How this instrument works
Molality is a measure of concentration defined as moles of solute per kilogram of solvent: m = moles of solute / kg of solvent, with units of mol/kg (sometimes written 'molal' or 'm'). Notice the denominator is the solvent alone, not the total solution — this is the detail that most often gets confused with molarity, which divides by the volume of the whole solution instead.
Because molality is built entirely from masses rather than a volume, it doesn't change as temperature changes. A litre of liquid expands slightly when heated and contracts when cooled, which shifts molarity (moles per litre) even though nothing was added or removed — but a kilogram of solvent stays a kilogram regardless of temperature or pressure, so molality stays fixed. That temperature-independence is exactly why molality is the standard unit for colligative property calculations like freezing-point depression and boiling-point elevation, which are derived assuming concentration doesn't drift with the very temperature change being calculated.
Molality and molarity are numerically close for dilute aqueous solutions, since a kilogram of water occupies almost exactly one litre near room temperature, but they diverge as concentration increases or as the solvent's density departs from water's. Chemists switch to molality whenever a calculation spans a temperature range, or whenever the solvent isn't water and its density isn't close to 1 kg/L, because using molarity in either situation would silently introduce a temperature- or density-dependent error.
- Enter the amount of dissolved substance, in moles, into Moles of solute (mol).
- Enter the mass of the solvent alone — not the whole solution — into Mass of solvent (kg).
- Read Molality (mol/kg) beneath the inputs; it recalculates as either value changes.
- If you only know the solvent's mass in grams, divide by 1000 before entering it, since this field expects kilograms.
Worked example — 0.5 mol dissolved in 0.25 kg of solvent
Enter 0.5 into Moles of solute (mol) and 0.25 into Mass of solvent (kg) — half a mole of solute dissolved into a quarter kilogram (250 g) of solvent. Molality reads 2.0 mol/kg: 0.5 / 0.25 = 2.0.
That 2.0 mol/kg figure stays fixed no matter what temperature the solution sits at, because it's anchored to the solvent's mass rather than the solution's volume. If this were a molarity calculation instead using the solution's volume, warming the liquid enough to expand it by even a percent or two would shift the reported concentration — molality is immune to that effect by construction.
Questions
What's the difference between molality and molarity?
Molality divides moles of solute by kilograms of solvent alone; molarity divides moles of solute by litres of the whole solution (solute plus solvent). Because molality's denominator is a mass, it never shifts with temperature; molarity's denominator is a volume, which expands or contracts as temperature changes, so molarity technically drifts slightly with temperature while molality does not.
Why use molality for freezing-point depression and boiling-point elevation?
Because those colligative-property formulas (ΔT = i x Kf x m or i x Kb x m) are derived on the assumption that concentration doesn't change as temperature changes during the measurement — which is exactly what molality guarantees and molarity does not. Using molarity in a freezing-point or boiling-point calculation would introduce an error that grows as the solution's temperature moves further from the volume's calibration point.
Are molality and molarity ever numerically equal?
They're close, though rarely exactly equal, for dilute solutions in water near room temperature, since a kilogram of water occupies almost exactly one litre (density approximately 1 kg/L). The two values diverge more noticeably as the solution becomes more concentrated, as temperature moves away from that reference point, or whenever the solvent isn't water and its density differs meaningfully from 1 kg/L.
Why does the mass of solvent go in kilograms rather than grams?
Because the formal SI-derived definition of molality is moles per kilogram, giving convenient mol/kg values for typical lab concentrations — moles per gram would produce inconveniently tiny decimal numbers for the same physical solution. If your solvent mass is recorded in grams, divide by 1000 before entering it here.
Does molality use the mass of the whole solution or just the solvent?
Just the solvent — this is the single detail that distinguishes molality from mass percent and from molarity's solution-volume denominator. If you have 0.5 mol of solute dissolved into 250 g of pure solvent, the denominator is 0.25 kg (the solvent alone), not the combined mass of solute plus solvent.