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EspañolMOLE (UNIT)
(Redirected from Millimole)
The 'mole' (symbol: mol) is the SI base unit that measures an amount of substance. One mole contains Avogadro's number (approximately 6.022) entities. A mole is much like "a dozen" in that both are absolute numbers (having no units) and can describe any type of elementary object, although the mole's use is usually limited to measurement of subatomic, atomic, and molecular structures.
A mole is the amount of substance of a system which contains as many elementary entities as there are atoms in 0.012 kilogram (or 12 grams) of carbon-12, where the carbon-12 atoms are unbound, at rest and in their ground state.[1] The number of atoms in 0.012 kilogram of carbon-12 is known as the Avogadro constant, and is determined empirically. The currently accepted value is 6.02214179(30) mol-1 (2007 CODATA).
According to the SI, the mole is not dimensionless, but has its very own dimension, namely "amount of substance", comparable to other dimensions such as mass and luminous intensity. The International System of Units (SI), , , , International Bureau of Weights and Measures, 2006, (By contrast, the SI specifically defines the radian and the steradian as special names for the dimensionless unit one.) The International System of Units (SI), , , , International Bureau of Weights and Measures, 2006, The SI additionally defines the Avogadro constant as having the unit reciprocal mole, as it is the ratio of a dimensionless quantity and a quantity with the unit mole. However, if in the future the kilogram is redefined in terms of a specific number of carbon-12 atoms (see below), then the value of Avogadro's number will be defined rather than measured, and the mole will cease to be a unit of physical significance.[2]
The relationship of the atomic mass unit (u[3])
to Avogadro's number means that a mole can also be defined as: ''That quantity of a substance whose mass in grams is the same as its formula weight.'' For example, iron has an relative atomic mass of 55.845 u, so a mole of iron has a mass of 55.845 grams. This notation is very commonly used by chemists and physicists.
Scientists and engineers (chemical engineers in particular) sometimes measure amount of substance in units of ''gram-moles'', ''kilogram-moles'', ''pound-moles'', or ''ounce-moles''; these measure the quantity of a substance whose mass in grams, kilograms, pounds, or ounces (respectively) is equal to its formula weight. The SI mole is identical to the gram-mole.
When the mole is used to specify the amount of a substance, the kind of elementary entities (particles) in the substance must be identified. The particles can be atoms, molecules, ions, formula units, electrons, photons or other particles. For example, one mole of water is equivalent to 18.016 grams of water and contains one mole of H2O molecules, but three moles of atoms (two moles H and one mole O).
When the substance of interest is a gas, the particles are usually molecules. However, the noble gases (He, Ar, Ne, Kr, Xe, Rn) are all monoatomic, that is each particle of gas is a single atom. All gases have the same molar volume of 22.4 litres per mole at STP (see Avogadro's Law).
A mole of atoms or molecules is also called a "gram atom" or "gram molecule", respectively.
The name ''mole'' (German ''Mol'') is attributed to Wilhelm Ostwald who introduced the concept in the year 1902. It is an abbreviation for ''molecule'' (German ''Molekül''), which is in turn derived from Latin ''moles'' "mass, massive structure".
He used it to express the gram molecular weight of a substance. So, for example, 1 mole of hydrochloric acid (HCl) has a mass of 36.5 grams (atomic masses Cl: 35.5 u, H: 1.0 u).
Prior to 1959 both the IUPAP and IUPAC used oxygen to define the mole, the chemists defining the mole as the number of atoms of oxygen which had mass 16 g, the physicists using a similar definition but with the oxygen-16 isotope only. The two organizations agreed in 1959/1960 to define the mole as such:
:''The mole is the amount of substance of a system which contains as many elementary entities as there are atoms in 0.012 kilogram of carbon-12; its symbol is "mol."''
This was adopted by the CIPM (International Committee for Weights and Measures) in 1967, and in 1971 it was adopted by the 14th CGPM (General Conference on Weights and Measures).
In 1980 the CIPM clarified the above definition, defining that the carbon-12 atoms are unbound and in their ground state.
As with other SI base units, there have been proposals to redefine the kilogram in such a way as to define some presently measured physical constants to fixed values. One proposed definition of the kilogram is:
:''The kilogram is the mass of exactly'' (6.0221415/0.012) ''unbound carbon-12 atoms at rest and in their ground state.'' [4]
This would have the effect of defining Avogadro's number to be precisely ''N''A = 6.0221415 elementary entities per mole, and, consequently, the mole would become merely a unit of counting, like the dozen.
Another proposed definition of ''NA'' is:
: ''NA'' = 602214141070409084099072 = 844468883
This has the nice properties of being a perfect cube, and being near the current experimental bounds of measurement.[5]
The mole is useful in chemistry because it allows different substances to be measured in a comparable way. Using the same number of moles of two substances, both amounts have the same number of molecules or atoms. The mole makes it easier to interpret chemical equations in practical terms. Thus the equation:
:2H2 + O2 → 2H2O
can be understood as "two moles of hydrogen plus one mole of oxygen yields two moles of water."
Moles are useful in chemical calculations, because they enable the calculation of yields and other values when dealing with particles of different mass.
Number of particles is a more useful unit in chemistry than mass or weight, because reactions take place between atoms (for example, two hydrogen atoms and one oxygen atom make one molecule of water) that have very different weights (one oxygen atom weighs almost 16 times as much as a hydrogen atom). However, the raw numbers of atoms in a reaction are not convenient, because they are very large; for example, just one mL of water contains over 3 (or 30,000,000,000,000,000,000,000) molecules.
Note: all of the following are accurate to approximately one significant figure.
★ Given that the volume of a grain of sand is approximately 10-12 m3[6], and given that the area of the United States is about 1013 m2[7], it therefore follows that a mole of sand grains would cover the United States in approximately one centimeter of sand.
★ A human body contains very roughly one hundred trillion cells[8]; there are roughly six billion people on Earth; so the total number of human cells on the planet is approximately 100
★ 6=6, which is very close to one mole.
★ Since the Earth has a radius of about 6400 km[9], its volume is approximately 1021 m3. Since about 500 large grapefruit will fit in one cubic meter[10], it therefore follows that a mole of grapefruit would have approximately the same volume as the Earth.
★ Avogadro constant
★ List of particles
★ Chemistry
★ Einstein (unit)
★ Faraday (unit)
★ Physics
★ Stoichiometry
★ Mole Day
★ Molarity
★ CODATA
1. Official SI Unit definitions
2. http://www.iop.org/EJ/article/0026-1394/42/2/001/met5_2_001.pdf
3. The symbol amu for atomic mass unit was replaced by the symbol u (unified atomic mass unit) in 1961. Before 1961 the symbol amu stood for different masses in chemistry and physics.
4. http://www.iop.org/EJ/abstract/0026-1394/42/2/001/
5. http://www.americanscientist.org/template/AssetDetail/assetid/54773
6. http://www.ingentaconnect.com/content/ap/ec/1999/00000048/00000005/art00470
7. http://www.daml.org/2001/12/factbook/us.html
8. A. S. Naidu, W. R. Bidlack, R. A. Clemens, "Probiotic Spectra of Lactic Acid Bacteria (LAB)", ''Critical Reviews in Food Science and Nutrition'', Volume 39, Number 1 / January 1999
9. http://scienceworld.wolfram.com/astronomy/EarthRadius.html
10. http://www.ams.usda.gov/standards/grpfrtfl.pdf
The 'mole' (symbol: mol) is the SI base unit that measures an amount of substance. One mole contains Avogadro's number (approximately 6.022) entities. A mole is much like "a dozen" in that both are absolute numbers (having no units) and can describe any type of elementary object, although the mole's use is usually limited to measurement of subatomic, atomic, and molecular structures.
| Contents |
| Definitions |
| Elementary entities |
| History |
| Proposed future definition |
| Utility of moles |
| Moles of everyday entities |
| See also |
| References |
Definitions
A mole is the amount of substance of a system which contains as many elementary entities as there are atoms in 0.012 kilogram (or 12 grams) of carbon-12, where the carbon-12 atoms are unbound, at rest and in their ground state.[1] The number of atoms in 0.012 kilogram of carbon-12 is known as the Avogadro constant, and is determined empirically. The currently accepted value is 6.02214179(30) mol-1 (2007 CODATA).
According to the SI, the mole is not dimensionless, but has its very own dimension, namely "amount of substance", comparable to other dimensions such as mass and luminous intensity. The International System of Units (SI), , , , International Bureau of Weights and Measures, 2006, (By contrast, the SI specifically defines the radian and the steradian as special names for the dimensionless unit one.) The International System of Units (SI), , , , International Bureau of Weights and Measures, 2006, The SI additionally defines the Avogadro constant as having the unit reciprocal mole, as it is the ratio of a dimensionless quantity and a quantity with the unit mole. However, if in the future the kilogram is redefined in terms of a specific number of carbon-12 atoms (see below), then the value of Avogadro's number will be defined rather than measured, and the mole will cease to be a unit of physical significance.[2]
The relationship of the atomic mass unit (u[3])
to Avogadro's number means that a mole can also be defined as: ''That quantity of a substance whose mass in grams is the same as its formula weight.'' For example, iron has an relative atomic mass of 55.845 u, so a mole of iron has a mass of 55.845 grams. This notation is very commonly used by chemists and physicists.
Scientists and engineers (chemical engineers in particular) sometimes measure amount of substance in units of ''gram-moles'', ''kilogram-moles'', ''pound-moles'', or ''ounce-moles''; these measure the quantity of a substance whose mass in grams, kilograms, pounds, or ounces (respectively) is equal to its formula weight. The SI mole is identical to the gram-mole.
Elementary entities
When the mole is used to specify the amount of a substance, the kind of elementary entities (particles) in the substance must be identified. The particles can be atoms, molecules, ions, formula units, electrons, photons or other particles. For example, one mole of water is equivalent to 18.016 grams of water and contains one mole of H2O molecules, but three moles of atoms (two moles H and one mole O).
When the substance of interest is a gas, the particles are usually molecules. However, the noble gases (He, Ar, Ne, Kr, Xe, Rn) are all monoatomic, that is each particle of gas is a single atom. All gases have the same molar volume of 22.4 litres per mole at STP (see Avogadro's Law).
A mole of atoms or molecules is also called a "gram atom" or "gram molecule", respectively.
History
The name ''mole'' (German ''Mol'') is attributed to Wilhelm Ostwald who introduced the concept in the year 1902. It is an abbreviation for ''molecule'' (German ''Molekül''), which is in turn derived from Latin ''moles'' "mass, massive structure".
He used it to express the gram molecular weight of a substance. So, for example, 1 mole of hydrochloric acid (HCl) has a mass of 36.5 grams (atomic masses Cl: 35.5 u, H: 1.0 u).
Prior to 1959 both the IUPAP and IUPAC used oxygen to define the mole, the chemists defining the mole as the number of atoms of oxygen which had mass 16 g, the physicists using a similar definition but with the oxygen-16 isotope only. The two organizations agreed in 1959/1960 to define the mole as such:
:''The mole is the amount of substance of a system which contains as many elementary entities as there are atoms in 0.012 kilogram of carbon-12; its symbol is "mol."''
This was adopted by the CIPM (International Committee for Weights and Measures) in 1967, and in 1971 it was adopted by the 14th CGPM (General Conference on Weights and Measures).
In 1980 the CIPM clarified the above definition, defining that the carbon-12 atoms are unbound and in their ground state.
Proposed future definition
As with other SI base units, there have been proposals to redefine the kilogram in such a way as to define some presently measured physical constants to fixed values. One proposed definition of the kilogram is:
:''The kilogram is the mass of exactly'' (6.0221415/0.012) ''unbound carbon-12 atoms at rest and in their ground state.'' [4]
This would have the effect of defining Avogadro's number to be precisely ''N''A = 6.0221415 elementary entities per mole, and, consequently, the mole would become merely a unit of counting, like the dozen.
Another proposed definition of ''NA'' is:
: ''NA'' = 602214141070409084099072 = 844468883
This has the nice properties of being a perfect cube, and being near the current experimental bounds of measurement.[5]
Utility of moles
The mole is useful in chemistry because it allows different substances to be measured in a comparable way. Using the same number of moles of two substances, both amounts have the same number of molecules or atoms. The mole makes it easier to interpret chemical equations in practical terms. Thus the equation:
:2H2 + O2 → 2H2O
can be understood as "two moles of hydrogen plus one mole of oxygen yields two moles of water."
Moles are useful in chemical calculations, because they enable the calculation of yields and other values when dealing with particles of different mass.
Number of particles is a more useful unit in chemistry than mass or weight, because reactions take place between atoms (for example, two hydrogen atoms and one oxygen atom make one molecule of water) that have very different weights (one oxygen atom weighs almost 16 times as much as a hydrogen atom). However, the raw numbers of atoms in a reaction are not convenient, because they are very large; for example, just one mL of water contains over 3 (or 30,000,000,000,000,000,000,000) molecules.
Moles of everyday entities
Note: all of the following are accurate to approximately one significant figure.
★ Given that the volume of a grain of sand is approximately 10-12 m3[6], and given that the area of the United States is about 1013 m2[7], it therefore follows that a mole of sand grains would cover the United States in approximately one centimeter of sand.
★ A human body contains very roughly one hundred trillion cells[8]; there are roughly six billion people on Earth; so the total number of human cells on the planet is approximately 100
★ 6=6, which is very close to one mole.
★ Since the Earth has a radius of about 6400 km[9], its volume is approximately 1021 m3. Since about 500 large grapefruit will fit in one cubic meter[10], it therefore follows that a mole of grapefruit would have approximately the same volume as the Earth.
See also
★ Avogadro constant
★ List of particles
★ Chemistry
★ Einstein (unit)
★ Faraday (unit)
★ Physics
★ Stoichiometry
★ Mole Day
★ Molarity
★ CODATA
References
1. Official SI Unit definitions
2. http://www.iop.org/EJ/article/0026-1394/42/2/001/met5_2_001.pdf
3. The symbol amu for atomic mass unit was replaced by the symbol u (unified atomic mass unit) in 1961. Before 1961 the symbol amu stood for different masses in chemistry and physics.
4. http://www.iop.org/EJ/abstract/0026-1394/42/2/001/
5. http://www.americanscientist.org/template/AssetDetail/assetid/54773
6. http://www.ingentaconnect.com/content/ap/ec/1999/00000048/00000005/art00470
7. http://www.daml.org/2001/12/factbook/us.html
8. A. S. Naidu, W. R. Bidlack, R. A. Clemens, "Probiotic Spectra of Lactic Acid Bacteria (LAB)", ''Critical Reviews in Food Science and Nutrition'', Volume 39, Number 1 / January 1999
9. http://scienceworld.wolfram.com/astronomy/EarthRadius.html
10. http://www.ams.usda.gov/standards/grpfrtfl.pdf
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