Information about Atomic Mass Unit
The unified atomic mass unit (u), or dalton (Da), is a small unit of mass used to express atomic and molecular masses. It is defined to be one twelfth of the mass of an unbound atom of the carbon-12 nuclide, at rest and in its ground state.
See 1 E-27 kg for a list of objects which have a mass of about 1 u.
The symbol amu for atomic mass unit is not a symbol for the unified atomic mass unit. Its use is an historical artifact (written during the time when the amu scales were used), an error (possibly deriving from confusion about historical usage), or correctly referring to the historical scales that used it. (see #History). Atomic masses are often written without any unit and then the unified atomic mass unit is implied. In biochemistry and molecular biology literature (particularly in reference to proteins), the term "dalton" is used, with the symbol Da. Because proteins are large molecules, they are typically referred to in kilodaltons, or "kDa", with one kilodalton being equal to 1000 daltons. The unified atomic mass unit, or dalton, is not an SI unit of mass, although it is accepted for use with SI under either name.
The unit is convenient because one hydrogen atom has a mass of approximately 1 u, and more generally an atom or molecule that contains n protons and neutrons will have a mass approximately equal to n u. (The reason is that a carbon-12 atom contains 6 protons, 6 neutrons and 6 electrons, with the protons and neutrons having about the same mass and the electron mass being negligible in comparison.The mass of the electron is approximately 1/1836 of the mass of the proton) This is an approximation, since it does not account for the mass contained in the binding energy of an atom's nucleus; this binding energy mass is not a fixed fraction of an atom's total mass. The differences which result from nuclear binding are generally less than 0.01 u, however. Chemical element masses, as expressed in u, would therefore all be close to whole number values (within 2% and usually within 1%) were it not for the fact that atomic weights of chemical elements are averaged values of the various stable isotope masses in the abundances which they naturally occur. [1] For example, chlorine has an atomic weight of 35.45 u because it is composed of 76% 35Cl (34.96 u) and 24% 37Cl (36.97 u).
Another reason the unit is used is that it is experimentally much easier and more precise to compare masses of atoms and molecules (determine relative masses) than to measure their absolute masses. Masses are compared with a mass spectrometer (see below).
Avogadro's number (NA) and the mole are defined so that one mole of a substance with atomic or molecular mass 1 u will have a mass of precisely 1 gram. For example, the molecular mass of a water molecule containing one 16O isotope and two 1H isotopes is 18.0106 u, and this means that one mole of this monoisotopic water has a mass of 18.0106 grams. Water and most molecules consist of a mixture of molecular masses due to naturally occurring isotopes. For this reason these sort of comparisons are more meaningful and practical using molar masses which are generally expressed in g/mol, not u. In other words the one-to-one relationship between daltons and g/mol is true but in order to be used accurately for any practical purpose any calculations must be with isotopically pure substances or involve much more complicated statistical averaging of multiple isotopic compositions.
Before 1961, the physical atomic mass unit (amu) was defined as 1⁄16 of the mass of one atom of oxygen-16, while the chemical atomic mass unit (amu) was defined as 1⁄16 of the average mass of an oxygen atom (taking the natural abundance of the different oxygen isotopes into account). Both units are slightly smaller than the unified atomic mass unit, which was adopted by the International Union of Pure and Applied Physics in 1960 and by the International Union of Pure and Applied Chemistry in 1961. Hence, before 1961 physicists as well as chemists used the symbol amu for their respective (and slightly different) atomic mass units. One still sometimes finds this usage in the scientific literature today. However, the accepted standard is now the unified atomic mass unit (symbol u), with: 1 u = 1.000 317 9 amu (physical scale) = 1.000 043 amu (chemical scale).
- 1 u = 1/NA gram = 1/ (1000 NA) kg (where NA is Avogadro's number)
- 1 u ≈ 1.660538782(83) × 10−27 kg ≈ 931.494028(23) MeV/c2
See 1 E-27 kg for a list of objects which have a mass of about 1 u.
The symbol amu for atomic mass unit is not a symbol for the unified atomic mass unit. Its use is an historical artifact (written during the time when the amu scales were used), an error (possibly deriving from confusion about historical usage), or correctly referring to the historical scales that used it. (see #History). Atomic masses are often written without any unit and then the unified atomic mass unit is implied. In biochemistry and molecular biology literature (particularly in reference to proteins), the term "dalton" is used, with the symbol Da. Because proteins are large molecules, they are typically referred to in kilodaltons, or "kDa", with one kilodalton being equal to 1000 daltons. The unified atomic mass unit, or dalton, is not an SI unit of mass, although it is accepted for use with SI under either name.
The unit is convenient because one hydrogen atom has a mass of approximately 1 u, and more generally an atom or molecule that contains n protons and neutrons will have a mass approximately equal to n u. (The reason is that a carbon-12 atom contains 6 protons, 6 neutrons and 6 electrons, with the protons and neutrons having about the same mass and the electron mass being negligible in comparison.The mass of the electron is approximately 1/1836 of the mass of the proton) This is an approximation, since it does not account for the mass contained in the binding energy of an atom's nucleus; this binding energy mass is not a fixed fraction of an atom's total mass. The differences which result from nuclear binding are generally less than 0.01 u, however. Chemical element masses, as expressed in u, would therefore all be close to whole number values (within 2% and usually within 1%) were it not for the fact that atomic weights of chemical elements are averaged values of the various stable isotope masses in the abundances which they naturally occur. [1] For example, chlorine has an atomic weight of 35.45 u because it is composed of 76% 35Cl (34.96 u) and 24% 37Cl (36.97 u).
Another reason the unit is used is that it is experimentally much easier and more precise to compare masses of atoms and molecules (determine relative masses) than to measure their absolute masses. Masses are compared with a mass spectrometer (see below).
Avogadro's number (NA) and the mole are defined so that one mole of a substance with atomic or molecular mass 1 u will have a mass of precisely 1 gram. For example, the molecular mass of a water molecule containing one 16O isotope and two 1H isotopes is 18.0106 u, and this means that one mole of this monoisotopic water has a mass of 18.0106 grams. Water and most molecules consist of a mixture of molecular masses due to naturally occurring isotopes. For this reason these sort of comparisons are more meaningful and practical using molar masses which are generally expressed in g/mol, not u. In other words the one-to-one relationship between daltons and g/mol is true but in order to be used accurately for any practical purpose any calculations must be with isotopically pure substances or involve much more complicated statistical averaging of multiple isotopic compositions.
History
The chemist John Dalton was the first to suggest the mass of one atom of hydrogen as the atomic mass unit. Francis Aston, inventor of the mass spectrometer, later used 1⁄16 of the mass of one atom of oxygen-16 as his unit.Before 1961, the physical atomic mass unit (amu) was defined as 1⁄16 of the mass of one atom of oxygen-16, while the chemical atomic mass unit (amu) was defined as 1⁄16 of the average mass of an oxygen atom (taking the natural abundance of the different oxygen isotopes into account). Both units are slightly smaller than the unified atomic mass unit, which was adopted by the International Union of Pure and Applied Physics in 1960 and by the International Union of Pure and Applied Chemistry in 1961. Hence, before 1961 physicists as well as chemists used the symbol amu for their respective (and slightly different) atomic mass units. One still sometimes finds this usage in the scientific literature today. However, the accepted standard is now the unified atomic mass unit (symbol u), with: 1 u = 1.000 317 9 amu (physical scale) = 1.000 043 amu (chemical scale).
See also
External links
units of measurement have played a crucial role in human endeavour from early ages up to this day. Disparate systems of measurement used to be very common. Now there is a global standard, the International System (SI) of units, the modern form of the metric system.
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Mass is a fundamental concept in physics, roughly corresponding to the intuitive idea of "how much matter there is in an object". Mass is a central concept of classical mechanics and related subjects, and there are several definitions of mass within the framework of relativistic
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atomic mass (ma) is the mass of an atom at rest, most often expressed in unified atomic mass units.[1] The atomic mass may be considered to be the total mass of protons, neutrons and electrons in a single atom (when the atom is motionless).
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molecular mass (abbreviated Mr) of a substance, formerly also called molecular weight and abbreviated as MW, is the mass of one molecule of that substance, relative to the unified atomic mass unit u (equal to 1/12 the mass of one atom of carbon-12).
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Carbon-12 is the more abundant of the two stable isotopes of the element carbon, accounting for 98.89% of carbon; it contains 6 protons, 6 neutrons and 6 electrons.
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A nuclide (from lat.: nucleus) is a nuclear species which is characterized by the number of protons and neutrons that every atomic nucleus of this species contains.
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Gram
Unit sign g
Measure Mass
Base Unit Kilogram
Multiple of Base 10−3
System SI, CGS, other
Common usage Commonly used in cooking and food labeling
Examples
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Unit sign g
Measure Mass
Base Unit Kilogram
Multiple of Base 10−3
System SI, CGS, other
Common usage Commonly used in cooking and food labeling
Examples
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kilogram or kilogramme (symbol: kg) is the SI base unit of mass. The kilogram is defined as being equal to the mass of the International Prototype Kilogram (IPK), which is almost exactly equal to the mass of one liter of water.
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The Avogadro constant (symbols: L, NA), also called the Avogadro number is the number of "entities" (usually, atoms or molecules) in one mole,[1][2] that is the number of carbon-12 atoms in 12 grams (0.
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To help compare different orders of magnitude, the following list describes various mass levels between 10−36 kg and 1053 kg.
Factor (kg) Value Item
10−36 1.
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Factor (kg) Value Item
10−36 1.
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Biochemistry is the study of the chemical processes in living organisms.[1] The word "biochemistry" comes from the Greek word βιοχημεία biochēmeia, which means "the chemistry of life.
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Molecular biology is the study of biology at a molecular level. The field overlaps with other areas of biology and chemistry, particularly genetics and biochemistry. Molecular biology chiefly concerns itself with understanding the interactions between the various systems of a cell,
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Proteins are large organic compounds made of amino acids arranged in a linear chain and joined together by peptide bonds between the carboxyl and amino groups of adjacent amino acid residues.
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molecule is defined as a sufficiently stable electrically neutral group of at least two atoms in a definite arrangement held together by strong chemical bonds.[1][2] In organic chemistry and biochemistry, the term molecule
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Si, si, or SI may refer to (all SI unless otherwise stated):
In language:
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In language:
- One of two Italian words:
- sì (accented) for "yes"
- si
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hydrogen atom is an atom of the chemical element hydrogen. It is composed of a single negatively-charged electron circling a single positively-charged nucleus of the hydrogen atom.
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atom (Greek ἄτομος or átomos meaning "indivisible") is the smallest particle still characterizing a chemical element.
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molecule is defined as a sufficiently stable electrically neutral group of at least two atoms in a definite arrangement held together by strong chemical bonds.[1][2] In organic chemistry and biochemistry, the term molecule
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Proton
The quark structure of the proton.
Composition: 2 up, 1 down
Family: Fermion
Group: Quark
Interaction: Gravity, Electromagnetic, Weak, Strong
Antiparticle: Antiproton
Discovered: Ernest Rutherford (1919)
Symbol: p+
Mass: 1.
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The quark structure of the proton.
Composition: 2 up, 1 down
Family: Fermion
Group: Quark
Interaction: Gravity, Electromagnetic, Weak, Strong
Antiparticle: Antiproton
Discovered: Ernest Rutherford (1919)
Symbol: p+
Mass: 1.
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Neutron
The quark structure of the neutron.
Composition: one up, two down
Family: Fermion
Group: Quark
Interaction: Gravity, Electromagnetic, Weak, Strong
Antiparticle: Antineutron
Discovered: James Chadwick[1]
Symbol: n
Mass: 1.
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The quark structure of the neutron.
Composition: one up, two down
Family: Fermion
Group: Quark
Interaction: Gravity, Electromagnetic, Weak, Strong
Antiparticle: Antineutron
Discovered: James Chadwick[1]
Symbol: n
Mass: 1.
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Electron
Theoretical estimates of the electron density for the first few hydrogen atom electron orbitals shown as cross-sections with color-coded probability density
Composition: Elementary particle
Family: Fermion
Group: Lepton
Generation: First
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Theoretical estimates of the electron density for the first few hydrogen atom electron orbitals shown as cross-sections with color-coded probability density
Composition: Elementary particle
Family: Fermion
Group: Lepton
Generation: First
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Binding energy is the mechanical energy required to disassemble a whole into separate parts. A bound system has a lower potential energy than its constituent parts; this is what keeps the system together.
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The nucleus of an atom is the very small dense region of an atom, in its center consisting of nucleons (protons and neutrons). The size (diameter) of the nucleus is in the range of 1.
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1, 3, 5, 7
(strongly acidic oxide)
Electronegativity 3.16 (Pauling scale)
Ionization energies
(more) 1st: 1251.2 kJmol−1
2nd: 2298 kJmol−1
3rd: 3822 kJmol−1
Atomic radius 100 pm
Atomic radius (calc.
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(strongly acidic oxide)
Electronegativity 3.16 (Pauling scale)
Ionization energies
(more) 1st: 1251.2 kJmol−1
2nd: 2298 kJmol−1
3rd: 3822 kJmol−1
Atomic radius 100 pm
Atomic radius (calc.
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Mass spectrometry (previously called mass spectroscopy ()[1] or informally, "mass-spec" and MS) is an analytical technique used to measure the mass-to-charge ratio of ions.
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The Avogadro constant (symbols: L, NA), also called the Avogadro number is the number of "entities" (usually, atoms or molecules) in one mole,[1][2] that is the number of carbon-12 atoms in 12 grams (0.
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The mole (symbol: mol) is the SI base unit that measures an amount of substance. One mole contains Avogadro's number (approximately 6.0221023) entities.
A mole is much like "a dozen" in that both are absolute numbers (having no units) and can describe any type of
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A mole is much like "a dozen" in that both are absolute numbers (having no units) and can describe any type of
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Gram
Unit sign g
Measure Mass
Base Unit Kilogram
Multiple of Base 10−3
System SI, CGS, other
Common usage Commonly used in cooking and food labeling
Examples
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Unit sign g
Measure Mass
Base Unit Kilogram
Multiple of Base 10−3
System SI, CGS, other
Common usage Commonly used in cooking and food labeling
Examples
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Water (H2O, HOH) is the most abundant molecule on Earth's surface, composing of about 70% of the Earth's surface as liquid and solid state in addition to being found in the atmosphere as a vapor.
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Molar mass, symbol M,[1] is the mass of one mole of a substance (chemical element or chemical compound).[2] It is a physical property which is characteristic of each pure substance.
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