The boiling point of a substance is the temperature at which the vapor pressure of the liquid equals the pressure surrounding the liquid and the liquid changes into a vapor.
A liquid in a vacuum has a lower boiling point than when that liquid is at atmospheric pressure. A liquid at high-pressure has a higher boiling point than when that liquid is at atmospheric pressure. In other words, the boiling point of a liquid varies depending upon the surrounding environmental pressure. For a given pressure, different liquids boil at different temperatures.
A phase transition is the transformation of thermodynamic system from one phase or state of matter to another.
A phase of a thermodynamic system and the states of matter have uniform physical properties.
Dihydrogen monoxide (DHMO)
Hydrogen hydroxide (HH or HOH)
The melting point (or, rarely, liquefaction point) of a solid is the temperature at which it changes state from solid to liquid at atmospheric pressure. At the melting point the solid and liquid phase exist in equilibrium. The melting point of a substance depends (usually slightly) on pressure and is usually specified at standard pressure. When considered as the temperature of the reverse change from liquid to solid, it is referred to as the freezing point or crystallization point. Because of the ability of some substances to supercool, the freezing point is not considered as a characteristic property of a substance. When the "characteristic freezing point" of a substance is determined, in fact the actual methodology is almost always "the principle of observing the disappearance rather than the formation of ice", that is, the melting point.
Boiling-point elevation describes the phenomenon that the boiling point of a liquid (a solvent) will be higher when another compound is added, meaning that a solution has a higher boiling point than a pure solvent. This happens whenever a non-volatile solute, such as a salt, is added to a pure solvent, such as water. The boiling point can be measured accurately using an ebullioscope. In sum, therefore, the boiling-point elevation means that when a solute is dissolved in a solvent, the boiling point of the resulting solution is higher than that of the pure solvent.
The boiling point elevation is a colligative property, which means that it is dependent on the presence of dissolved particles and their number, but not their identity. It is an effect of the dilution of the solvent in the presence of a solute. It is a phenomenon that happens for all solutes in all solutions, even in ideal solutions, and does not depend on any specific solute-solvent interactions. The boiling point elevation happens both when the solute is an electrolyte, such as various salts, and a nonelectrolyte. In thermodynamic terms, the origin of the boiling point elevation is entropic and can be explained in terms of the vapor pressure or chemical potential of the solvent. In both cases, the explanation depends on the fact that many solutes are only present in the liquid phase and do not enter into the gas phase (except at extremely high temperatures).
Freezing food preserves it from the time it is prepared to the time it is eaten. Since early times, farmers, fishermen, and trappers have preserved their game and produce in unheated buildings during the winter season. Freezing food slows down decomposition by turning residual moisture into ice, inhibiting the growth of most bacterial species. In the food commodity industry, the process is called IQF or Individually Quick Frozen (flash freezing).
Preserving food in domestic kitchens during the 20th and 21st centuries is achieved using household freezers. Accepted advice to householders was to freeze food on the day of purchase. An initiative by a supermarket group in 2012 (backed by the UK's Waste & Resources Action Programme) promotes advising the freezing of food "as soon as possible up to the product's 'use by' date". The Food Standards Agency was reported as supporting the change, providing the food had been stored correctly up to that time.
Freezing-point depression describes the process in which adding a solute to a solvent decreases the freezing point of the solvent.
Examples include salt in water, alcohol in water, or the mixing of two solids such as impurities in a finely powdered drug. In such cases, the added compound is the solute, and the original solid can be thought of as the solvent. The resulting solution or solid-solid mixture has a lower freezing point than the pure solvent or solid did. This phenomenon is what causes sea water, (a mixture of salt (and other things) in water) to remain liquid at temperatures below 0 °C (32 °F), the freezing point of pure water.