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Amount of Substance
In chemistry, the amount of substance in a given sample of matter is the number of discrete atomic-scale particles in it divided by the Avogadro constant. The particles or entities may be molecules, atoms, ions, electrons, or other, depending on the context. The value of the Avogadro constant NA has been set to . The amount of substance is sometimes referred to as the chemical amount.
The mole (symbol "mol") is a unit of the International System of Units defined (since 2019) by fixing the Avogadro constant at the given value. Historically, the mole was defined as the amount of substance in 12 grams of the carbon-12 isotope. As a consequence, the mass of one mole of a chemical compound, in grams, is numerically equal (for all practical purposes) to the mass of one molecule of the compound, in daltons, and the molar mass of an isotope in grams per mole is equal to the mass number. For example, a molecule of water has a mass of about 18.015 daltons on average, whereas a mole of water (which contains water molecules) has a total mass of about 18.015 grams.
In chemistry, because of the law of multiple proportions, it is often much more convenient to work with amounts of substances (that is, number of moles or of molecules) than with masses (grams) or volumes (liters). For example, the chemical fact "1 molecule of oxygen will react with 2 molecules of hydrogen to make 2 molecules of water " can also be stated as "1 mole of will react with 2 moles of to form 2 moles of water". The same chemical fact, expressed in terms of masses, would be "32 g of oxygen will react with approximately 2.0156 g of hydrogen to make approximately 18.0152 g of water" (and the numbers would depend on the isotopic composition of the reagents). In terms of volume, the numbers would depend on the pressure and temperature of the reagents and products. For the same reasons, the concentrations of reagents and products in liquids are often specified in moles per liter, rather than grams per liter.
The amount of substance is also a convenient concept in thermodynamics. For example, the pressure of a certain quantity of a noble gas in a recipient of a given volume, at a given temperature, is directly related to the number of molecules in the gas (through the ideal gas law), not to its mass.
This technical sense of the term "amount of substance" should not be confused with the general sense of "amount" in the English language. The latter may refer to other measurements such as mass or volume, rather than the number of particles. There are proposals to replace "amount of substance" with more easily-distinguishable terms, such as enplethy and stoichiometric amount.
The IUPAC recommends that "amount of substance" should be used instead of "number of moles", just as the quantity mass should not be called "number of kilograms".
Nature of the particles
To avoid ambiguity, the nature of the particles should be specified in any measurement of the amount of substance: thus, 1 mol of molecules of oxygen is about 32 grams, whereas 1 mol of atoms of oxygen is about 16 grams.
For salts (ionic solids) and for polymers of indeterminate molecular size, the term usually refers to the number of instances of the conventional chemical formula of the substance. Thus, for example, 1 mol of calcium chloride is understood to contain one mole of calciumcations and two moles of chlorideanions , even though the ions are not bound into separate three-atom molecules. Likewise, 1 mol of solid silicon dioxiden, which has a three-dimensional covalent lattice structure, is understood to contain one mole of silicon atoms and two moles of oxygen atoms.
For substances that normally exist in partially dissociated or polymerized form, the amount of substance usually refers to its nominal formula, without taking those changes into account. For example, "a solution of 1 mol of formaldehyde in water" is generally understood to contain one mole of carbon atoms, even though some of the formaldehyde may be in the form of methanediol or oligomers like paraformaldehyde and metaformaldehyde . Note that, because of the latter, the number of carbon-containing molecules in the solution may be substantially less than one mole.
Molar quantities (per mole)
The quotient of some extensive physical quantity of a homogeneous sample by its amount of substance is an intensive property of the substance, usually named by the prefix molar.
For example, the ratio of the mass of a sample by its amount of substance is the molar mass, whose SI unit is kilograms (or, more usually, grams) per mole; which is about 18.015 g/mol for water, and 55.845 g/mol for iron. From the volume, one gets the molar volume, which is about 17.962 milliliter/mol for liquid water and 7.092 mL/mol for iron at room temperature. From the heat capacity, one gets the molar heat capacity, which is about 75.385 J/K/mol for water and about 25.10 J/K/mol for iron.
Amount concentration (moles per liter)
Another important derived quantity is the amount of substance concentration (also called amount concentration, or substance concentration in clinical chemistry; which is defined as the amount of a specific substance in a sample of a solution (or some other mixture), divided by the volume of the sample.
The SI unit of this quantity is the mole (of the substance) per liter (of the solution). Thus, for example, the amount concentration of sodium chloride in ocean water is typically about 0.599 mol/L.
It must be noted that the denominator is the volume of the solution, not of the solvent. Thus, for example, one liter of standard vodka contains about 0.40 L of ethanol (315 g, 6.85 mol) and 0.60 L of water. The amount concentration of ethanol is therefore (6.85 mol of ethanol)/(1 L of vodka) = 6.85 mol/L, not (6.85 mol of ethanol)/(0.60 L of water), which would be 11.4 mol/L.
In chemistry, it is customary to read the unit "mol/L" as molar, and denote it by the symbol "M" (both following the numeric value). Thus, for example, each liter of a "0.5 molar" or "0.5 M" solution of urea in water contains 0.5 moles of that molecule. By extension, the amount concentration is also commonly called the molarity of the substance of interest in the solution. However, as of May 2007, these terms and symbols are not condoned by IUPAC.
This quantity should not be confused with the mass concentration, which is the mass of the substance of interest divided by the volume of the solution (about 35 g/L for sodium chloride in ocean water).
Amount fraction (moles per mole)
Confusingly, the amount concentration, or "molarity", should also be distinguished from "molar concentration", which should be the number of moles (molecules) of the substance of interest divided by the total number of moles (molecules) in the solution sample. This quantity is more properly called the amount fraction.
The alchemists, and especially the early metallurgists, probably had some notion of amount of substance, but there are no surviving records of any generalization of the idea beyond a set of recipes. In 1758, Mikhail Lomonosov questioned the idea that mass was the only measure of the quantity of matter, but he did so only in relation to his theories on gravitation. The development of the concept of amount of substance was coincidental with, and vital to, the birth of modern chemistry.
1777: Wenzel publishes Lessons on Affinity, in which he demonstrates that the proportions of the "base component" and the "acid component" (cation and anion in modern terminology) remain the same during reactions between two neutral salts.
1792: Richter publishes the first volume of Stoichiometry or the Art of Measuring the Chemical Elements (publication of subsequent volumes continues until 1802). The term "stoichiometry" is used for the first time. The first tables of equivalent weights are published for acid-base reactions. Richter also notes that, for a given acid, the equivalent mass of the acid is proportional to the mass of oxygen in the base.
1805: Dalton publishes his first paper on modern atomic theory, including a "Table of the relative weights of the ultimate particles of gaseous and other bodies".
The concept of atoms raised the question of their weight. While many were skeptical about the reality of atoms, chemists quickly found atomic weights to be an invaluable tool in expressing stoichiometric relationships.
1808: Publication of Dalton's A New System of Chemical Philosophy, containing the first table of atomic weights (based on H = 1).
1811: Avogadro hypothesizes that equal volumes of different gases (at same temperature and pressure) contain equal numbers of particles, now known as Avogadro's law.
1813/1814: Berzelius publishes the first of several tables of atomic weights based on the scale of O = 100.
1815: Prout publishes his hypothesis that all atomic weights are integer multiple of the atomic weight of hydrogen. The hypothesis is later abandoned given the observed atomic weight of chlorine (approx. 35.5 relative to hydrogen).
The ideal gas law was the first to be discovered of many relationships between the number of atoms or molecules in a system and other physical properties of the system, apart from its mass. However, this was not sufficient to convince all scientists of the existence of atoms and molecules, many considered it simply being a useful tool for calculation.
1834: Faraday states his Laws of electrolysis, in particular that "the chemical decomposing action of a current is constant for a constant quantity of electricity".
1887: Arrhenius describes the dissociation of electrolyte in solution, resolving one of the problems in the study of colligative properties.
1893: First recorded use of the term mole to describe a unit of amount of substance by Ostwald in a university textbook.
1897: First recorded use of the term mole in English.
By the turn of the twentieth century, the concept of atomic and molecular entities was generally accepted, but many questions remained, not least the size of atoms and their number in a given sample. The concurrent development of mass spectrometry, starting in 1886, supported the concept of atomic and molecular mass and provided a tool of direct relative measurement.
1905: Einstein's paper on Brownian motion dispels any last doubts on the physical reality of atoms, and opens the way for an accurate determination of their mass.
^Avogadro, Amedeo (1811). "Essai d'une maniere de determiner les masses relatives des molecules elementaires des corps, et les proportions selon lesquelles elles entrent dans ces combinaisons". Journal de Physique. 73: 58-76.English translation.
^Berzelius' first atomic weight measurements were published in Swedish in 1810: Hisinger, W.; Berzelius, J.J. (1810). "Forsok rorande de bestamda proportioner, havari den oorganiska naturens bestandsdelar finnas forenada". Afh. Fys., Kemi Mineral. 3: 162.