Updated March 04, 2020
By Riti Gupta
Reviewed by: Lana Bandoim, B.S.
Chemical solutions have a concentration associated with them. The concentration allows you to know how much of a substance you are working with. Concentration is often reported in molarity (M), which is defined as mole of solute per liter solution (mol/L).
While molarity is usually a convenient way to talk about concentration, sometimes there aren't actually that many moles in the solution.
What does that mean? Well, say you have a 0.001 M solution of NaOH. It can be kind of annoying to say 0.001 M. Instead, scientists will move the decimal point to the right three places. This makes the new number 1. But the units are no longer moles/L, but rather millimoles/L or mmol/L. These units are called millimolar or mM.
While in many cases scientists prefer to talk about concentration in terms of molarity, it can sometimes be more useful to talk about mass per liter solvent. For example, this could be grams per liter (g/L) or milligrams/L (mg/L). Below, you will see an example as to why this is the case.
Given that there are so many different ways to talk about concentration, it is necessary to have a way to convert between them. Below you will see examples of how to convert from mmol/L to mg/L and visa versa.
In order to convert from milligrams to mmol, you will need to make use of the molar mass of the substance in question. Since the molar mass is given in grams per mol (g/mol) you will also need to take into account that there is a conversion factor from milligrams to grams (there are 1,000 milligrams in 1 gram).
Take a look at the following example:
Say you have a 259 milligram per liter solution of NaOH (259 mg/L) and you want to see what the mmol/L or mM concentration is. You need to do some dimensional analysis given the fact that there are 1,000 milligrams in a gram and the molar mass of NaOH (39.997 g/mol).
First, convert the milligrams to grams:
Now, the units are g/L, and you can use the molar mass to convert to mol/L:
Finally, you can convert from moles to millimoles. There are 1,000 millimoles in 1 mole.
A one liter solution that contains 259 milligrams of NaOH is thus a 6.48 mM solution.
This entire dimensional analysis can be done in one step as well:
Thus, a 259 mg/L solution of NaOH is the same as a 6.48 mM solution of NaOH.
Say you have a 10 mM solution of MgCl2. What is the concentration in mg/L? In other words, how many milligrams of the salt would you need to weigh out and add to a liter of solution?
First, begin by converting millimolar to molar:
Next, you can use the molar mass to find how many grams of MgCl2 there are in 0.010 moles or MgCl2.
Finally, you can convert the number of grams to milligrams:
This indicates that in a 10 mM solution of MgCl2 is the same as a 962 mg/L solution of MgCl2.
You can see why talking about the concentration in millimolar is a bit easier: It's a nice round number. On the other hand, the concentration of 962 mg/L tells you exactly how much salt to weigh out when making the solution.
This is a prime example of why knowing how to convert from mmol/L to mg/L (and visa versa) can be helpful in the lab.
Enter the total number of mmol and the total number of L into the calculator to convert into mg/dl.
Mmol/L to mg/dl Formula
The following formula is used to convert mmol/L to mg/dl.
mg/dl = 18 * mmol/l
- Where mg/dl is the milligrams per deciliter
- mmol/l is the millimoles per liter
How do you convert mmol/L to mg/dL? In order to convert milligrams per deciliter multiply the total amount of mmol/L by 18. This will convert the value into mg/dL.
How do you convert mg/dL to mmol/L? To convert the mg/dL into mmol/L, simply divide the total mg/dL by 18. The formula for this process is shown above.
Example Problem
how to calculate mg/dl from mmol/l?
First, determine the total number of millimoles per liter.
For this example, the total number of mmol/l is found to be 150.
Next, use the formula for the conversion above to determine the amount of mg/dl.
mg/dl = 18 * mmol/l
mg/dl = 18 * 150
mg/dl = 2700
"Nmol" redirects here. For the mathematical technique, see Method of lines. The mole, symbol mol, is the SI base unit of amount of substance.[1][2][3] The quantity amount of substance is a measure of how many elementary entities of a given substance are in an object or sample. Depending on what the substance is, an elementary entity may be an atom, a molecule, an ion, an ion pair, or a subatomic particle such as an electron. For example, if beaker A contains 10 moles of water (a chemical compound) and beaker B contains 10 moles of mercury (a chemical element), they contain equal amounts of substance and beaker B contains exactly 1 atom of mercury for each molecule of water in beaker A, despite the two beakers' containing different volumes and very different masses of liquid.
The mole is defined as exactly 6.02214076×1023 elementary entities.[clarification needed] This definition was adopted in November 2018 and came into force on May 20, 2019, superseding the previous definition of a mole as a number of elementary entities equal to that of 12 grams of carbon-12, the most common isotope of carbon. Because a dalton, a unit commonly used to measure atomic mass, is exactly 1/12 of the mass of a carbon-12 atom, the definition of the mole in use before 2019 entailed that the mass of one mole of a compound or element in grams was numerically equal to the average mass of one molecule or atom of the substance in daltons, and that the number of daltons in a gram was equal to the number of elementary entities in a mole. Because the mass of a nucleon (i.e. a proton or neutron) is approximately 1 dalton and the nucleons in an atom's nucleus make up the overwhelming majority of its mass, the pre-2019 definition also entailed that the mass of one mole of a substance was roughly equivalent to the number of nucleons in one atom or molecule of that substance. For example, a water molecule formed from the most common isotope of oxygen and of hydrogen contains 10 protons plus 8 neutrons for a total mass of 18.015 daltons, and a mole of water has a mass of 18.015 grams.
The number of elementary entities in 1 mole is known as the Avogadro number. Prior to 2019, it could only be estimated based on experimental data. The value 6.02214076×1023 was adopted based on the best estimates available in 2018, allowing the new definition to very closely approximate the earlier definition and avoid the need to recalibrate measuring equipment or update published data tables.
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 2H2 + O2 → 2H2O can be interpreted to mean that for each 2 mol dihydrogen (H2) and 1 mol dioxygen (O2) that react, 2 mol of water (H2O) form. The mole may also be used to measure the amount 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 in mole(s) per unit volume of solution, for which the unit typically used is moles per litre (mol/L), commonly abbreviated M.
The term gram-molecule (g mol) was formerly used for "mole of molecules",[4] and gram-atom (g atom) for "mole of atoms". For example, 1 mole of MgBr2 is 1 gram-molecule of MgBr2 but 3 gram-atoms of MgBr2.[5][6] The mole is essentially a count of particles.[7] 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.
The mass of 1 mole of a substance is equal to its relative atomic or molecular mass in grams.
The molar mass of a substance is the ratio of the mass of a sample of that substance to its amount of substance. The amount of substance is the number of moles in the sample.
For most practical purposes, the numerical value of the molar mass expressed with the unit gram per mole is the same as that of the mean mass of one molecule of the substance expressed with the unit dalton. For example, the molar mass of water is 18.015 g/mol.[8] Other methods include the use of the molar volume or the measurement of electric charge.[8] 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.[8] 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 molar mass of calcium-40 is 39.96259098±0.00000022 g/mol, whereas the molar mass of calcium-42 is 41.95861801±0.00000027 g/mol, and of calcium with the normal isotopic mix is 40.078±0.004 g/mol.
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 (mol/L) are used.
The molar fraction or mole fraction of a substance in a mixture (such as a solution) 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 NaCl is dissolved in 100 g of water, the amounts of the two substances in the solution will be (20 g)/(58.443 g/mol) = 0.34221 mol and (100 g)/(18.015 g/mol) = 5.5509 mol, respectively; and the molar fraction of NaCl will be 0.34221/(0.34221 + 5.5509) = 0.05807.
In a mixture of gases, the partial pressure of each component is proportional to its molar ratio.
Molar mass
Molar concentration
Molar fraction
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 (atomic mass) was published by John Dalton (1766–1844) 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 (an unproven hypothesis at the time) to make practical use of the tables. This would lead to some confusion between atomic masses (promoted by proponents of atomic theory) and equivalent weights (promoted by its opponents and which sometimes differed from relative atomic masses by an integer factor), which would last throughout much of the nineteenth century.
Jöns Jacob Berzelius (1779–1848) 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 (1816–56), Henri Victor Regnault (1810–78) and Stanislao Cannizzaro (1826–1910) 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 (1860). 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 (molecule).[9][10][11] The related concept of equivalent mass had been in use at least a century earlier.[12]
Standardization
Developments in mass spectrometry led to the adoption of oxygen-16 as the standard substance, in lieu of natural oxygen.[citation needed]
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.[4][7] The four different definitions were equivalent to within 1%.
| Atomic mass of hydrogen = 1 | 1.00794(7) | −0.788% |
| Atomic mass of oxygen = 16 | 15.9994(3) | +0.00375% |
| Relative atomic mass of 16O = 16 | 15.9949146221(15) | +0.0318% |
Since the definition of the gram was not mathematically tied to that of the dalton, the number of molecules per mole NA (the Avogadro constant) had to be determined experimentally. The experimental value adopted by CODATA in 2010 is NA = (6.02214129±0.00000027)×1023 mol−1.[13] In 2011 the measurement was refined to (6.02214078±0.00000018)×1023 mol−1.[14]
The mole was made the seventh SI base unit in 1971 by the 14th CGPM.[15]
2019 redefinition of SI base units
In 2011, the 24th meeting of the General Conference on Weights and Measures (CGPM) 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.[2]
Such changes officially came into effect on 20 May 2019. Following such changes, "one mole" of a substance was redefined as containing "exactly 6.02214076×1023 elementary entities" of that substance.[16][17]
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;[7][18]
- The SI thermodynamic mole is irrelevant to analytical chemistry and could cause avoidable costs to advanced economies[19]
- The mole is not a true metric (i.e. measuring) unit, rather it is a parametric unit, and amount of substance is a parametric base quantity[20]
- the SI defines numbers of entities as quantities of dimension one, and thus ignores the ontological distinction between entities and units of continuous quantities[21]
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 that 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, the most notable one being 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.[citation needed]
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 (1000 mol) in industrial-scaled processes, the numerical value of molarity remains the same.
For convenience in avoiding conversions in the imperial (or American customary units), some engineers adopted the pound-mole (notation 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,[22] 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 (notation kg-mol), which is defined as the number of entities in 12 kg of 12C, and often referred to the mole as the gram-mole (notation g-mol), when dealing with laboratory data.[22]
Late 20th-century chemical engineering practice came to use the kilomole (kmol), 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 = 6.02×1023 photons.[23]
October 23, denoted 10/23 in the US, is recognized by some as Mole Day.[24] It is an informal holiday in honor of the unit among chemists. The date is derived from the Avogadro number, which is approximately 6.022×1023. It starts at 6:02 a.m. and ends at 6:02 p.m. Alternatively, some chemists celebrate June 2 (06/02), June 22 (6/22), or 6 February (06.02), a reference to the 6.02 or 6.022 part of the constant.[25][26][27]
- Einstein (unit)
- Element-reactant-product table
- Faraday (unit)
- Mole fraction – Proportion of a constituent in a mixture
- Dalton (unit) – Standard unit of mass for atomic-scale chemical species
- Molecular mass – Mass of a given molecule in daltons
- Molar mass – Mass per amount of substance
- ^ IUPAC Gold Book. "IUPAC - mole (M03980)". International Union of Pure and Applied Chemistry. doi:10.1351/goldbook.M03980. S2CID 241546445. {{cite journal}}: Cite journal requires |journal= (help)
- ^ a b "On the revision of the International System of Units - International Union of Pure and Applied Chemistry". IUPAC | International Union of Pure and Applied Chemistry. 16 November 2018. Retrieved 1 March 2021.
- ^ BIPM (20 May 2019). "Mise en pratique for the definition of the mole in the SI". BIPM.org. Retrieved 18 February 2022.
- ^ a b International Bureau of Weights and Measures (2006), The International System of Units (SI) (PDF) (8th ed.), pp. 114–15, ISBN 92-822-2213-6, archived (PDF) from the original on 2021-06-04, retrieved 2021-12-16
- ^ Wang, Yuxing; Bouquet, Frédéric; Sheikin, Ilya; Toulemonde, Pierre; Revaz, Bernard; Eisterer, Michael; Weber, Harald W.; Hinderer, Joerg; Junod, Alain; et al. (2003). "Specific heat of MgB2 after irradiation". Journal of Physics: Condensed Matter. 15 (6): 883–893. arXiv:cond-mat/0208169. Bibcode:2003JPCM...15..883W. doi:10.1088/0953-8984/15/6/315. S2CID 16981008.
- ^ Lortz, R.; Wang, Y.; Abe, S.; Meingast, C.; Paderno, Yu.; Filippov, V.; Junod, A.; et al. (2005). "Specific heat, magnetic susceptibility, resistivity and thermal expansion of the superconductor ZrB12". Phys. Rev. B. 72 (2): 024547. arXiv:cond-mat/0502193. Bibcode:2005PhRvB..72b4547L. doi:10.1103/PhysRevB.72.024547. S2CID 38571250.
- ^ a b c de Bièvre, Paul; Peiser, H. Steffen (1992). "'Atomic Weight' — The Name, Its History, Definition, and Units" (PDF). Pure and Applied Chemistry. 64 (10): 1535–43. doi:10.1351/pac199264101535.
- ^ a b c International Bureau of Weights and Measures. Realising the mole Archived 2008-08-29 at the Wayback Machine. Retrieved 25 September 2008.
- ^ Helm, Georg (1897). "The Principles of Mathematical Chemistry: The Energetics of Chemical Phenomena". transl. by Livingston, J.; Morgan, R. New York: Wiley: 6. {{cite journal}}: Cite journal requires |journal= (help)
- ^ Some sources place the date of first usage in English as 1902. Merriam–Webster proposes Archived 2011-11-02 at the Wayback Machine an etymology from Molekulärgewicht (molecular weight).
- ^ Ostwald, Wilhelm (1893). Hand- und Hilfsbuch zur Ausführung Physiko-Chemischer Messungen [Handbook and Auxiliary Book for Conducting Physical-Chemical Measurements]. Leipzig, Germany: Wilhelm Engelmann. p. 119. From p. 119: "Nennen wir allgemein das Gewicht in Grammen, welches dem Molekulargewicht eines gegebenen Stoffes numerisch gleich ist, ein Mol, so ... " (If we call in general the weight in grams, which is numerically equal to the molecular weight of a given substance, a "mol", then ... )
- ^ mole, n.8, Oxford English Dictionary, Draft Revision Dec. 2008
- ^ physics.nist.gov/ Archived 2015-06-29 at the Wayback Machine Fundamental Physical Constants: Avogadro Constant
- ^ Andreas, Birk; et al. (2011). "Determination of the Avogadro Constant by Counting the Atoms in a 28Si Crystal". Physical Review Letters. 106 (3): 30801. arXiv:1010.2317. Bibcode:2011PhRvL.106c0801A. doi:10.1103/PhysRevLett.106.030801. PMID 21405263. S2CID 18291648.
- ^ "BIPM – Resolution 3 of the 14th CGPM". www.bipm.org. Archived from the original on 9 October 2017. Retrieved 1 May 2018.
- ^ CIPM Report of 106th Meeting Archived 2018-01-27 at the Wayback Machine Retrieved 7 April 2018
- ^ "Redefining the Mole". NIST. NIST. 2018-10-23. Retrieved 24 October 2018.
- ^ Barański, Andrzej (2012). "The Atomic Mass Unit, the Avogadro Constant, and the Mole: A Way to Understanding". Journal of Chemical Education. 89 (1): 97–102. Bibcode:2012JChEd..89...97B. doi:10.1021/ed2001957.
- ^ Price, Gary (2010). "Failures of the global measurement system. Part 1: the case of chemistry". Accreditation and Quality Assurance. 15 (7): 421–427. doi:10.1007/s00769-010-0655-z. S2CID 95388009.
- ^ Johansson, Ingvar (2010). "Metrological thinking needs the notions of parametric quantities, units, and dimensions". Metrologia. 47 (3): 219–230. Bibcode:2010Metro..47..219J. doi:10.1088/0026-1394/47/3/012.
- ^ Cooper, G.; Humphry, S. (2010). "The ontological distinction between units and entities". Synthese. 187 (2): 393–401. doi:10.1007/s11229-010-9832-1. S2CID 46532636.
- ^ a b Himmelblau, David (1996). Basic Principles and Calculations in Chemical Engineering (6 ed.). pp. 17–20. ISBN 978-0-13-305798-0.
- ^ "Lighting Radiation Conversion". Archived from the original on March 11, 2016. Retrieved March 10, 2016.
- ^ History of National Mole Day Foundation, Inc. Archived 2010-10-23 at the Wayback Machine.
- ^ Happy Mole Day! Archived 2014-07-29 at the Wayback Machine, Mary Bigelow. SciLinks blog, National Science Teachers Association. October 17, 2013.
- ^ What Is Mole Day? – Date and How to Celebrate. Archived 2014-07-30 at Wikiwix, Anne Marie Helmenstine. About.com.
- ^ The Perse School (Feb 7, 2013), The Perse School celebrates moles of the chemical variety, Cambridge Network, archived from the original on 2015-02-11, retrieved Feb 11, 2015, As 6.02 corresponds to 6th February, the School has adopted the date as their 'Mole Day'.
- ChemTeam: The Origin of the Word 'Mole' at the Wayback Machine (archived December 22, 2007)
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