Animated Solution for Chemistry - States of Matter: 'X' melts at low temperature and is a bad conductor of electricity in both liquid and solid state. X is

Imagine you are a chemical detective, and you've just been handed a vial containing a mysterious substance labeled simply as 'X'. Your mission is to identify it using only two macroscopic clues: it melts at a relatively low temperature, and it stubbornly refuses to conduct electricity, whether it is frozen solid or melted into a liquid puddle.

These macroscopic properties are not random; they are direct whispers from the microscopic world of atoms and bonds. To crack this case, we need to translate these physical behaviors into chemical structure.

In the realm of solid-state chemistry, crystalline solids are broadly classified into four distinct families based on the nature of the forces holding them together: Molecular, Ionic, Metallic, and Covalent (Network) solids.

Each family has a unique fingerprint. Metallic solids, like a copper wire, are bathed in a sea of delocalized electrons, making them excellent conductors. Ionic solids, like table salt, are rigid lattices of charged ions; they are insulators when solid but become conductive when melted, as the ions are finally free to move. Covalent network solids, like diamond, are massive, interconnected webs of strong covalent bonds, resulting in incredibly high melting points. Finally, molecular solids consist of discrete molecules held together by weak intermolecular forces, leading to low melting points and a lack of electrical conductivity.

Let's apply our clues to the given options. Our first clue, a low melting point, immediately tells us that the forces holding the solid together are weak. This strongly points towards a molecular solid. Our second clue, being a bad conductor in both solid and liquid states, confirms that there are absolutely no mobile charge carriers—no free electrons and no free ions.

Now, let's interrogate the suspects: - Silicon carbide ($SiC$): This is a covalent network solid. Its strong, three-dimensional network of covalent bonds gives it a massive melting point. It doesn't fit the profile. - Mercury ($Hg$): This is a metallic solid. While it does have a low melting point (it's a liquid at room temperature!), its sea of free electrons makes it an excellent conductor of electricity. Suspect dismissed. - Zinc sulphide ($ZnS$): This is an ionic solid. It has a high melting point due to strong electrostatic forces between the $Z{n}^{2+}$ and $S^{2-}$ ions. Furthermore, if we were to melt it, those ions would become mobile, and it would conduct electricity. It fails on both counts.

This leaves us with our final suspect: Carbon tetrachloride ($CC{l}_4$).

Carbon tetrachloride is a non-polar molecule. In its solid state, the individual $CC{l}_4$ molecules are held together only by weak London dispersion forces. Because these forces are so weak, it takes very little thermal energy to break them apart, resulting in a low melting point.

Furthermore, all the electrons in $CC{l}_4$ are tightly bound in covalent bonds between the carbon and chlorine atoms. There are no free electrons, and the molecule itself is neutral, meaning there are no ions. Therefore, it cannot conduct electricity in either the solid or liquid state.

Carbon tetrachloride perfectly matches every clue. The mystery is solved!