Question

Use a reference source such as the CRC Handbook of Chemistry and Physics to compare the melting and boiling points of the following pairs of inorganic substances: (a) \(\mathrm{W}\) and \(\mathrm{WF}_{6}\), (b) \(\mathrm{SO}_{2}\) and \(\mathrm{SF}_{4}\), (c) \(\mathrm{SiO}_{2}\) and \(\mathrm{SiCl}_{4}\). Account for the major differences observed in terms of likely structures and bonding.

Short answer

The differences in melting and boiling points of the given pairs of substances can be attributed to their structures and bonding types. In the case of \(\mathrm{W}\) and \(\mathrm{WF}_{6}\), the metallic bond in tungsten leads to much higher melting and boiling points compared to the covalent bonding in \(\mathrm{WF}_{6}\). For \(\mathrm{SO}_{2}\) and \(\mathrm{SF}_{4}\), the stronger intermolecular forces in \(\mathrm{SO}_{2}\) due to higher polarity result in its higher boiling point. Lastly, the higher melting and boiling points of \(\mathrm{SiO}_{2}\) compared to those of \(\mathrm{SiCl}_{4}\) are due to its network covalent solid structure with strong Si-O bonds.

Step-by-step solution

Look up the melting and boiling points of each substance

Using a reference source like the CRC Handbook of Chemistry and Physics, we can find the melting and boiling points for each substance. The values are listed below (note that these are approximate values): (a) \(\mathrm{W}\) (tungsten): melting point = 3422°C, boiling point = 5555°C \(\mathrm{WF}_{6}\) (tungsten hexafluoride): melting point = 2.3°C, boiling point = 17.1°C (b) \(\mathrm{SO}_{2}\) (sulfur dioxide): melting point = -72.4°C, boiling point = -10°C \(\mathrm{SF}_{4}\) (sulfur tetrafluoride): melting point = -121.8°C, boiling point = -38°C (c) \(\mathrm{SiO}_{2}\) (silicon dioxide): melting point = 1610°C, boiling point = 2230°C \(\mathrm{SiCl}_{4}\) (silicon tetrachloride): melting point = -70.2°C, boiling point = 57.65°C

Analyzing structures and types of bonding in the given substances

For each pair of substances, we can analyze the type of bonding present and the likely structures, which will help us understand the differences in their melting and boiling points. (a) \(\mathrm{W}\) (tungsten): metallic bonding \(\mathrm{WF}_{6}\) (tungsten hexafluoride): covalent bonding, strong W-F bonds (b) \(\mathrm{SO}_{2}\) (sulfur dioxide): covalent bonding, bent molecule structure \(\mathrm{SF}_{4}\) (sulfur tetrafluoride): covalent bonding, see-saw molecule structure (c) \(\mathrm{SiO}_{2}\) (silicon dioxide): covalent bonding, network covalent solid, strong Si-O bonds \(\mathrm{SiCl}_{4}\) (silicon tetrachloride): covalent bonding, tetrahedral molecule structure

Accounting for the differences in melting and boiling points

(a) The difference in melting and boiling points between \(\mathrm{W}\) and \(\mathrm{WF}_{6}\) can be attributed to the type of bonding present in each substance. In tungsten, metallic bonding allows the metal atoms to be held together by a 'sea' of electrons which leads to a strong bond and high melting and boiling points. In contrast, \(\mathrm{WF}_{6}\) has covalent bonds with a molecular structure instead of a metallic lattice, which results in weaker interactions between molecules. This leads to lower melting and boiling points. (b) Both \(\mathrm{SO}_{2}\) and \(\mathrm{SF}_{4}\) have covalent bonding, but their molecular structures cause slight differences in their melting and boiling points. \(\mathrm{SO}_{2}\) has a bent molecule structure with polar bonds and a net dipole moment, while \(\mathrm{SF}_{4}\) has a see-saw molecule structure with polar bonds but a slightly lower polarity than \(\mathrm{SO}_{2}\). The higher polarity in \(\mathrm{SO}_{2}\) results in stronger intermolecular forces, leading to its higher boiling point compared to that of \(\mathrm{SF}_{4}\). (c) The difference in melting and boiling points between \(\mathrm{SiO}_{2}\) and \(\mathrm{SiCl}_{4}\) can be attributed to their structures and type of bonding. \(\mathrm{SiO}_{2}\) is a network covalent solid with strong covalent bonds between Si and O atoms extending throughout the solid. This leads to high melting and boiling points. In contrast, \(\mathrm{SiCl}_{4}\) has covalent bonds leading to a tetrahedral molecular structure, which results in weaker intermolecular forces as well as lower melting and boiling points compared to \(\mathrm{SiO}_{2}\).

Key concepts

Phase Transition Temperatures

Understanding phase transition temperatures, namely melting points and boiling points, is crucial for students studying inorganic substances. These temperatures represent the specific conditions at which a substance changes from solid to liquid (melting) or from liquid to gas (boiling). For inorganic substances, differences in these temperatures often relate to strength of bonds and structural arrangement of atoms or ions within a material.

For instance, the melting point of tungsten (W) is significantly higher than that of tungsten hexafluoride (WF₆) primarily due to the different bonding types and strength within each substance. Understanding why substances have varying phase transition temperatures requires examining their bonding and structure, which is covered in more depth in subsequent sections.

It is essential to identify the factors affecting these temperatures, such as types of interparticle forces and the presence of molecular structure and polarity. In your studies, consider not only the values relating to phase transition temperatures, but also their correlation with intermolecular forces, the presence of a crystal lattice (in solids), and the packing of atoms, which can help explain why one substance melts or boils at a significantly different temperature than another.

Chemical Structure and Bonding

The chemical structure and bonding of a substance directly influence its melting and boiling points. In the case of W and WF₆, the former exhibits metallic bonding where a 'sea' of delocalized electrons contributes to a strong cohesive force between metal atoms, resulting in high phase transition temperatures. WF₆, whereas, demonstrates covalent bonding with distinct molecules held by comparatively weaker forces, which account for its lower melting and boiling points.

Similar investigations into molecular geometries, such as the bent structure of SO₂ and the see-saw shape of SF₄, reveal how molecular shape and polarity impact the strength of intermolecular forces. This explains their variance in boiling points despite both having covalent bonding. A student should note that substances with higher polarity will typically have stronger intermolecular attractions, hence higher boiling points. The network covalent structure of SiO₂ demonstrates how extended covalent bonding through a solid can dramatically raise transition temperatures compared to SiCl₄'s discrete tetrahedral molecules.

Inorganic Compound Properties

Inorganic compound properties encompass a broad spectrum, but when it comes to explaining melting and boiling points, one must consider the types of interactions that hold the compounds together. Each inorganic compound carries its unique physical properties determined by its chemical makeup and structure. For instance, SiO₂ stands out due to the robust covalent bonds forming a three-dimensional network, this translates to its high melting and boiling points. In contrast, SiCl₄ consists of individual SiCl₄ molecules relying on weaker van der Waals forces, hence its lower phase transition temperatures.

By exploring the properties of these inorganic compounds, students can associate the macroscale observable properties, like phase transition temperatures, to the microscale interactions and structures. This connection between the chemical bonding and properties serves as a foundation for understanding a wide range of material behaviors and is fundamental in the field of chemistry.