Problem 21 - Chemistry The Molecular Science

Question

Briefly explain the variations in the boiling points in this table. In your discussion be sure to mention the types of intermolecular forces involved. $$ \begin{array}{lc} \hline \text { Compound } & \text { Boiling Point }\left({ }^{\circ} \mathrm{C}\right) \\ \hline \mathrm{NH}_{3} & -33.4 \\ \mathrm{PH}_{3} & -87.5 \\ \mathrm{AsH}_{3} & -62.4 \\ \mathrm{SbH}_{3} & -17 \\ \hline \end{array} $$

Step-by-Step Solution

Answer

NH₃ has higher boiling point due to hydrogen bonding. Others rise in boiling point with molecular size due to stronger dispersion forces.

Step 1: Identify Intermolecular Forces

The compounds listed (NH₃, PH₃, AsH₃, SbH₃) all exhibit certain intermolecular forces that influence their boiling points. The dominant forces at play here are hydrogen bonding, dipole-dipole interactions, and London dispersion forces.

Step 2: Analyze Hydrogen Bonding

Among these compounds, NH₃ is capable of hydrogen bonding due to the presence of a highly electronegative nitrogen atom bonded to hydrogen. This results in a relatively strong intermolecular force compared to the others, leading to a higher boiling point of -33.4°C.

Step 3: Consider Dipole-Dipole Interactions

While PH₃, AsH₃, and SbH₃ lack significant hydrogen bonding, they do exhibit dipole-dipole interactions because of the polar bonds caused by the difference in electronegativity between hydrogen and the central atom (P, As, Sb). These interactions are weaker than hydrogen bonds, but stronger than just dispersion forces alone.

Step 4: Evaluate London Dispersion Forces

All these compounds have London dispersion forces, which increase with greater molar mass and surface area. As we move down the group from NH₃ to SbH₃, the molecules become larger and exhibit stronger dispersion forces, increasingly affecting boiling points.

Step 5: Assess Trends in Boiling Points

Notice that for PH₃, AsH₃, and SbH₃, the boiling point rises as the molar mass increases (from PH₃ to SbH₃), indicating that dispersion forces become more significant than dipole-dipole forces. However, NH₃'s boiling point remains distinctly higher due to hydrogen bonding despite its lower molar mass.

Key Concepts

Intermolecular Forces

In the world of chemistry, understanding boiling points calls for a grasp of intermolecular forces. These forces are the attractions that occur between molecules and significantly impact the boiling point of substances. Boiling happens when a liquid turns into a gas, requiring molecules to overcome these attractions.
There are three main types of intermolecular forces to consider:

• Hydrogen bonding: Strong, especially in molecules with H bonded to N, O, or F.
• Dipole-dipole interactions: Occur between molecules with polar bonds.
• London dispersion forces: These are the weakest, present in all molecules, but more prominent in larger ones.

The interplay of these forces explains why different substances boil at different temperatures.

Hydrogen Bonding

Hydrogen bonding is a special type of dipole-dipole attraction occurring in molecules where hydrogen is covalently bonded to a highly electronegative element, namely nitrogen, oxygen, or fluorine. This bond occurs because the high electronegativity of these atoms attracts the shared electrons strongly, creating a polar bond.
In ammonia ( ext{NH}_3), hydrogen bonding has a large role. The nitrogen atom is very electronegative, pulling electron density away from hydrogen atoms and creating a dipole. When ext{NH}_3 e molecules come close, the slight positive charge on the hydrogen of one molecule can be attracted to the lone pair of electrons on the nitrogen atom of another.

• This creates a strong link between molecules, raising the boiling point.
• For ext{NH}_3 e, this results in a boiling point of -33.4 °C, higher than expected considering its size.

Hydrogen bonding is a key reason why molecules like water and ammonia have significantly higher boiling points than similar-sized molecules without hydrogen bonds.

London Dispersion Forces

Even though they're the weakest, London dispersion forces are present in every molecule. They arise due to momentary shifts in electron cloud distribution, leading to temporary dipoles.
As molecules become larger with more electrons, these forces become stronger:\[ \text{Larger molecules have more electrons, increasing dispersion forces.} \]
In the compounds we are comparing, from ext{PH}_3 to ext{SbH}_3, the size of the molecules increases. This increase leads to stronger London dispersion forces:

• These forces become more influential in determining boiling points as molecular weight increases.
• In turn, this contributes to the rising boiling points from ext{PH}_3 (at -87.5 °C) to ext{SbH}_3 (at -17 °C).

Thus, London dispersion forces explain the increase in boiling points for larger molecules, where their impact outweighs dipole-dipole interactions as size grows.

Dipole-Dipole Interactions

Dipole-dipole interactions occur in molecules with a permanent dipole moment, due to their polar covalent bonds. Such bonds happen when atoms have different electronegativities, creating partial charges that attract oppositely charged regions of neighboring molecules.
In compounds like ext{PH}_3, ext{AsH}_3, and ext{SbH}_3, the central atom bonded to hydrogen creates a polar environment due to differences in electronegativity. While these dipole-dipole interactions are generally weaker than hydrogen bonds:

• In smaller molecules (e.g., ext{PH}_3), they do influence boiling points, though less so than hydrogen bonds.
• These interactions add to overall molecular attraction but are overshadowed in larger molecules by London dispersion forces.

Understanding dipole-dipole interactions helps clarify why polar molecules, though affected by other forces, tend to have varying boiling points based on polarity and size.