The melting point is the temperature at which the disruptive vibrations of the particles of the solid overcome the attractive forces operating within the solid. At some point, the amplitude of vibration becomes so large that the atoms start to invade the space of their nearest neighbors and disturb them, and the melting process initiates. As a solid is heated, its particles vibrate more rapidly as the solid absorbs kinetic energy. The motion of individual atoms, ions, or molecules in a solid is restricted to vibrational motion about a fixed point. The atoms in a solid are tightly bound to each other, either in a regular geometric lattice (crystalline solids, which include metals and ordinary ice) or irregularly (an amorphous solid such as common window glass), and are typically low in energy. Solids are similar to liquids in that both are condensed states, with particles that are far closer together than those of a gas. The first theory explaining the mechanism of melting in bulk was proposed by Lindemann, who used the vibration of atoms in the crystal to explain the melting transition. When considered as the temperature of the reverse change from liquid to solid, it is called the freezing point or crystallization point. The melting point of a substance depends on pressure and is usually specified at standard pressure. Adding heat will convert the solid into a liquid with no temperature change. In thermodynamics, the melting point defines a condition where the solid and liquid can exist in equilibrium. When considered as the temperature of the reverse change from vapor to liquid, it is called the condensation point. The pressure at which vaporization (boiling) starts to occur for a given temperature is called the saturation pressure. The temperature at which vaporization (boiling) starts to occur for a given pressure is called the saturation temperature or boiling point. In thermodynamics, saturationdefines a condition in which a mixture of vapor and liquid can exist together at a given temperature and pressure. Whenever you make an observation in the laboratory that just doesn't "feel" right, trust your instinct, pause, ensure safety, and more thoughtfully assess the situation before proceeding.Note that these points are associated with the standard atmospheric pressure. Carry out the experiment in a fume hood behind appropriate shielding, wear appropriate PPE for the situation, and alert your lab-mates what you are planning to do. Whenever you're going to be sealing something with the potential for evaporation/expansion, a quick Henry's Law calculation should be carried out to ensure that the system won't reach a critical pressure and rupture the container. When using LN2 cooling, never use an experimental set-up that contains air or oxygen. If there is an oxidizable material already in the flask (e.g., solvent and/or other organic material), the liquid oxygen may react rapidly even at the LN2 cold bath temperature. An empty vessel exposed to an oxygen headspace, most typically simply open to air, will condense/liquify the oxygen. Use of LN2 always introduces an added layer of risk because it has a lower boiling point (-196 ☌) than oxygen (-183 ☌). Do not let ammonia gas go into the fume hood as it will oxidize anything in it (including the monitor for air flow). Make sure to have a trap to capture ammonia that did not condense. This is the case for many things, not just ammonia-we are too quick to reach for LN2 for many applications (freeze/pump/thaw degassing, for example). There's no need to use LN2 (-196 ☌) when a dry ice +acetone/IPA bath (-78 ☌) would suffice.
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