masses of the elements are given on the periodic table. For example, the average mass of one iron atom is 55.8 u, where u means atomic mass units. The atomic mass unit is defined so that the atomic mass of an atom of the most common carbon isotope is exactly 12 u, and the mass of 1 mol of the most common isotope of carbon atoms is exactly 12 g. The mass of 1 mol of a pure substance is called its molar mass. For example, the molar mass of iron is 55.847 g, and the molar mass of platinum is 195.08 g. Relative masses of elements are demonstrated in Figure 12.4. The molar mass is the mass in grams of the average atomic mass. 

The reciprocal Angstrom unit is also conventionally used for momentum (A ) throughout the neutron scattering literature. There was neither a common optical spectroscopic alternative (as there was to justify the adoption of the wavenmnber, cm, for the unit of energy) nor chemical reason (as there was to justify the adoption of the atomic mass unit, u or amu, for the unit of mass). For quantities not defined by lUPAC we have mostly used symbols consistent with the neutron scattering literature [14,15] but have, on rare occasions, been forced to invent our own symbol for the sake of clarity. We provide a table of symbols and units (p. xix). 

---