melting point

(noun)

The temperature at which the solid and liquid phases of a substance are in equilibrium; it is relatively insensitive to changes in pressure.

Related Terms

  • boiling point

Examples of melting point in the following topics:

  • Crystalline Solids

    • Most organic compounds have melting points below 200 ºC.
    • If two crystalline compounds (A & B) are thoroughly mixed, the melting point of that mixture is normally depressed and broadened, relative to the characteristic sharp melting point of each pure component.
    • The lowest mixture melting point, e, is called the eutectic point.
    • The A:B complex has a melting point of 54 ºC, and the phase diagram displays two eutectic points, the first at 50 ºC, the second at 30 ºC.
    • The compound was first prepared in England in 1946, and had a melting point of 58 ºC.
  • Boiling & Melting Points

    • The melting points of crystalline solids cannot be categorized in as simple a fashion as boiling points.
    • Spherically shaped molecules generally have relatively high melting points, which in some cases approach the boiling point.
    • This structure or shape sensitivity is one of the reasons that melting points are widely used to identify specific compounds.
    • Notice that the boiling points of the unbranched alkanes (pentane through decane) increase rather smoothly with molecular weight, but the melting points of the even-carbon chains increase more than those of the odd-carbon chains.
    • The last compound, an isomer of octane, is nearly spherical and has an exceptionally high melting point (only 6º below the boiling point).
  • Variation of Physical Properties Within a Group

    • The physical properties (notably, melting and boiling points) of the elements in a given group vary as you move down the table.
    • For example, pure carbon can exist as diamond, which has a very high melting point, or as graphite, whose melting point is still high but much lower than that of diamond.
    • Different groups exhibit different trends in boiling and melting points.
    • For Groups 1 and 2, the boiling and melting points decrease as you move down the group.
    • The noble gases (Group 18) decrease in their boiling and melting points down the group.
  • Fats & Oils

    • Thus, the melting points of triglycerides reflect their composition, as shown by the following examples.
    • Natural mixed triglycerides have somewhat lower melting points, the melting point of lard being near 30 º C, whereas olive oil melts near -6 º C.
  • Physical Properties of Carboxylic Acids

    • The table at the beginning of this page gave the melting and boiling points for a homologous group of carboxylic acids having from one to ten carbon atoms.
    • The boiling points increased with size in a regular manner, but the melting points did not.
    • Unbranched acids made up of an even number of carbon atoms have melting points higher than the odd numbered homologs having one more or one less carbon.
    • In the table of fatty acids we see that the presence of a cis-double bond significantly lowers the melting point of a compound.
    • Thus, palmitoleic acid melts over 60º lower than palmitic acid, and similar decreases occur for the C18 and C20 compounds.
  • Variation of Physical Properties Across a Period

    • Another physical property that varies across a period is the melting point of the corresponding halide.
    • The melting point is correlated to the strength of intermolecular bonds within the element.
    • All of the alkali halides and alkaline earth halides are solids at room temperature and have melting points in the hundreds of degrees centigrade.
    • For example, the melting point of sodium chloride (NaCl) is 808 °C.
    • In contrast, the melting points of the non-metal halides from Periods 2 and 3, such as CCl4, PCl3, and SCl2, are below 0 °C, so these materials are liquids at room temperature.
  • Molecular Crystals

    • Molecular solids tend to be soft or deformable, have low melting points, and are often sufficiently volatile to evaporate directly into the gas phase.
    • Whereas the characteristic melting point of metals and ionic solids is ~1000 °C, most molecular solids melt well below ~300 °C.
    • Because dispersion forces and the other van der Waals forces increase with the number of atoms, large molecules are generally less volatile, and have higher melting points than smaller ones.
    • A molecular solid, white phosphorus has a relatively low density of 1.82 g/cm3 and melting point of 44.1 °C; it is a soft material which can be cut with a knife.
    • How does changing the Van der Waals attraction or charging the atoms affect the melting and boiling point of the substance?
  • Liquid to Solid Phase Transition

    • For most substances, the melting and freezing points are the same temperature; however, certain substances possess different solid-liquid transition temperatures.
    • This is a first-order thermodynamic phase transition, which means that as long as solid and liquid coexist, the equilibrium temperature of the system remains constant and equal to the melting point.
    • Crystallization of pure liquids usually begins at a lower temperature than the melting point, due to the high activation energy of homogeneous nucleation.
    • The melting point of water at one atmosphere of pressure is very close to 0 °C (32 °F, 273.15 K), and in the presence of nucleating substances the freezing point of water is close to the melting point.
    • How quickly do the more energetic atoms melt the solid?
  • Phase Changes and Energy Conservation

    • To boil or melt one mole of a substance, a certain amount of energy is required.
    • If that amount of energy is added to a mole of that substance at boiling or freezing point, all of it will melt or boil, but the temperature won't change.
    • Temperature increases linearly with heat, until the melting point .
    • Once all ice has been melted, the temperature again rises linearly with heat added.
    • The curve ends at a point called the critical point, because at higher temperatures the liquid phase does not exist at any pressure.
  • Latent Heat

    • For example, consider water dripping from icicles melting on a roof warmed by the Sun.
    • The temperature of a glass of lemonade initially at 0 ºC stays at 0 ºC until all the ice has melted.
    • Even more energy is required to vaporize water; it would take 2256 kJ to change 1 kg of liquid water at the normal boiling point (100ºC at atmospheric pressure) to steam (water vapor).
    • Once at this temperature, the ice begins to melt until all the entire sample has melted, absorbing a total of 79.8 cal/g of heat.
    • Heat from the air transfers to the ice causing it to melt.
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