Atomic mass units defined

How is the atomic mass unit defined How has this definition evolved ... 

The SI unit for the atomic mass m is given in kg. The additional unit used in chemistry is the atomic mass unit defined as... 

Because the masses of nuclides are so small, they are normally reported as a multiple of the atomic mass constant, ma (formerly atomic mass unit, amu). The atomic mass constant is defined as exactly V12 the mass of one atom of carbon-12 ... 

The weight of atoms and their constituents can be given in kilograms. A proton, for example, weighs 1.67 x 10 kilograms, but its weight or mass can be expressed more conveniently in a measure called the atomic mass unit (amu). One amu is defined as 1/12 the mass of a carbon atom that consists of six protons. 

The formula weight of a substance is equal to its number of grams per mole. Avogadro s number is the number of atomic mass units in 1 g. It is defined in that manner so that the atomic weight of an element (in amu) is numerically equal to the number of grams of the element per mole. Consider helium, with atomic weight 4.0 ... 

MS equipment is evaluated on several performance metrics. Mass accuracy, mass resolution, and mass range are standard parameters frequently assessed to determine the suitability of an instrument. Mass accuracy is defined as the extent to which a mass analyzer reflects true m/z values and is measured in atomic mass units (amu), parts per million (ppm), or percent accuracy. 

Unified Atomic Mass Unit (u) A non-SI unit of mass defined as one twelfth of the mass of one atom of 12C in its ground state and 1.66 x 10-27 kg. The term atomic mass unit (amu) is not recommended to use since it is ambiguous. It has been used to denote atomic masses measured relative to a single atom of 160, or to the isotope-averaged mass of an oxygen atom, or to a single atom of 12C. 

The unified atomic mass unit (u), previously symbolized as AMU or amu, is defined to be 1/12 of the mass of one atom of isotope 12 of carbon. Therefore,... 

Note Care has to be taken when mass values from dated literature are cited. Prior to 1961 physicists defined the atomic mass unit [amu] based on Vie of the mass of one atom of nuclide 0. The definition of chemists was based on the relative atomic mass of oxygen which is somewhat higher resulting from the nuclides and contained in natural oxygen. 

At one time, the hydrogen atom with one proton and no neutron was used as the standard to define 1 atomic mass unit (1 amu). Today, chemists use carbon-12, the most abundant isotope of carbon for the standard amu, which is defined as 1/12 of the C-12 atom. Therefore, the actual atomic weight for an element is in average mass units (numbers), taking into account all the isotopes (atoms) of that element. 

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. 

By means of mass spectrometry, the mass of atoms and molecules, via mass-separated charged atomic or polyatomic ions, can be determined by measuring the mass-to-charge ratio m/z), whereby the mass of an atom or a molecule is not measured in g or kg, but in a multiple of the atomic mass constant mu (atomic mass unit). The atomic mass unit mu is defined as one-twelfth the mass of a neutral 12C atom, ma (12C), in its ground state ... 

Another convenient unit for describing the mass of a single atom or molecule is the atomic mass unit (formerly amu, now commonly denoted u). One atomic mass unit (1 u) is defined as one-twelfth the mass of an atom of carbon-12. Since the experimentally measured mass of an atom of carbon-12 is 1.9926 X l(T23g, 1 u = 1.6606 X 10 24g. The atomic mass unit is convenient for describing the mass of a peak observed by mass spectrometry (see Box 3-2). 

ATOMIC MASS UNIT, UNIFIED (u). The atomic mass unit (unified) is l/12th of the mass of anatom of the 12Cnuclide. (Use of the prior atomic mass unit (amu), defined by reference to oxygen, is no longer preferred.)... 

One of the most important nuclear properties that can be measured is the mass. Nuclear or atomic masses are usually given in atomic mass units (amu or u) or their energy equivalent. The mass unit u is defined so that the mass of one atom of 12C is equal to 12.0000. .. u. Note we said atom. For convenience, the masses of atoms rather than nuclei are used in all calculations. When needed, the nuclear mass mllucl can be calculated from the relationship... 

The atomic mass unit (amu) is defined as 1/12 the mass of a carbon-12 isotope. The relative atomic mass of an element is the weighted average of the isotopes relative to I/12(of jhe carbon-12 isotope. For example, the atomic mass of neon is 20.17 amu and is calculated from the following data neon-19 (amu of 19.99245, natural abundance of 90.92%), neon-20 (amu of 20.99396, natural abu dan c of 0,260%) and ncon-21 (amu of 21.99139, natural abundance ofc 82%) ... 

Atoms are so tiny that even the smallest speck of dust visible to the naked eye contains about 1016 atoms. Thus, the mass of a single atom in grams is much too small a number for convenience, and chemists therefore use a unit called an atomic mass unit (amu), also known as a dalton (Da). One amu is defined as exactly one-twelfth the mass of an atom of and is equal to 1.660 539 X 10 24 g ... 

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. 

Elements differ from one another according to how many protons their atoms contain, a value called the atomic number (Z) of the element. The sum of an atom s protons and neutrons is its mass number (A). Although all atoms of a specific element have the same atomic number, different atoms of an element can have different mass numbers, depending on how many neutrons they have. Atoms with identical atomic numbers but different mass numbers are called isotopes. Atomic masses are measured using the atomic mass unit (amu), defined as 1/12 the mass of a 12C atom. Because both protons and neutrons have a mass of approximately 1 amu, the mass of an atom in atomic mass units (the isotopic mass) is numerically close to the atom s mass number. The element s atomic mass is a weighted mass average for naturally occurring isotope mixtures. 

Prior to 1961, the atomic mass unit was defined as 1/16 the mass of the atomic mass of oxygen. That is, the atomic mass of oxygen was defined as exactly 16 amu. What was the mass of a 12C atom prior to 1961 if the atomic mass of oxygen on today s scale is 15.9994 amu ... 

Mass Unit The unified atomic mass unit, or u, is the fundamental unit of mass for most mass spectrometrists. The Dalton, or Da, is also generally accepted and is commonly used in descriptions of large, biological molecules. The mass unit is defined as one-twelfth of the mass of carbon-12. Atomic mass unit, or amu, is technically incorrect but still commonly used. The unit Thomson (Th) has been used as a unit of m jz. However, Th is not accepted by most mass spectrometry journals and the International Union of Pure and Applied Chemistry (IUPAC). Therefore, m/z used for labeling the x-axis of mass spectra is unit less. 

Because the mass of an atom is very small, it is convenient to define a special unit that avoids large negative exponents. This unit, called the atomic mass unit and designated by the symbol u (some authors use the abbreviation amu), is defined as exactly 1/12 the mass of a 12C atom. 

In the future, it will end in the value of the Avogadro constant and the value 1 of the atomic mass unit u defined as the mass of 1/12th of the mass of the 12C atom. This will happen when the definition of the kg (now the mass of the prototype of the kilogram ) will have changed into the mass of a number of12C atoms, i.e. of the mass of NAj-m(12C). 1000/12) . 

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. 

The subatomic particles differ in mass and charge. Their masses are expressed by the atomic mass unit, u (also called the dalton), which is also used to express the masses of individual atoms, and molecules (aggregates of atoms). The atomic mass unit is defined as a mass equal to exactly 1/12 that of an atom of carbon-12, the isotope of carbon that contains six protons and six neutrons in its nucleus. 

Carbon has the distinction of being chosen as the basis for the unit of atomic mass. The atomic mass unit (amu) is defined as l/nth of the mass of the 12C atom. 

Nuclear binding energy This is the energy needed to disassemble a nucleus into its component nucleons, and is an important demonstration of Einstein s formula. It can best be understood by considering the nucleus 12C. The atomic mass unit, by being defined as i/l2th of the 12C atomic mass, causes A for 12C to be identically equal to zero thatis, A(12C) = o. But the 12C atom is assembled from six H atoms (sixprotons with their six electrons) and six neutrons, each of which is more massive than mu. Their atomic mass excesses are... 

Proton—a positively charged particle located in the atom s nucleus. The electrical charge has a magnitude of+ 1.6 X 10 19 coulombs (C) however, for simplicity, it is often referred to by its relative charge of +1.0 (charge relative to an electron). The mass of a proton is about 1.67 X 10 24 g. The gram is not a practical unit to describe the mass of subatomic particles, so instead we use the atomic mass unit, or amu. An amu is defined as the mass of a carbon atom containing 6 protons and 6 neutrons. The mass of a proton is 1.0073 amu. 

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