Explain the observed trend in the boiling points of these compounds.
H₂Te -2 °C
H₂Se -41.5 °C
H₂S -60.7 °C
H₂O 100 °C

Verified step by step guidance

1. Identify the compounds and their respective boiling points: H₂Te (-2 °C), H₂Se (-41.5 °C), H₂S (-60.7 °C), and H₂O (100 °C).
2. Recognize that these compounds are hydrides of Group 16 elements (chalcogens) in the periodic table.
3. Understand that boiling points are influenced by intermolecular forces, such as hydrogen bonding, dipole-dipole interactions, and London dispersion forces.
4. Note that H₂O has a significantly higher boiling point due to strong hydrogen bonding, which is absent or weaker in the other hydrides.
5. Observe that the trend in boiling points for H₂Te, H₂Se, and H₂S is influenced by increasing molecular weight and size, leading to stronger London dispersion forces.

Key Concepts

Intermolecular Forces

Intermolecular forces are the attractive forces between molecules that influence physical properties such as boiling points. The main types include hydrogen bonding, dipole-dipole interactions, and London dispersion forces. Stronger intermolecular forces typically result in higher boiling points, as more energy is required to separate the molecules.

Hydrogen Bonding

Hydrogen bonding is a specific type of strong dipole-dipole interaction that occurs when hydrogen is bonded to highly electronegative atoms like oxygen, nitrogen, or fluorine. In the case of water (H2O), the presence of hydrogen bonds significantly raises its boiling point compared to other similar compounds, such as H2S and H2Se, which lack this strong interaction.

Molecular Weight and Structure

Molecular weight and structure also play a crucial role in determining boiling points. Generally, larger molecules with greater molecular weights exhibit higher boiling points due to increased London dispersion forces. However, the presence of polar bonds and the ability to form hydrogen bonds can override this trend, as seen in the comparison of H2O with the other compounds listed.