Question

For each pair of solids, determine which solid has the higher melting point and explain why. (a) \(\mathrm{Ti}(\mathrm{s})\) and \(\mathrm{Ne}(\mathrm{s})\) (b) \(\mathrm{H}_{2} \mathrm{O}\) (s) and \(\mathrm{H}_{2} \mathrm{~S}(s)\) (c) \(\mathrm{Kr}(s)\) and \(\mathrm{Xe}(s)\) (d) \(\mathrm{NaCl}(s)\) and \(\mathrm{CH}_{4}(s)\)

Short answer

Titanium, water, Xenon, and Sodium chloride each have higher melting points compared to Neon, hydrogen sulfide, Krypton, and methane respectively, due to the nature and strength of their intermolecular forces.

Step-by-step solution

Comparing Melting Points of Ti(s) and Ne(s)

Titanium (Ti) is a metal with metallic bonding, while Neon (Ne) is a noble gas with weak van der Waals forces when solidified. Metallic bonds are much stronger than van der Waals forces, so Ti has a higher melting point than Ne.

Comparing Melting Points of H2O(s) and H2S(s)

Water (H2O) has hydrogen bonding, which is a strong type of dipole-dipole interaction, while hydrogen sulfide (H2S) has weaker dipole-dipole interactions and van der Waals forces. Hydrogen bonds are stronger than the other types of intermolecular forces present in H2S, thus H2O has a higher melting point.

Comparing Melting Points of Kr(s) and Xe(s)

Both Krypton (Kr) and Xenon (Xe) are noble gases and thus held together by van der Waals forces when solid. The strength of these forces increases with the size of the atoms, which correlates with their molar mass. Xe has a higher molar mass than Kr, resulting in stronger van der Waals forces, and consequently, a higher melting point.

Comparing Melting Points of NaCl(s) and CH4(s)

Sodium chloride (NaCl) is an ionic compound with strong ionic bonds, while methane (CH4) is a molecular compound with weak van der Waals forces. The ionic bonds in NaCl require more energy to break them, hence NaCl has a higher melting point than CH4.

Key concepts

Intermolecular Forces

Understanding intermolecular forces is critical when comparing the melting points of different substances. Intermolecular forces are the forces of attraction between molecules. These forces include hydrogen bonding, dipole-dipole interactions, and van der Waals forces, which comprise London dispersion forces and dipole-induced dipole interactions.

One way to visualize intermolecular forces is by imagining a crowd at a concert. Some people (molecules) have strong connections (bonds) to each other and stick closely together, while others stand apart with less interaction. The strength of these connections determines how much energy is needed to 'break' them apart, or in chemical terms, raise the temperature to a point where the substance transitions from solid to liquid; this is the melting point.

So, when comparing two substances, we look at the type and strength of their intermolecular forces. Generally, the substance with stronger intermolecular forces will have a higher melting point because more energy is required to overcome these forces.

Metallic Bonding

Metallic bonding is a type of chemical bonding that occurs between atoms of metallic elements. It's characterized by a sea of electrons that are shared and delocalized across a lattice of metal cations. These electrons, not held by any specific atom, can move freely throughout the metal. This unique structure presents both the high electrical conductivity and the substantial strength seen in metals.

Imagine a metallic bond as a strong, communal society—a group where resources (electrons) are shared, and everyone (metals ions) contributes to the group's cohesion and stability. This communal strength is why metals have high melting points. Each metal ion is bonded to several neighbors, so a considerable amount of energy is needed to break these numerous bonds and cause the metal to melt.

Hydrogen Bonding

Hydrogen bonding is a special type of dipole-dipole interaction that occurs when a hydrogen atom, which is bonded to a highly electronegative atom like oxygen, nitrogen, or fluorine, interacts with lone electron pairs on another electronegative atom. This bond is stronger than a typical van der Waals interaction yet weaker than ionic or covalent bonds.

In a metaphorical sense, think of hydrogen bonds like the hands of LEGO blocks snapping together. They're strong enough to hold shapes effectively but can be pulled apart with the right amount of force. Substances like water possess such bonds, making their intermolecular forces stronger than those in other molecular compounds without hydrogen bonding, thus yielding a higher melting point.

Ionic Bonds

Ionic bonds are electrostatic attractions between positively charged ions (cations) and negatively charged ions (anions). These bonds are formed when electrons are transferred from one atom to another, leading to a substance composed of ions. Ionic bonding is typically very strong because of the significant electrostatic forces between the oppositely charged ions.

You might liken ionic bonds to the relationship between opposite poles of magnets that firmly stick together—hard to separate unless a considerable force is applied. Therefore, when a substance like sodium chloride forms ionic bonds, it takes a lot of energy (in the form of heat) to disrupt these interactions and melt the substance, thus they typically have high melting points.

Van der Waals Forces

The term 'van der Waals forces' is an umbrella for various weak forces including dipole-dipole interactions, dipole-induced dipole interactions, and London dispersion forces. These are the kinds of forces that occur between non-polar molecules or between noble gases. The weakest of these, the London dispersion forces, are present in all molecules, regardless of their polarity, due to momentary changes in electron density that create temporary dipoles.

Consider van der Waals forces as acquaintances who only occasionally shake hands rather than the tight embrace of close friends. Because these forces are relatively weak, substances bound by van der Waals forces usually have low melting points. For instance, noble gases in their solid form only require a small amount of heat to overcome these weak intermolecular forces and transition to the liquid state.