Question

Consider the following melting point data: $$ \begin{array}{|llllllll|} \hline \text { Compound: } & \mathrm{NaCl} & \mathrm{MgCl}_{2} & \mathrm{AlCl}_{3} & \mathrm{SiCl}_{4} & \mathrm{PCl}_{3} & \mathrm{SCl}_{2} & \mathrm{Cl}_{2} \\ \operatorname{mp}\left({ }^{\circ} \mathrm{C}\right): & 801 & 708 & 190 & -70 & -91 & -78 & -101 \\ \text { Compound: } & \mathrm{NaF} & \mathrm{MgF}_{2} & \mathrm{AlF}_{3} & \mathrm{SiF}_{4} & \mathrm{PF}_{5} & \mathrm{SF}_{6} & \mathrm{~F}_{2} \\ \operatorname{mp}\left({ }^{\circ} \mathrm{C}\right): & 997 & 1396 & 1040 & -90 & -94 & -56 & -220 \\ \hline \end{array} $$ Account for the trends in melting points in terms of interparticle forces.

Short answer

The observed trends in melting points of the given compounds can be accounted for by analyzing the interparticle forces present. For ionic compounds (NaCl, MgCl2, AlCl3, NaF, MgF2, AlF3), stronger electrostatic attractions due to higher charge density result in higher melting points. In covalent compounds (SiCl4, SiF4, PCl3, PF5, SCl2, SF6), the weaker London dispersion forces affect the melting points, but the differences are not significant enough to show a clear trend. Overall, the identified trends reflect the strength of the interparticle forces, be it ionic or covalent bonding.

Step-by-step solution

Identify the types of bonds and forces present in the compounds

Looking at the given compounds, we can identify that they are all binary compounds formed between a halogen (Cl or F) and elements from Groups 1, 2, 13, 14, 15, and 16. We need to consider the types of bonds and forces present in the compounds: - For Group 1 elements (Na): Ionic bonds with halogens (NaCl, NaF) - For Group 2 elements (Mg): Ionic bonds with halogens (MgCl2, MgF2) - For Group 13 elements (Al): Ionic bonds with halogens (AlCl3, AlF3) - For Group 14 elements (Si): Covalent bonds with halogens (SiCl4, SiF4) - For Group 15 elements (P): Covalent bonds with halogens (PCl3, PF5) - For Group 16 elements (S): Covalent bonds with halogens (SCl2, SF6) - For halogens (Cl2, F2): Covalent bonds

Analyze the melting point trends across the periods

The melting points of the ionic compounds (NaCl, MgCl2, AlCl3, NaF, MgF2, and AlF3) are generally higher than those of the covalent compounds (SiCl4, SiF4, PCl3, PF5, SCl2, SF6, and Cl2 isoclinal). This is because ionic compounds have strong electrostatic attractions between their positive and negative ions, which require more energy to overcome, resulting in higher melting points. Now, let's analyze the trends in melting points for each period: - Across Period 3 (NaCl, MgCl2, AlCl3, SiCl4, PCl3, SCl2): The melting points generally decrease along the period until SiCl4, after which they stay relatively constant. - Across Period 2 (NaF, MgF2, AlF3, SiF4, PF5, SF6): The melting points generally increase along the period to AlF3, and then decrease in the following compounds. The increase in melting points in Period 2 is due to the increase in charge density, which results in a larger electrostatic force between the ions in the ionic compounds, thus increasing their melting points. The decrease in melting points in covalent compounds in both periods is due to weaker London dispersion forces between the molecules.

Explain the trends in terms of interparticle forces

To account for the observed trends in melting points, we can explain the changes in interparticle forces as follows: - For ionic compounds, the strength of the electrostatic attraction between ions increases with higher charge density. This results in higher melting points for compounds with smaller ions and a higher charge (e.g., in Period 2: NaF, MgF2, and AlF3). - For covalent compounds, the strength of the London dispersion forces is directly related to the number of electrons in the molecule (as it affects its polarizability). However, the differences in dispersion forces are not strong enough to show a clear trend in melting points across the period. In conclusion, the melting points of the given compounds are mainly determined by the type of interparticle force present (ionic or covalent bonding) and the strength of these forces as influenced by factors such as charge density and polarizability.

Key concepts

Understanding Interparticle Forces

To make sense of melting point trends in chemistry, we must first delve into the concept of interparticle forces. These forces are what hold the particles within a substance together—be they atoms, ions, or molecules. In solids, these forces are strong enough to maintain the particles in close proximity, forming a defined shape and volume. The type of interparticle forces present will dictate how much energy is required to change a solid into a liquid, which we observe as the melting point.

Understanding these forces gives us insight into why different substances melt at varying temperatures. For example, ionic compounds tend to have higher melting points due to the strong electrostatic attractions between positively and negatively charged ions. In contrast, molecular compounds are usually held together by weaker forces, such as dipole-dipole interactions or London dispersion forces, resulting in lower melting points.

Charge Density and Melting Points

A key factor that affects the strength of interparticle forces in ionic compounds is charge density. A high charge density means that the ions have a larger charge in relation to their size, which intensifies the electrostatic attractions between them. Hence, with other factors being equal, compounds with higher charge densities melt at higher temperatures due to the greater forces holding the ions together.

Ionic and Covalent Bonds

The type of chemical bond—ionic or covalent—plays a pivotal role in a compound's melting point. Ionic bonds are formed between elements with a large difference in electronegativity, typically a metal and a nonmetal. As a result, one atom donates electrons to another, creating oppositely charged ions that attract each other strongly. This ionic bonding results in a lattice structure that requires substantial amounts of energy to break, consequently leading to high melting points, as seen in NaCl and MgF2.

Conversely, covalent bonds arise when two nonmetals share electrons, and the resulting molecules are held together by less intense forces. In simple molecular structures, such as SiCl4 or PCl3, these compounds exhibit lower melting points since the intermolecular forces between the molecules, like London dispersion or dipole-dipole interactions, require less energy to overcome when compared to the ionic counterparts.

Understanding Bond Strengths

The strength of the ionic or covalent bond contributes significantly to the observed melting points. Stronger ionic bonds, with higher ionic charges and smaller ionic radii, result in substances that are more difficult to melt. This trend can be seen when comparing the melting points of NaF to MgF2, where the higher charged Mg2+ ion contributes to MgF2's significantly higher melting point.

London Dispersion Forces

Among the various intermolecular forces, London dispersion forces are the weakest but are present in all molecular compounds. They originate from the temporary dipoles that occur when electron density within a molecule fluctuates, leading to a temporary uneven distribution of charge. Every so often, one side of a molecule becomes slightly more negative, inducing a temporary dipole in neighboring molecules.

These forces increase with the size of the molecules, or more specifically with the number of electrons. Larger atoms or molecules with more electrons have more pronounced shifts in electron density, resulting in stronger London dispersion forces. Even though these forces are weak individually, they can become substantial when many interactions are present.

Polarizability and its Effect

The concept of polarizability refers to how easily an electron cloud is distorted, which affects the strength of London dispersion forces. Larger molecules are more polarizable and, therefore, have stronger London dispersion forces, which can lead to higher melting points. For example, within a homologous series of hydrocarbons, as the molecular mass increases, so does the melting point, mainly due to the increasing strength of London dispersion forces.