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In my textbook, it is written: The addition of impurities in a solid decreases the melting point of the solid.

How does it do so? Why can't it increase the melting point?

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  • $\begingroup$ A solid is a lattice held together by intermolecular forces rigidly (solids are rigid relative to a liquid). As you increase temperature, you add thermal energy to the lattice. At a certain temperature, the thermal energy overpowers the intermolecular forces (lattice energy), and the lattice breaks down. This is the melting point. The lattice energy stems from ordered structures. Impurities interrupt the crystalline structure, thereby weakening the intermolecular forces, which lowers the lattice energy and therefore the amount of thermal energy needed to melt, ie. the melting point decreases. $\endgroup$
    – Blaise
    Oct 13 '18 at 11:55
  • $\begingroup$ A liquid btw isnt a rigidly ordered structure. This flexibility allows it to incorporate impurities more efficiently into a network of maximum intermolecular forces, and so impurities can increase the boiling point. The inability to efficiently bring inpurities within an existing ordered crystal is why the mp wont increase. $\endgroup$
    – Blaise
    Oct 13 '18 at 12:04
  • $\begingroup$ Agreed. At first, I thought there's no way impurities lower mp. Like, when I add an impurity in any solvent, I'm kind of not allowing it to provide more 'space' for it to melt effectively because of it's addition. But, The kind of 'interuption' that impurities add lowers it's lattice energy and so the mp decreases. Ain't it? $\endgroup$
    – Yena shah
    Oct 13 '18 at 12:11
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one of the most prevalent (and simplest) model which explains this fact is colligative freezing point depression.

Freezing (or melting) point is temperature at which solid and liquid phase are in equilibrium, and their vapor pressures are equal. Adding (non-volatile) impurity in solid phase lowers the vapor pressure of solid phase (because you lower the "concentration" of your compound in solid phase), and phase equilibrium can be achieved at lower temperature.

Difference between melting points for pure and impure compounds can be quantified with following equation: $$ \Delta T = i K_f b $$ where $\Delta T$ is melting point difference, $i$ is Van't Hoff factor which tells you on how many species does the impurity dissociate, $K_f$ is cryoscopic constant which tells you how great the effect is, and $b$ is molality of the impurity.

Basically, all impurities lower the melting point, if they do not interact strongly with the given compound. In some cases (i.e. urea/choline chloride) freezing point depression can be in order of 100 °C! Organic compounds in particular have pretty high cryoscopic constants, which can cause headaches in organic synthesis (in most cases, you end up with some kind of oil which has to be purified to obtain solid product).

Elevation of freezing point is possible only if pressure is decreased, but the effect is so minuscule that it is not practically useful.

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