# Molar mass

Last updated
Molar mass
Common symbols
M
SI unit kg/mol
Other units
g/mol

In chemistry, the molar mass of a chemical compound is defined as the mass of a sample of that compound divided by the amount of substance in that sample, measured in moles. [1] The molar mass is a bulk, not molecular, property of a substance. The molar mass is an average of many instances of the compound, which often vary in mass due to the presence of isotopes. Most commonly, the molar mass is computed from the standard atomic weights and is thus a terrestrial average and a function of the relative abundance of the isotopes of the constituent atoms on Earth. The molar mass is appropriate for converting between the mass of a substance and the amount of a substance for bulk quantities.

## Contents

The molecular weight is very commonly used as a synonym of molar mass, particularly for molecular compounds; however, the most authoritative sources define it differently (see molecular mass).

The formula weight is a synonym of molar mass that is frequently used for non-molecular compounds, such as ionic salts.

The molar mass is an intensive property of the substance, that does not depend on the size of the sample. In the International System of Units (SI), the base unit of molar mass is kg/mol. However, for historical reasons, molar masses are almost always expressed in g/mol.

The mole was defined in such as way that the molar mass of a compound, in g/mol, is numerically equal (for all practical purposes) to the average mass of one molecule, in daltons. Thus, for example, the average mass of a molecule of water is about 18.0153 daltons, and the molar mass of water is about 18.0153 g/mol.

For chemical elements without isolated molecules, such as carbon and metals, the molar mass is computed dividing by the number of moles of atoms instead. Thus, for example, the molar mass of iron is about 55.845 g/mol.

Between 1971 and 2019, SI defined the "amount of substance" as a separate dimension of measurement, and the mole was defined as the amount of substance that has as many constituent particles as there are atoms in 12 grams of carbon-12. In that period, the molar mass of carbon-12 was thus exactly 12 g/mol, by definition. Since 2019, a mole of any substance has been redefined in the SI as the amount of that substance containing an exactly defined number of particles, N = 6.02214076×1023. Therefore, the molar mass of a compound now is simply the mass of this number of molecules of the compound.

## Molar masses of elements

Before reading this section, it must be understood that the 2019 redefinition of the SI base units concluded that the molar mass constant is not exactly 1×10−3 kg/mol, but Mu = 0.99999999965(30)×10−3 kg⋅mol−1. [2]

The molar mass of atoms of an element is given by the relative atomic mass of the element multiplied by the molar mass constant, Mu 1.000000×10−3 kg/mol = 1.000000 g/mol. [3] For normal samples from earth with typical isotope composition, the atomic weight can be approximated by the standard atomic weight [4] or the conventional atomic weight.

M(H) = 1.00797(7) × 1.000000 g/mol = 1.00797(7) g/mol
M(S) = 32.065(5) × 1.000000 g/mol = 32.065(5) g/mol
M(Cl) = 35.453(2) × 1.000000 g/mol = 35.453(2) g/mol
M(Fe) = 55.845(2) × 1.000000 g/mol = 55.845(2) g/mol.

Multiplying by the molar mass constant ensures that the calculation is dimensionally correct: standard relative atomic masses are dimensionless quantities (i.e., pure numbers) whereas molar masses have units (in this case, grams/mole).

Some elements are usually encountered as molecules, e.g. hydrogen (H
2
), sulfur (S
8
), chlorine (Cl
2
). The molar mass of molecules of these elements is the molar mass of the atoms multiplied by the number of atoms in each molecule:

M(H
2
) = 2 × 1.007 97(7) × 1.000000 g/mol = 2.01588(14) g/mol
M(S
8
) = 8 × 32.065(5) × 1.000000 g/mol = 256.52(4) g/mol
M(Cl
2
) = 2 × 35.453(2) × 1.000000 g/mol = 70.906(4) g/mol.

## Molar masses of compounds

The molar mass of a compound is given by the sum of the relative atomic mass A
r
of the atoms which form the compound multiplied by the molar mass constant M
u
:

${\displaystyle M=M_{\rm {u}}M_{\rm {r}}=M_{\rm {u}}\sum _{i}{A_{\rm {r}}}_{i}.}$

Here, M
r
is the relative molar mass, also called formula weight. For normal samples from earth with typical isotope composition, the standard atomic weight or the conventional atomic weight can be used as an approximation of the relative atomic mass of the sample. Examples are:

M(NaCl) = [22.98976928(2) + 35.453(2)] × 1.000000 g/mol = 58.443(2) g/mol
M(C
12
H
22
O
11
) = ([12 × 12.0107(8)] + [22 × 1.00794(7)] + [11 × 15.9994(3)]) × 1.000000 g/mol = 342.297(14) g/mol.

An average molar mass may be defined for mixtures of compounds. [1] This is particularly important in polymer science, where different polymer molecules may contain different numbers of monomer units (non-uniform polymers). [5] [6]

## Average molar mass of mixtures

The average molar mass of mixtures ${\displaystyle {\bar {M}}}$ can be calculated from the mole fractions ${\displaystyle x_{i}}$ of the components and their molar masses ${\displaystyle M_{i}}$:

${\displaystyle {\bar {M}}=\sum _{i}x_{i}M_{i}.}$

It can also be calculated from the mass fractions ${\displaystyle w_{i}}$ of the components:

${\displaystyle {\frac {1}{\bar {M}}}=\sum _{i}{\frac {w_{i}}{M_{i}}}.}$

As an example, the average molar mass of dry air is 28.97 g/mol. [7]

Molar mass is closely related to the relative molar mass (M
r
) of a compound, to the older term formula weight (F.W.), and to the standard atomic masses of its constituent elements. However, it should be distinguished from the molecular mass (which is confusingly also sometimes known as molecular weight), which is the mass of one molecule (of any single isotopic composition) and is not directly related to the atomic mass, the mass of one atom (of any single isotope). The dalton, symbol Da, is also sometimes used as a unit of molar mass, especially in biochemistry, with the definition 1 Da = 1 g/mol, despite the fact that it is strictly a unit of mass (1 Da = 1 u = 1.66053906660(50)×10−27 kg, as of 2018 CODATA recommended values).

Gram atomic mass is another term for the mass, in grams, of one mole of atoms of that element. "Gram atom" is a former term for a mole.

Molecular weight (M.W.) is an older term for what is now more correctly called the relative molar mass (M
r
). [8] This is a dimensionless quantity (i.e., a pure number, without units) equal to the molar mass divided by the molar mass constant. [9]

### Molecular mass

The molecular mass (m) is the mass of a given molecule: it is usually measured in daltons (Da or u). [10] Different molecules of the same compound may have different molecular masses because they contain different isotopes of an element. This is distinct but related to the molar mass, which is a measure of the average molecular mass of all the molecules in a sample and is usually the more appropriate measure when dealing with macroscopic (weigh-able) quantities of a substance.

Molecular masses are calculated from the atomic masses of each nuclide, while molar masses are calculated from the standard atomic weights [11] of each element. The standard atomic weight takes into account the isotopic distribution of the element in a given sample (usually assumed to be "normal"). For example, water has a molar mass of 18.0153(3) g/mol, but individual water molecules have molecular masses which range between 18.0105646863(15) u (1H
2
16O) and 22.0277364(9) u (2H
2
18O).

The distinction between molar mass and molecular mass is important because relative molecular masses can be measured directly by mass spectrometry, often to a precision of a few parts per million. This is accurate enough to directly determine the chemical formula of a molecule. [12]

### DNA synthesis usage

The term formula weight (F.W.) has a specific meaning when used in the context of DNA synthesis: whereas an individual phosphoramidite nucleobase to be added to a DNA polymer has protecting groups and has its molecular weight quoted including these groups, the amount of molecular weight that is ultimately added by this nucleobase to a DNA polymer is referred to as the nucleobase's formula weight (i.e., the molecular weight of this nucleobase within the DNA polymer, minus protecting groups).

## Precision and uncertainties

The precision to which a molar mass is known depends on the precision of the atomic masses from which it was calculated, and value of the molar mass constant. Most atomic masses are known to a precision of at least one part in ten-thousand, often much better [4] (the atomic mass of lithium is a notable, and serious, [13] exception). This is adequate for almost all normal uses in chemistry: it is more precise than most chemical analyses, and exceeds the purity of most laboratory reagents.

The precision of atomic masses, and hence of molar masses, is limited by the knowledge of the isotopic distribution of the element. If a more accurate value of the molar mass is required, it is necessary to determine the isotopic distribution of the sample in question, which may be different from the standard distribution used to calculate the standard atomic mass. The isotopic distributions of the different elements in a sample are not necessarily independent of one another: for example, a sample which has been distilled will be enriched in the lighter isotopes of all the elements present. This complicates the calculation of the standard uncertainty in the molar mass.

A useful convention for normal laboratory work is to quote molar masses to two decimal places for all calculations. This is more accurate than is usually required, but avoids rounding errors during calculations. When the molar mass is greater than 1000 g/mol, it is rarely appropriate to use more than one decimal place. These conventions are followed in most tabulated values of molar masses. [14] [15]

## Measurement

Molar masses are almost never measured directly. They may be calculated from standard atomic masses, and are often listed in chemical catalogues and on safety data sheets (SDS). Molar masses typically vary between:

1–238 g/mol for atoms of naturally occurring elements;
10–1000 g/mol for simple chemical compounds;
1000–5000000 g/mol for polymers, proteins, DNA fragments, etc.

While molar masses are almost always, in practice, calculated from atomic weights, they can also be measured in certain cases. Such measurements are much less precise than modern mass spectrometric measurements of atomic weights and molecular masses, and are of mostly historical interest. All of the procedures rely on colligative properties, and any dissociation of the compound must be taken into account.

### Vapour density

The measurement of molar mass by vapour density relies on the principle, first enunciated by Amedeo Avogadro, that equal volumes of gases under identical conditions contain equal numbers of particles. This principle is included in the ideal gas equation:

${\displaystyle pV=nRT\ }$

where n is the amount of substance. The vapour density (ρ) is given by

${\displaystyle \rho ={{nM} \over {V}}.\ }$

Combining these two equations gives an expression for the molar mass in terms of the vapour density for conditions of known pressure and temperature.

${\displaystyle M={{RT\rho } \over {p}}\ }$

### Freezing-point depression

The freezing point of a solution is lower than that of the pure solvent, and the freezing-point depression (ΔT) is directly proportional to the amount concentration for dilute solutions. When the composition is expressed as a molality, the proportionality constant is known as the cryoscopic constant (K
f
) and is characteristic for each solvent. If w represents the mass fraction of the solute in solution, and assuming no dissociation of the solute, the molar mass is given by

${\displaystyle M={{wK_{f}} \over {\Delta T}}.\ }$

### Boiling-point elevation

The boiling point of a solution of an involatile solute is higher than that of the pure solvent, and the boiling-point elevation (ΔT) is directly proportional to the amount concentration for dilute solutions. When the composition is expressed as a molality, the proportionality constant is known as the ebullioscopic constant (K
b
) and is characteristic for each solvent. If w represents the mass fraction of the solute in solution, and assuming no dissociation of the solute, the molar mass is given by

${\displaystyle M={{wK_{b}} \over {\Delta T}}.\ }$

## Related Research Articles

The molecular mass (m) is the mass of a given molecule: it is measured in daltons. Different molecules of the same compound may have different molecular masses because they contain different isotopes of an element. The related quantity relative molecular mass, as defined by IUPAC, is the ratio of the mass of a molecule to the unified atomic mass unit and is unitless. The molecular mass and relative molecular mass are distinct from but related to the molar mass. The molar mass is defined as the mass of a given substance divided by the amount of a substance and is expressed in g/mol. The molar mass is usually the more appropriate figure when dealing with macroscopic (weigh-able) quantities of a substance.

Stoichiometry is the calculation of reactants and products in chemical reactions in chemistry.

The mole (symbol: mol) is the unit of measurement for amount of substance in the International System of Units (SI). A mole of a substance or a mole of particles is defined as exactly 6.02214076×1023 particles, which may be atoms, molecules, ions, or electrons. In short, for particles, 1 mol = 6.02214076×1023.

The Avogadro constant (NA or L) is the proportionality factor that relates the number of constituent particles (usually molecules, atoms or ions) in a sample with the amount of substance in that sample. Its SI unit is the reciprocal mole, and it is defined as NA = 6.02214076×1023 mol−1. It is named after the Italian scientist Amedeo Avogadro.

Lorenzo Romano Amedeo Carlo Avogadro, Count of Quaregna and Cerreto (, also, Italian: [ameˈdɛːo avoˈɡaːdro]; 9 August 1776 – 9 July 1856) was an Italian scientist, most noted for his contribution to molecular theory now known as Avogadro's law, which states that equal volumes of gases under the same conditions of temperature and pressure will contain equal numbers of molecules. In tribute to him, the number of elementary entities (atoms, molecules, ions or other particles) in 1 mole of a substance, 6.02214076×1023, is known as the Avogadro constant, one of the seven SI base units and represented by NA.

The dalton or unified atomic mass unit is a unit of mass widely used in physics and chemistry. It is defined as 1/12 of the mass of an unbound neutral atom of carbon-12 in its nuclear and electronic ground state and at rest. The atomic mass constant, denoted mu is defined identically, giving mu = m(12C)/12 = 1 Da.

The volumetric heat capacity of a material is the heat capacity of a sample of the substance divided by the volume of the sample. Informally, it is the amount of energy that must be added, in the form of heat, to one unit of volume of the material in order to cause an increase of one unit in its temperature. The SI unit of volumetric heat capacity is joule per kelvin per cubic meter, J/K/m3 or J/(K·m3).

Relative atomic mass or atomic weight is a dimensionless physical quantity defined as the ratio of the average mass of atoms of a chemical element in a given sample to the atomic mass constant. The atomic mass constant is defined as being 1/12 of the mass of a carbon-12 atom. Since both quantities in the ratio are masses, the resulting value is dimensionless; hence the value is said to be relative.

Avogadro's law is an experimental gas law relating the volume of a gas to the amount of substance of gas present. The law is a specific case of the ideal gas law. A modern statement is:

Avogadro's law states that "equal volumes of all gases, at the same temperature and pressure, have the same number of molecules."

For a given mass of an ideal gas, the volume and amount (moles) of the gas are directly proportional if the temperature and pressure are constant.

In chemistry, the amount of substance in a given sample of matter is defined as the number of discrete atomic-scale particles in it divided by the Avogadro constant NA. In a truly atomistic view, the amount of substance is simply the number of particles that constitute the substance. 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 defined as 6.02214076×1023 mol−1. In the truly atomistic view, 1 mol = 6.02214076×1023 particles (the Avogadro number) and therefore the conversion constant is simply NA = 1. The amount of substance is sometimes referred to as the chemical amount.

The molar heat capacity of a chemical substance is the amount of energy that must be added, in the form of heat, to one mole of the substance in order to cause an increase of one unit in its temperature. Alternatively, it is the heat capacity of a sample of the substance divided by the amount of substance of the sample; or also the specific heat capacity of the substance times its molar mass. The SI unit of specific heat is joule per kelvin per mole, J⋅K−1⋅mol−1.

Equivalent weight is the mass of one equivalent, that is the mass of a given substance which will combine with or displace a fixed quantity of another substance. The equivalent weight of an element is the mass which combines with or displaces 1.008 gram of hydrogen or 8.0 grams of oxygen or 35.5 grams of chlorine. These values correspond to the atomic weight divided by the usual valence; for oxygen as example that is 16.0g / 2.

In linear polymers the individual polymer chains rarely have exactly the same degree of polymerization and molar mass, and there is always a distribution around an average value. The molar mass distribution describes the relationship between the number of moles of each polymer species (Ni) and the molar mass (Mi) of that species. The molar mass distribution of a polymer may be modified by polymer fractionation.

Monoisotopic mass (Mmi) is one of several types of molecular masses used in mass spectrometry. The theoretical monoisotopic mass of a molecule is computed by taking the sum of the accurate masses of the primary isotope of each atom in the molecule. For small molecules made up of low atomic number elements the monoisotopic mass is observable as an isotopically pure peak in a mass spectrum. This differs from the nominal molecular mass, which is the sum of the mass number of the primary isotope of each atom in the molecule and is an integer. It also is different from the molar mass, which is a type of average mass. For some atoms like carbon, oxygen, hydrogen, nitrogen, and sulfur the Mmi of these elements is exactly the same as the mass of its natural isotope, which is the lightest one. However, this does not hold true for all atoms. Iron's most common isotope has a mass number of 56, while the stable isotopes of iron vary in mass number from 54 to 58. Monoisotopic mass is typically expressed in daltons (Da), also called unified atomic mass units (u).

The standard atomic weight of a chemical element is the weighted arithmetic mean of the relative isotopic masses of all isotopes of that element weighted by each isotope's abundance on Earth. The standard atomic weight of each chemical element is determined and published by the Commission on Isotopic Abundances and Atomic Weights (CIAAW) of the International Union of Pure and Applied Chemistry (IUPAC) based on natural, stable, terrestrial sources of the element. The standard atomic weight is the most common and practical atomic weight used by scientists.

The mass recorded by a mass spectrometer can refer to different physical quantities depending on the characteristics of the instrument and the manner in which the mass spectrum is displayed.

The molar mass constant, usually denoted by Mu, is a physical constant defined as the ratio of the molar mass of an element and its relative mass.

This glossary of chemistry terms is a list of terms and definitions relevant to chemistry, including chemical laws, diagrams and formulae, laboratory tools, glassware, and equipment. Chemistry is a physical science concerned with the composition, structure, and properties of matter, as well as the changes it undergoes during chemical reactions; it features an extensive vocabulary and a significant amount of jargon.

In chemistry, the mass fraction of a substance within a mixture is the ratio of the mass of that substance to the total mass of the mixture. Expressed as a formula, the mass fraction is:

The atomic mass is the mass of an atom. Although the SI unit of mass is kilogram, the atomic mass is often expressed in the non-SI unit dalton where 1 dalton is defined as ​112 of the mass of a single carbon-12 atom, at rest. The protons and neutrons of the nucleus account for nearly all of the total mass of atoms, with the electrons and nuclear binding energy making minor contributions. Thus, the numeric value of the atomic mass when expressed in daltons has nearly the same value as the mass number. Conversion between mass in kilograms and mass in daltons can be done using the atomic mass constant .

## References

1. International Union of Pure and Applied Chemistry (1993). Quantities, Units and Symbols in Physical Chemistry , 2nd edition, Oxford: Blackwell Science. ISBN   0-632-03583-8 . p. 41. Electronic version.
2. "2018 CODATA Value: molar mass constant". The NIST Reference on Constants, Units, and Uncertainty. NIST. 20 May 2019. Retrieved 2019-05-20.Cite has empty unknown parameter: |month= (help)
3. Mohr, Peter J.; Taylor, Barry N.; Newell, David B. (2011). "CODATA Recommended Values of the Fundamental Physical Constants: 2010". Database developed by J. Baker, M. Douma, and S. Kotochigova. National Institute of Standards and Technology, Gaithersburg, MD 20899.
4. Wieser, M. E. (2006), "Atomic Weights of the Elements 2005" (PDF), Pure and Applied Chemistry , 78 (11): 2051–66, doi:
5. "International union of pure and applied chemistry, commission on macromolecular nomenclature, note on the terminology for molar masses in polymer science". Journal of Polymer Science: Polymer Letters Edition. 22 (1): 57. 1984. Bibcode:1984JPoSL..22...57.. doi:10.1002/pol.1984.130220116.
6. Metanomski, W. V. (1991). Compendium of Macromolecular Nomenclature. Oxford: Blackwell Science. pp. 47–73. ISBN   0-632-02847-5.
7. The Engineering ToolBox Molecular Mass of Air
8. IUPAC , Compendium of Chemical Terminology , 2nd ed. (the "Gold Book") (1997). Online corrected version:  (2006) " relative molar mass ". doi : 10.1351/goldbook.R05270
9. The technical definition is that the relative molar mass is the molar mass measured on a scale where the molar mass of unbound carbon 12 atoms, at rest and in their electronic ground state, is 12. The simpler definition given here is equivalent to the full definition because of the way the molar mass constant is itself defined.
10. International Bureau of Weights and Measures (2006), The International System of Units (SI) (PDF) (8th ed.), p. 126, ISBN   92-822-2213-6, archived (PDF) from the original on 2017-08-14
11. "Atomic Weights and Isotopic Compositions for All Elements". NIST . Retrieved 2007-10-14.
12. "Author Guidelines – Article Layout". RSC Publishing . Retrieved 2007-10-14.
13. Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. p. 21. ISBN   978-0-08-037941-8.
14. See, e.g., Weast, R. C., ed. (1972). Handbook of Chemistry and Physics (53rd ed.). Cleveland, OH: Chemical Rubber Co.
15. Possolo, Antonio; van der Veen, Adriaan M. H.; Meija, Juris; Hibbert, D. Brynn (2018-01-04). "Interpreting and propagating the uncertainty of the standard atomic weights (IUPAC Technical Report)". Pure and Applied Chemistry. 90 (2): 395–424. doi:10.1515/pac-2016-0402.