Molality, also called molal concentration, is a measure of the concentration of a solute in a solution in terms of amount of substance in a specified amount of mass of the solvent. This contrasts with the definition of molarity which is based on a specified volume of solution.
A commonly used unit for molality in chemistry is mol/kg. A solution of concentration 1 mol/kg is also sometimes denoted as 1 molal.
Contents

Definition 1

Origin 2

Unit 3

Usage considerations 4

Relation to other compositional properties 5

Mass fraction 5.1

Mole fraction 5.2

Molar concentration (Molarity) 5.3

Mass concentration 5.4

Equal ratios 5.5

Example of conversion 5.6

Osmolality 5.7

Relation to apparent (molar) properties 6

References 7
Definition
The molality (b), of a solution is defined as the amount of substance (in mol) of solute, n_{solute}, divided by the mass (in kg) of the solvent, m_{solvent}:^{[1]}

b = \frac{n_{solute}}{m_{solvent}}
Origin
The term molality is formed in analogy to molarity which is the molar concentration of a solution. The earliest of the intensive property molality and of its adjectival unit, the nowdeprecated molal, appears to have been published by G. N. Lewis and M. Randall in the 1923 publication of Thermodynamics and the Free Energies of Chemical Substances.^{[2]} Though the two terms are subject to being confused with one another, the molality and molarity of a weak aqueous solution are nearly the same, as one kilogram of water (solvent) occupies the volume of 1 liter at room temperature and a small amount of solute has little effect on the volume.
Unit
The SI unit for molality is mol/kg.
A solution with a molality of 3 mol/kg is often described as "3 molal" or "3 m". However, following the SI system of units, the National Institute of Standards and Technology, the United States authority on measurement, considers the term "molal" and the unit symbol "m" to be obsolete, and suggests mol/kg or a related unit of the SI.^{[3]} This recommendation has not been universally implemented in academia yet.
Usage considerations

Advantages
Compared to molar concentration or mass concentration, the preparation of a solution of a given molality requires only a good scale: both solvent and solute need to be weighed, as opposed to measured volumetrically, which would be subject to variations in density due to the ambient conditions of temperature and pressure; this is an advantage because, in chemical compositions, the mass, or the amount, of a pure known substance is more relevant than its volume: a contained measured amount of substance may change in volume with ambient conditions, but its amount and mass are unvarying, and chemical reactions occur in proportions of mass, not volume. The massbased nature of molality implies that it can be readily converted into a mass ratio (or mass fraction, "w," ratio),

b\,M_{solute}=\frac{m_{solute}}{m_{solvent}}=\frac{w_{solute}}{w_{solvent}},
where the symbol M stands for molar mass in kg/mole, or into a mole ratio (or mole fraction, "x," ratio)

b\,M_{solvent}=\frac{n_{solute}}{n_{solvent}}=\frac{x_{solute}}{x_{solvent}}.
The advantage of molality over other massbased fractions is the fact that the molality of one solute in a singlesolvent solution is independent of the presence or absence of other solutes.

Problem areas
Unlike all the other compositional properties listed in "Relation" section (below), molality depends on the choice of the substance to be called “solvent” in an arbitrary mixture. If there is only one pure liquid substance in a mixture, the choice is clear, but not all solutions are this clearcut: in an alcoholwater solution, either one could be called the solvent; in an alloy, or solid solution, there is no clear choice and all constituents may be treated alike. In such situations, mass or mole fraction is the preferred compositional specification.
Relation to other compositional properties
In what follows, the solvent may be given the same treatment as the other constituents of the solution, such that the molality of the solvent of an nsolute solution, say b_{0}, is found to be nothing more than the reciprocal of its molar mass, M_{0}:

b_0=\frac=M_0^{1}.
Mass fraction
The conversions to and from the mass fraction, w, of the solute in a singlesolute solution are

w=(1+(b\,M)^{1})^{1},\ b=\frac{w}{(1w)M} ,
where b is the molality and M is the molar mass of the solute.
More generally, for an nsolute/onesolvent solution, letting b_{i} and w_{i} be, respectively, the molality and mass fraction of the ith solute,

w_i=w_0\,b_i M_i,\ b_i=\frac{w_i}{w_0 M_i},
where M_{i} is the molar mass of the ith solute, and w_{}0 is the mass fraction of the solvent, which is expressible both as a function of the molalities as well as a function of the other mass fractions,

w_0=\left(1+\sum_{j=1}^{n}{b_j M_j}\right)^{1}=1\sum_{j=1}^{n}{w_j}.
Mole fraction
The conversions to and from the mole fraction, x, of the solute in a singlesolute solution are

x=(1+(M_0\,b)^{1})^{1},\ b=\frac{x}{M_0(1x)},
where M_{0} is the molar mass of the solvent.
More generally, for an nsolute/onesolvent solution, letting x_{i} be the mole fraction of the ith solute,

x_i=x_0 M_0\,b_i,\ b_i=\frac{b_0 x_i}{x_0},
where x_{0} is the mole fraction of the solvent, expressible both as a function of the molalities as well as a function of the other mole fractions:

x_0=\left(1+M_0\sum_{j=1}^{n}{b_j}\right)^{1}=1\sum_{j=1}^{n}{x_j}.
Molar concentration (Molarity)
The conversions to and from the molar concentration, c, for onesolute solutions are

c = \frac{\rho\, b}{1+ b M},\ b=\frac{c}{\rhocM},
where ρ is the mass density of the solution, b is the molality, and M is the molar mass of the solute.
For solutions with n solutes, the conversions are

c_i =c_0 M_0\,b_i,\ b_i=\frac{b_0 c_i}{c_0},
where the molar concentration of the solvent c_{0} is expressible both as a function of the molalities as well as a function of the molarities:

c_0=\frac{\rho\,b_0}{1+\displaystyle\sum_{j=1}^{n}{b_j M_j}}=\frac{\rho\displaystyle\sum_{j=1}^{n}{c_i M_i}}{M_0}.
Mass concentration
The conversions to and from the mass concentration, ρ_{solute}, of a singlesolute solution are

\rho_{solute} = \frac{\rho\,b\,M}{1+b M},\ b=\frac{\rho_{solute}}{M (\rho\rho_{solute})},
where ρ is the mass density of the solution, b is the molality, and M is the molar mass of the solute.
For the general nsolute solution, the mass concentration of the ith solute, ρ_{i}, is related to its molality, b_{i}, as follows:

\rho_i = \rho_0\,b_i M_i,\ b_i=\frac{\rho_i}{\rho_0 M_i},
where the mass concentration of the solvent, ρ_{0}, is expressible both as a function of the molalities as well as a function of the mass concentrations:

\rho_0=\frac{\rho}{1+\displaystyle\sum_{j=1}^n b_j M_j}=\rho\sum_{j=1}^{n}{\rho_i}.
Equal ratios
Alternatively, we may use just the last two equations given for the compositional property of the solvent in each of the preceding sections, together with the relationships given below, to derive the remainder of properties in that set:

\frac{b_i}{b_j}=\frac{x_i}{x_j}=\frac{c_i}{c_j}=\frac{\rho_i\,M_j}{\rho_j\,M_i}=\frac{w_i\,M_j}{w_j\,M_i},
where i and j are subscripts representing all the constituents, the n solutes plus the solvent.
Example of conversion
An acid mixture consists of 0.76/0.04/0.20 mass fractions of (70% HNO_{3}) / (49% HF) /(H_{2}O), where the percentages refer to mass fractions of the bottled acids carrying a balance of H_{2}O. The first step is determining the mass fractions of the constituents:

\begin{align} w_{HNO_3}&=0.70\times 0.76=0.532\\ w_{HF}&=0.49\times 0.04=0.0196\\ w_{H_2O}&=1w_{HNO_3}w_{HF}=0.448\\ \end{align}
The approximate molar masses in kg/mol are

M_{HNO_3}=0.063,\ M_{HF}=0.020,\ M_{H_2O}=0.018.
First derive the molality of the solvent, in mol/kg,

b_{H_2O}=(M_{H_2O})^{1}=\frac{1}{0.018},
and use that to derive all the others by use of the equal ratios:

\frac{b_{HNO_3}}{b_{H_2O}}=\frac{w_{HNO_3}M_{H_2O}}{w_{H_2O}M_{HNO_3}} \ \therefore b_{HNO_3}=18.83.
Actually, b_{H2O} cancels out, because it is not needed. In this case, there is a more direct equation: we use it to derive the molality of HF:

b_{HF}=\frac{w_{HF}}{w_{H_2O}M_{HF}}=2.19.
The mole fractions may be derived from this result:

\begin{align} &x_{H_2O}=(1+M_{H_2O}(b_{HNO_3}+b_{HF}))^{1}=0.726,\\ &\frac{x_{HNO_3}}{x_{H_2O}}=\frac{b_{HNO_3}}{b_{H_2O}}\ \therefore x_{HNO_3}=0.246,\\ &x_{HF}=1x_{HNO_3}x_{H_2O}=0.029. \end{align}
Osmolality
Osmolality is a variation of molality that takes into account only solutes that contribute to a solution's osmotic pressure. It is measured in osmoles of the solute per kilogram of water. This unit is frequently used in medical laboratory results in place of osmolarity, because it can be measured simply by depression of the freezing point of a solution, or cryoscopy (see also: osmostat and colligative properties).
Relation to apparent (molar) properties
Molality appears in the expression of the apparent (molar) volume of a solute as a function of the molality b of that solute (and density of the solution and solvent):

{}^\phi\tilde{V}_1 = \frac{1}{b}( \frac{1}{\rho}  \frac{1}{\rho_0^0}) + \frac{M_1}{\rho}
For multicomponent systems the relation is slightly modified by the sum of molalities of solutes.

{}^\phi\tilde{V}_i = \frac{1}{\sum b_j}( \frac{1}{\rho}  \frac{1}{\rho_0^0}) + \frac{\sum b_j M_j}{\sum b_j \rho}
References

^ IUPAC, Compendium of Chemical Terminology, 2nd ed. (the "Gold Book") (1997). Online corrected version: (2006–) "molality".

^ www.OED.com. Oxford University Press. 2011.

^ "NIST Guide to SI Units". Retrieved 20071217.
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