Mole (unit): Wikis

  

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The mole (symbol mol) is the SI base unit[1] of amount of substance; one of a few units used to measure this physical quantity. The name "mole" is an 1897 translation[2][3] of the German Mol, coined by Wilhelm Ostwald in 1893,[4] although the related concept of equivalent mass had been in use at least a century earlier. The name is assumed[5] to be derived from the German word Molekül (molecule). The names gram-atom and gram-molecule have also been used in the same sense as "mole",[1][6] but these names are now obsolete.

The mole is definied as the amount of substance of a system that contains as many "elementary entities" (e.g. atoms, molecules, ions, electrons) as there are atoms in 12 g of carbon-12 (12C).[1] A mole has 6.0221415×1023[7] atoms or molecules of the pure substance being measured. A mole of a substance has mass in grams exactly equal to the substance's molecular or atomic weight. That is to say, a substance's atomic or molecular mass in atomic mass units is the same as its molar mass in grams. Because of this, one can measure the number of moles in a pure substance by weighing it and comparing the result to its molecular or atomic weight.

The current definition of the mole was approved during the 1960s:[1][6] Prior to that, there had been definitions based on the atomic weight of hydrogen (about one gram of hydrogen-1 gas, excluding its heavy isotopes), the atomic weight of oxygen, and the relative atomic mass of oxygen-16: the four different definitions are equivalent to within 1%.

The most common method of measuring an amount of substance is to measure its mass and then to divide by the molar mass of the substance.[8] Molar masses may be easily calculated from tabulated values of atomic weights and the molar mass constant (which has a convenient defined value of 1 g/mol). Other methods include the use of the molar volume or the measurement of electric charge.[8]

Contents

The mole as a unit

Since its adoption into the International System of Units, there have been a number of criticisms of the concept of the mole being a unit like the metre or the second.[6] These criticisms may be briefly summarised as:

  • amount of substance is not a true physical quantity (or dimension), and is redundant to mass, so should not have its own base unit;
  • the mole is simply a shorthand way of referring to a large number.

In chemistry, it has been known since Proust's Law of definite proportions (1794) that knowledge of the mass of each of the components in a chemical system is not sufficient to define the system. Amount of substance can be described as mass divided by Proust's "definite proportions", and contains information which is missing from the measurement of mass alone. As demonstrated by Dalton's Law of partial pressures (1803), a measurement of mass is not even necessary to measure the amount of substance (although in practice it is usual). There are many physical relationships between amount of substance and other physical quantities, most notably the ideal gas law (where the relationship was first demonstrated in 1857). The term "mole" was first used in a textbook describing these colligative properties.

The second misconception, that the mole is simply a counting aid, has even found its way into elementary chemistry textbooks.[9] These books and others often contend that the mole is defined in terms of the Avogadro constant, rather than the other way around, and so is equal to 6.0221415×1023 elementary entities.

Consider the measurement of one mole of silicon. As silicon is a solid at room temperature, the convenient method of measurement is weighing. By consulting published tables, it can easily be found that the atomic weight of silicon is 28.0855.[10] Multiplying by the molar mass constant Mu gives the molar mass in any desired mass units: assuming the measurement is to be made in grams, Mu = 1 g/mol, and so the molar mass of silicon is 28.0855 g/mol. Hence, 28.0855 g of silicon is equivalent to one mole of silicon, without the Avogadro constant ever having come into play.

Counting (or calculating) the number of atoms in 28.0855 g of silicon is one way of determining the Avogadro constant, NA, and a way which is currently receiving a lot of attention (see below) although, as of the 2006 CODATA values of the physical constants, it is not the most accurate. It is only a method of determining NA because it is known by other means that 28.0855 g of silicon is equivalent to one mole. Those other means are:

  • the very accurate determination of the ratios of the masses of each of the three silicon nuclides to the mass of an atom of carbon-12, in such a way that it is known that a silicon-28 atom is [27.976 926 5327(20)/12] times as massive as a carbon-12 atom;[10][11]
  • the determination of the isotopic abundance of silicon in the samples used to make the measurements, allowing the calculation of the atomic weight of silicon in each individual sample;
  • the definition of 12 g of carbon-12 atoms to be equivalent to one mole.

History

The first table of atomic weights was published by John Dalton (1766–1844) in 1805, based on a system in which the atomic weight of hydrogen was defined as 1. These atomic weights were based on the stoichiometric proportions of chemical reactions and compounds, a fact which greatly aided their acceptance: it was not necessary for a chemist to subscribe to atomic theory (an unproven hypothesis at the time) to make practical use of the tables. This would lead to some confusion between atomic weights (promoted by proponents of atomic theory) and equivalent weights (promoted by its opponents and which sometimes differed from atomic weights by an integer factor), which would last throughout much of the nineteenth century.

Jöns Jacob Berzelius (1779–1848) was instrumental in the determination of atomic weights to ever increasing accuracy. He was also the first chemist to use oxygen as the standard to which other weights were referred. Oxygen is a useful standard, as, unlike hydrogen, it forms compounds with most other elements, especially metals. However he chose to fix the atomic weight of oxygen as 100, an innovation which did not catch on.

Charles Frédéric Gerhardt (1816–56), Henri Victor Regnault (1810–78) and Stanislao Cannizzaro (1826–1910) expanded on Berzelius' work, resolving many of the problems of unknown stoichiometry of compounds, and the use of atomic weights attracted a large consensus by the time of the Karlsruhe Congress (1860). The convention had reverted to defining the atomic weight of hydrogen as 1, although at the level of precision of measurements at that time—relative uncertainties of around 1%—this was numerically equivalent to the later standard of oxygen = 16. However the chemical convenience of having oxygen as the primary atomic weight standard became ever more evident with advances in analytical chemistry and the need for ever more accurate atomic weight determinations.

Scale basis Scale basis
relative to 12C = 12
Relative deviation
from the 12C = 12 scale
Atomic weight of hydrogen = 1 1.007 94(7) −0.788%
Atomic weight of oxygen = 16 15.9994(3) +37.5 ppm
Relative atomic mass of 16O = 16 15.994 914 6221(15) +318 ppm

Other units called "mole"

Chemical engineers use the concept extensively, but the unit is rather small for industrial use. For convenience in avoiding conversions, some American engineers adopted the pound-mole (noted lb-mol or lbmol), which is defined as the number of entities in 12 lb of 12C. One lb-mol is equal to 453.59237 mol.[12] Chemical engineers also use * moles( star moles) to denote two amounts of mass(or more) that has the same latent heat. This allows the assumption of equal-mole overflow to hold true in distillation columns.[citation needed]

In the metric system, chemical engineers once used the kilogram-mole (noted kg-mol), which is defined as the number of entities in 12 kg of 12-C, and often referred to the mole as the gram-mole (noted g-mol), when dealing with laboratory data.[12] However modern chemical engineering practice is to use the kilomole (kmol), which is identical to the kilogram-mole, but whose name and symbol adopt the SI convention for standard multiples of metric units.

Proposed future definition

Kilogram

As with other SI base units, there have been proposals to redefine the kilogram in such a way as to define some currently measured physical constants to fixed values. One proposed definition of the kilogram is:[13]

The kilogram is the mass of exactly (6.0221415×10230.012) unbound carbon-12 atoms at rest and in their ground state.

This would have the effect of defining the Avogadro constant to be precisely 6.0221415×1023 elementary entities per mole.

Holiday

October the 23rd (10/23), "Mole Day", is an informal holiday in honour of the unit among chemists in North America. The date is derived from the Avogadro constant, which is approximately 6.02×1023.

See also

References

  1. ^ a b c d International Bureau of Weights and Measures (2006), The International System of Units (SI) (8th ed.), pp. 114–15, ISBN 92-822-2213-6, http://www.bipm.org/utils/common/pdf/si_brochure_8_en.pdf 
  2. ^ Helm, Georg (1897), The Principles of Mathematical Chemistry: The Energetics of Chemical Phenomena, transl. by Livingston, J.; Morgan, R., New York: Wiley, p. 6 
  3. ^ Some sources place the date of first usage in English as 1902. Merriam–Webster proposes an etymology from Molekulärgewicht (molecular weight).
  4. ^ Ostwald, Wilhelm (1893). Hand- und Hilfsbuch zur ausführung physiko-chemischer Messungen. Leipzig. p.  119. 
  5. ^ mole, n.8, Oxford English Dictionary, Draft Revision Dec. 2008
  6. ^ a b c de Bièvre, P.; Peiser, H.S. (1992), "'Atomic Weight'—The Name, Its History, Definition, and Units", Pure Appl. Chem. 64 (10): 1535–43, doi:10.1351/pac199264101535, http://www.iupac.org/publications/pac/1992/pdf/6410x1535.pdf 
  7. ^ Mohr, Peter J.; Taylor, Barry N.; Newell, David B. (2008). "CODATA Recommended Values of the Fundamental Physical Constants: 2006". Rev. Mod. Phys. 80: 633–730. doi:10.1103/RevModPhys.80.633. http://physics.nist.gov/cuu/Constants/codata.pdf.  Direct link to value..
  8. ^ a b International Bureau of Weights and Measures. "Realising the mole." Retrieved 25 September 2008.
  9. ^ Kotz, John C.; Treichel, Paul M.; Townsend, John R. (2008), Chemistry and Chemical Reactivity (7th ed.), Brooks/Cole, ISBN 0495387037, http://cengagesites.com/academic/kotz.cfm?site=2719&section=home 
  10. ^ a b National Institute of Standards and Technology. "Atomic Weights and Isotopic Compositions for All Elements". http://physics.nist.gov/cgi-bin/Compositions/stand_alone.pl?ele=&ascii=html&isotype=some. Retrieved 2008-09-25. 
  11. ^ It should be emphasised that relative atomic masses are measured as ratios of the masses of two nuclides. They cannot be measured (at least not to this level of accuracy) as absolute values of the mass of each nuclide in yoctograms.
  12. ^ a b Himmelblau, David (1996). Basic Principles and Calculations in Chemical Engineering (6 ed.). p. 17–20. ISBN 0-13-305798-4. 
  13. ^ Mills, Ian M.; Mohr, Peter J.; Quinn, Terry J.; Taylor, Barry N.; Williams, Edwin R. (2005). "Redefinition of the kilogram: a decision whose time has come". Metrologia 42: 71–80. doi:10.1088/0026-1394/42/2/001. http://www.iop.org/EJ/article/0026-1394/42/2/001/met5_2_001.pdf.  Abstract.


Simple English

The article about the animal is at Mole

Mole is the SI unit of measurement used to measure the number of things, usually atoms or molecules. One mole of something is equal to 6.0221415×10²³ of it. So, "One mole of hydrogen atoms" means 6.0221415×10²³ hydrogen atoms. "One mole of grapefruits" means 6.0221415×10²³ grapefruits. We call this number Avogadro's number. We use this number because it is the number of carbon atoms in 12 grams of carbon-12, which is the most common kind of carbon. We can measure anything in moles, but it is not very useful for most things because the numbers are so big. For example, one mole of grapefruits would be as big as the earth.

Because different molecules and atoms do not have the same mass, one mole of one thing does not weigh the same as one mole of something else. Atoms and molecule mass is measured in u. One u is equal to one gram per mole. This means that if an atom has a mass of one u, one mole of this atom weighs one gram.

Math with the Mole

Moles = mass (g) / Relative mass (grams per mole) Example: How many moles are there in 20 grams of hydrogen? A value of 1 can be used for hydrogen's relative mass, although the correct value is slightly larger. So: moles = mass/relative mass = 20/1 = 20 moles.

Moles = concentration (mol/dm3) x volume (dm3) Example: How many moles are there in 100cm3 of 0.1M H2SO4? 1 dm3 is the same as 1000 cm3, so the value in cubic centimetres needs to be divided by 1000. 100/1000 x 0.1 = 0.01 moles.

A methane molecule is made from one carbon atom and four hydrogen atoms. Carbon has a mass of 12.011 u and hydrogen has a mass of 1.008 u. This means that the mass of one methane molecule is 12.011 u + (4*1.008u), or 16.043 u. This means that one mole of methane has a mass of 16.043 grams.

A mole can be thought of as two bags of different sized balls. One bag contains tennis balls and the other footballs. There are the same amount of balls in both bags but the mass of the footballs is much larger.It is a different way to measure things. Moles measure the number of particles, not the mass. So both bags contain three moles.

A mole is simply a unit of the number of things. Units are invented when existing units can not describe something well enough. Chemical reactions often take place at levels where using grams wouldn't make sense, yet using absolute numbers of atoms/molecules/ions would be confusing, too.



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