In chemistry, the molality, b (or m), of a solution is defined as the amount of substance of solute (in moles), nsolute, divided by the mass in kg of the solvent, msolvent:
Molality is an intensive property of solutions, such as sodium chloride (table salt) in water .
The earliest definition of molality was most likely coined by G. N. Lewis and M. Randall in their 1923 publication Thermodynamics and the Free Energies of Chemical Substances. The two words "molality" and "molarity" are apt to be confused with one another, and in fact the molality and molarity of a weak aqueous solution happen to be nearly the same, as one kilogram of water (the solvent) occupies one liter of volume at room temperature, and the small amount of solute would have little effect on the volume of the solvent.
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, suggesting instead mol/kg or another related SI unit. Despite this, their recommendation has not been universally implemented in scientific literature, and many references to molality still exist.
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. This is because chemical reactions occur in proportions of mass, not volume. Since volume is subject to variation due to temperature and pressure, weighing is in fact an advantage because mass does not vary with ambient conditions. The mass-based nature of molality implies that it can be readily converted into a mass ratio (or mass fraction, "w," ratio):
where the symbol M stands for molar mass. This works because the units of molality are mol solute/kg solvent, and if we multiply the moles of solute by its molar mass (in grams per mole), we obtain grams of solute. The equation thus ultimately yields grams solute per kg of solvent, a mass ratio.
Likewise, a mole ratio of the two substances (solvent and solute) can be obtained:
Again, this works because the denominator in molality is the mass of the solute, and multiplying 1/kg by kg/mol (easily calculated from grams/mol) gives us a unit of moles, so we now have mol solute/mol solvent. These ratios are much more difficult, if not impossible, to determine for solutions of known molarity. However, if the solvent is reactive, and one needs to know the stoichiometry between the solvent and the solute, knowing the molality can be very important and much more useful. More commonly, however, the solvent is unreactive (or is simply required in excess), and the mole or mass ratios of the solute and the solvent do not need to be known, making molarity a more useful tool.