If the boiling point of H2O is 373 K, the boiling point of H2S will be
(a) less than 300 K
(b) equal

Q: Why doesn't Sulfur form hydrogen bonds like Oxygen?

Sulfur has a larger atomic radius and lower electronegativity compared to Oxygen. This means the $\text{S—H}$ bond is not polar enough to create the strong dipole required for hydrogen bonding.

Q: How many hydrogen bonds can a single water molecule form?

A single water molecule can form up to four hydrogen bonds: two using its partially positive hydrogen atoms, and two using the lone pairs on its partially negative oxygen atom.

Q: Why is $\text{H}_2\text{S}$ a gas at room temperature?

Because $\text{H}_2\text{S}$ molecules cannot form hydrogen bonds, they are only held together by weak dipole-dipole interactions and London dispersion forces. These weak forces are easily overcome by thermal energy at room temperature.

Q: Does $\text{H}_2\text{S}$ have any intermolecular forces at all?

Yes! While it lacks hydrogen bonding, it still experiences weak van der Waals forces, specifically dipole-dipole interactions (since it is a bent, polar molecule) and London dispersion forces.

Q: What happens to the boiling point of other Group 16 hydrides?

Water has an anomalously high boiling point due to hydrogen bonding. For the rest of the group ($\text{H}_2\text{S}$, $\text{H}_2\text{Se}$, $\text{H}_2\text{Te}$), the boiling point steadily increases as the molar mass and size of the electron cloud increase, which strengthens the London dispersion forces.

Question

If the boiling point of $	ext{H}_2	ext{O}$ is $373	ext{ K}$, the boiling point of $	ext{H}_2	ext{S}$ will be

Options:

(A) less than $300	ext{ K}$

(B) equal to $373	ext{ K}$

(C) more than $373	ext{ K}$

(D) greater than $300	ext{ K}$ but less than $373	ext{ K}$

Sigma Percentile · Accepted Answer

### The Setup
Imagine you are looking at two very similar molecules: water ($	ext{H}_2	ext{O}$) and hydrogen sulfide ($	ext{H}_2	ext{S}$). Both Oxygen and Sulfur belong to Group 16 of the periodic table, so you might expect their hydrides to behave similarly. However, their physical states at room temperature are drastically different. Water is a life-giving liquid, while hydrogen sulfide is a foul-smelling gas. Why does this happen?

### The Magic of Hydrogen Bonding
The secret lies in the atomic properties of Oxygen. Oxygen is highly electronegative and has a very small atomic radius. When it bonds with Hydrogen, it pulls the shared electrons strongly towards itself, creating a significant partial negative charge on the Oxygen and a partial positive charge on the Hydrogen.

Sulfur, on the other hand, is much larger and less electronegative. The $	ext{S—H}$ bond is not nearly as polar as the $	ext{O—H}$ bond. Because of this, Sulfur cannot participate in the special, strong intermolecular force known as **hydrogen bonding**.

### The Network of Water
Because of its strong dipole, a single water molecule can form up to four intermolecular hydrogen bonds with its neighbors. This creates a massive, sticky, three-dimensional network of molecules.

This extensive molecular association means that the molecules hold onto each other very tightly. To break these bonds and turn liquid water into a gas, you need to supply a lot of thermal energy. This is exactly why water has an unusually high boiling point of $373	ext{ K}$.

### The Loner Hydrogen Sulfide
Now, let's shift our focus to hydrogen sulfide. Without the ability to form hydrogen bonds, $	ext{H}_2	ext{S}$ molecules only experience weak van der Waals forces, specifically dipole-dipole interactions and London dispersion forces.

These weak forces are easily overcome by the ambient thermal energy present at room temperature. As a result, the molecules don't stick together well, making $	ext{H}_2	ext{S}$ a gas under normal conditions.

### Final Calculation
Room temperature is approximately $298	ext{ K}$. Since we know that $	ext{H}_2	ext{S}$ is a gas at this temperature, it must have already boiled!

Therefore, its boiling point must be significantly lower than room temperature. This leads us to the undeniable conclusion that the boiling point of $	ext{H}_2	ext{S}$ is **less than $300	ext{ K}$**.

Animated Solution for Chemistry - Chemical Bonding and Molecular Structure: If the boiling point of $H_{2} O$ is $373 K$ , the boiling point of $H_{2} S$ will be

Select Answer:

Visualized Solution

$H_{2} O vs H_{2} S$

$Atomic Properties$

$Hydrogen Bonding in Water$

$Boiling Point of Water$

$Hydrogen Sulfide$

$Intermolecular Forces in H_{2} S$

$Final Conclusion$

The Sigma Insight: Bond Parameters and Resonance

The Setup Imagine you are looking at two very similar molecules: water ( $H_{2} O$ ) and hydrogen sulfide ( $H_{2} S$ )

The Magic of Hydrogen Bonding The secret lies in the atomic properties of Oxygen

The Network of Water Because of its strong dipole, a single water molecule can form up to four intermolecular hydrogen bonds with its neighbors

The Loner Hydrogen Sulfide Now, let's shift our focus to hydrogen sulfide

Final Calculation Room temperature is approximately $298 K$

Similar Questions

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Select the correct statement.

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Animated Solution for Chemistry - Chemical Bonding and Molecular Structure: .math-text-wrapper .katex { font-size: 1.05em !important; } .math-text-wrapper .katex-display { margin: 1rem 0 !important; } If the boiling point of H2​O is 373 K, the boiling point of H2​S will be

Select Answer:

Visualized Solution

.math-text-wrapper .katex { font-size: 1.05em !important; } .math-text-wrapper .katex-display { margin: 1rem 0 !important; } H2​O vs H2​S

.math-text-wrapper .katex { font-size: 1.05em !important; } .math-text-wrapper .katex-display { margin: 1rem 0 !important; } Atomic Properties

.math-text-wrapper .katex { font-size: 1.05em !important; } .math-text-wrapper .katex-display { margin: 1rem 0 !important; } Hydrogen Bonding in Water

.math-text-wrapper .katex { font-size: 1.05em !important; } .math-text-wrapper .katex-display { margin: 1rem 0 !important; } Boiling Point of Water

.math-text-wrapper .katex { font-size: 1.05em !important; } .math-text-wrapper .katex-display { margin: 1rem 0 !important; } Hydrogen Sulfide

.math-text-wrapper .katex { font-size: 1.05em !important; } .math-text-wrapper .katex-display { margin: 1rem 0 !important; } Intermolecular Forces in H2​S

.math-text-wrapper .katex { font-size: 1.05em !important; } .math-text-wrapper .katex-display { margin: 1rem 0 !important; } Final Conclusion