How a correct statement can lead to misleading or practically irrelevant conclusions
Pure water boils at approximately 100 degrees Celsius (°C) at standard atmospheric pressure, roughly the pressure at sea level. Dissolving a non-volatile substance such as salt in the water raises its boiling point. This is one of the so-called colligative properties: in dilute solutions, these depend on the concentration of dissolved particles rather than their chemical identity, although they do depend on the solvent.
Adding ordinary table salt to water raises its boiling point by an amount that can be estimated quite easily (see Note 1). This well-known physical effect is sometimes invoked in cooking advice, particularly in Italy, a country that prides itself on being the home of pasta served with ragù, tomato sauce and countless other accompaniments. Some cooks recommend adding salt before the water boils, and certainly before adding the pasta, arguing that the pasta will then cook at a higher temperature. Salt does indeed raise the boiling point, but it does so whether it is added before or after the water starts boiling, once it has dissolved and the water is boiling again.
Let us do the sums. For a typical quantity of around 10 grams of salt per litre of water, the estimated increase is about 0.17 °C. It is hard to see how a rise from approximately 100 to 100.17 °C, less than two-tenths of a degree, could make any meaningful practical difference to the time needed to bring the water to the boil or to the quality of the finished pasta.
Altitude has a much greater effect on cooking. At 1,000 metres above sea level, water boils at about 97 °C; at 2,000 metres, at about 93 °C. Differences of this size have a far more noticeable effect on how pasta cooks. In an ordinary open pan, turning up the heat will not raise the temperature of the boiling water: we need to allow more cooking time, or use a pressure cooker (see Note 2).
A similar effect, but in the opposite direction, occurs with the freezing point of pure water, which is approximately 0 °C at standard atmospheric pressure. Adding salt lowers this temperature, an effect used to help melt snow and ice on roads in winter. Salt application rates can, for example, range from 10 to 40 grams per square metre, depending on the conditions. Without knowing how much water is present, however, even a rough estimate of the resulting salt concentration is difficult. For water, the freezing-point depression constant is 1.86 °C kg mol⁻¹, more than three times the boiling-point elevation constant. At the same dilute-solution particle concentration, the fall in freezing point is therefore more than three times the rise in boiling point. The practical effectiveness of road salt also depends on the concentration of the brine that forms and on the temperature; the larger constant alone does not explain it.
Notes
1) For a dilute solution of a non-volatile solute, the boiling-point elevation can be estimated using:
ΔTb ≈ Kb m i
Here, ΔTb is the rise in boiling point, Kb is the boiling-point elevation constant of the solvent (approximately 0.51 °C kg mol⁻¹ for water), m is the molality of the solute, in moles per kilogram of solvent, and i is the van ’t Hoff factor, which accounts for the number of dissolved particles produced per formula unit of solute.
For 10 grams of sodium chloride (NaCl) in one litre of water, approximated as one kilogram, the amount of NaCl is 10/58.44 ≈ 0.171 mol. Each formula unit yields two ions, Na⁺ and Cl⁻, so taking i = 2 gives approximately 0.342 mol of ions per kilogram of water. Hence:
ΔTb ≈ 2 × 0.51 × 0.171 ≈ 0.17 °C
This is an ideal dilute-solution estimate, rather than an exact result: interactions between ions cause deviations from ideal behaviour.
2) A pressure cooker produces a much larger effect in the opposite direction to altitude: the higher pressure can raise the boiling point of water to around 120 °C, depending on the operating pressure.
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