Mole (unit)
The mole is the unit of measurement for amount of substance in the International System of Units. A mole of a substance or a mole of particles is defined as exactly particles, which may be atoms, molecules, ions, or electrons. In short, for particles 1 mol =.
The definition was adopted in November 2018 as one of the seven SI base units, revising the previous definition that specified it as the number of atoms in 12 grams of carbon-12, an isotope of carbon.
The number was chosen so that the mass of one mole of a chemical compound, in grams, is numerically equal to the average mass of one molecule of the compound, in daltons. Thus, for example, one mole of water contains molecules, whose total mass is about 18.015 grams – and the mean mass of one molecule of water is about 18.015 daltons.
The mole is widely used in chemistry as a convenient way to express amounts of reactants and products of chemical reactions. For example, the chemical equation can be interpreted to mean that 2 mol dihydrogen and 1 mol dioxygen react to form 2 mol water. The mole may also be used to represent the number of atoms, ions, electrons, or other entities. The concentration of a solution is commonly expressed by its molarity, defined as the amount of dissolved substance per unit volume of solution, for which the unit typically used is moles per litre, commonly abbreviated M.
The term gram-molecule was formerly used for "mole of molecules", and gram-atom for "mole of atoms". For example, 1 mole of MgBr2 is 1 gram-molecule of MgBr2 but 3 gram-atoms of MgBr2.
Concepts
Nature of the particles
The mole is essentially a count of particles. Usually the particles counted are chemically identical entities, individually distinct. For example, a solution may contain a certain number of dissolved molecules that are more or less independent of each other. However, in a solid the constituent particles are fixed and bound in a lattice arrangement, yet they may be separable without losing their chemical identity. Thus the solid is composed of a certain number of moles of such particles. In yet other cases, such as diamond, where the entire crystal is essentially a single molecule, the mole is still used to express the number of atoms bound together, rather than a count of multiple molecules. Thus, common chemical conventions apply to the definition of the constituent particles of a substance, in other cases exact definitions may be specified.Molar mass
The molar mass of a substance is the mass of 1 mole of that substance, in multiples of the gram. The amount of substance is the number of moles in the sample.For most practical purposes, the magnitude of molar mass is numerically the same as that of the mean mass of one molecule, expressed in daltons. For example, the molar mass of water is 18.015 g/mol. Other methods include the use of the molar volume or the measurement of electric charge.
The number of moles of a substance in a sample is obtained by dividing the mass of the sample by the molar mass of the compound. For example, 100 g of water is about 5.551 mol of water.
The molar mass of a substance depends not only on its molecular formula, but also on the distribution of isotopes of each chemical element present in it. For example, the mass of one mole of calcium-40 is, whereas the mass of one mole of calcium-42 is, and of one mole of calcium with the normal isotopic mix is.
Molar concentration
The molar concentration, also called molarity, of a solution of some substance is the number of moles per unit of volume of the final solution. In the SI its standard unit is mol/m3, although more practical units, such as mole per litre are used.Molar fraction
The molar fraction or mole fraction of a substance in a mixture is the number of moles of the compound in one sample of the mixture, divided by the total number of moles of all components. For example, if 20 g of is dissolved in 100 g of water, the amounts of the two substances in the solution will be / = 0.34221 mol and / = 5.5509 mol, respectively; and the molar fraction of will be 0.34221/ = 0.05807.In a mixture of gases, the partial pressure of each component is proportional to its molar ratio.
History
The history of the mole is intertwined with that of molecular mass, atomic mass units, and the Avogadro number.The first table of standard atomic weight was published by John Dalton in 1805, based on a system in which the relative atomic mass of hydrogen was defined as 1. These relative atomic masses were based on the stoichiometric proportions of chemical reaction and compounds, a fact that greatly aided their acceptance: It was not necessary for a chemist to subscribe to atomic theory to make practical use of the tables. This would lead to some confusion between atomic masses and equivalent weights, which would last throughout much of the nineteenth century.
Jöns Jacob Berzelius was instrumental in the determination of relative atomic masses to ever-increasing accuracy. He was also the first chemist to use oxygen as the standard to which other masses 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 mass of oxygen as 100, which did not catch on.
Charles Frédéric Gerhardt, Henri Victor Regnault and Stanislao Cannizzaro expanded on Berzelius' works, resolving many of the problems of unknown stoichiometry of compounds, and the use of atomic masses attracted a large consensus by the time of the Karlsruhe Congress. The convention had reverted to defining the atomic mass 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 mass standard became ever more evident with advances in analytical chemistry and the need for ever more accurate atomic mass determinations.
The name mole is an 1897 translation of the German unit Mol, coined by the chemist Wilhelm Ostwald in 1894 from the German word Molekül. However, the related concept of equivalent mass had been in use at least a century earlier.
Standardization
Developments in mass spectrometry led to the adoption of oxygen-16 as the standard substance, in lieu of natural oxygen.The oxygen-16 definition was replaced with one based on carbon-12 during the 1960s. The mole was defined by International Bureau of Weights and Measures as "the amount of substance of a system which contains as many elementary entities as there are atoms in 0.012 kilogram of carbon-12." Thus, by that definition, one mole of pure 12C had a mass of exactly 12 g. The four different definitions were equivalent to within 1%.
| Scale basis | Scale basis relative to C = 12 | Relative deviation from the C = 12 scale |
| Atomic mass of hydrogen = 1 | 1.00794 | −0.788% |
| Atomic mass of oxygen = 16 | 15.9994 | +0.00375% |
| Relative atomic mass of O = 16 | 15.9949146221 | +0.0318% |
Since the definition of the gram was not mathematically tied to that of the dalton, the number of molecules per mole NA had to be determined experimentally. The experimental value adopted by CODATA in 2010 is.
In 2011 the measurement was refined to.
The mole was made the seventh SI base unit in 1971 by the 14th CGPM.
2019 redefinition of SI base units
In 2011, the 24th meeting of the General Conference on Weights and Measures agreed to a plan for a possible revision of the SI base unit definitions at an undetermined date.On 16 November 2018, after a meeting of scientists from more than 60 countries at the CGPM in Versailles, France, all SI base units were defined in terms of physical constants. This meant that each SI unit, including the mole, would not be defined in terms of any physical objects but rather they would be defined by constants that are, in their nature, exact.
Such changes officially came into effect on 20 May 2019. Following such changes, "one mole" of a substance was redefined as "exactly elementary entities" of that substance.
Criticism
Since its adoption into the International System of Units in 1971, numerous criticisms of the concept of the mole as a unit like the metre or the second have arisen:- the number of molecules, etc. in a given amount of material is a fixed dimensionless quantity that can be expressed simply as a number, not requiring a distinct base unit;
- the official mole is based on an outdated continuum concept of matter;
- the SI thermodynamic mole is irrelevant to analytical chemistry and could cause avoidable costs to advanced economies;
- the mole is not a true metric unit, rather it is a parametric unit, and amount of substance is a parametric base quantity;
- the SI defines numbers of entities as quantities of dimension one, and thus ignores the ontological distinction between entities and units of continuous quantities.
Similar units
Like chemists, chemical engineers use the unit mole extensively, but different unit multiples may be more suitable for industrial use. For example, the SI unit for volume is the cubic metre, a much larger unit than the commonly used litre in the chemical laboratory. When amount of substance is also expressed in kmol in industrial-scaled processes, the numerical value of molarity remains the same.For convenience in avoiding conversions in the imperial, some engineers adopted the pound-mole, which is defined as the number of entities in 12 lb of 12C. One lb-mol is equal to, which value is the same as the number of grams in an international avoirdupois pound.
In the metric system, chemical engineers once used the kilogram-mole, which is defined as the number of entities in 12 kg of 12C, and often referred to the mole as the gram-mole, when dealing with laboratory data.
Late 20th-century chemical engineering practice came to use the kilomole, which is numerically identical to the kilogram-mole, but whose name and symbol adopt the SI convention for standard multiples of metric units – thus, kmol means 1000 mol. This is equivalent to the use of kg instead of g. The use of kmol is not only for "magnitude convenience" but also makes the equations used for modelling chemical engineering systems coherent. For example, the conversion of a flowrate of kg/s to kmol/s only requires the molecular mass without the factor 1000 unless the basic SI unit of mol/s were to be used.
Greenhouse and growth chamber lighting for plants is sometimes expressed in micromoles per square metre per second, where 1 mol photons = photons.