3. At 298 K and 1 atm, Br2 is a liquid with a high vapor pressure, and Cl2 is a gas. Those observati — Bonding Chemistry Question

STEPS:

1. Distinguish between intermolecular and intramolecular forces:
* Intermolecular forces (IMFs) are the attractive forces that exist *between* individual molecules (such as London dispersion forces). These forces govern physical macro-properties like phase states (solid, liquid, gas), boiling points, and vapor pressures.
* Intramolecular forces are the covalent chemical bonds *within* a single molecule (like the $\text{Cl}-\text{Cl}$ or $\text{Br}-\text{Br}$ single covalent bonds) that hold the individual atoms together.
2. Analyze the physical states of $\text{Br}_2$ and $\text{Cl}_2$ at $298\text{ K}$ and $1\text{ atm}$:
* Bromine ($\text{Br}_2$) is a liquid, meaning its molecules are held closely together in a condensed phase.
* Chlorine ($\text{Cl}_2$) is a gas, meaning its molecules have completely overcome their mutual attractions and are free to move far apart from one another.
3. Compare the strength of their intermolecular forces:
* Because it requires more thermal energy (higher temperature) to disrupt the attractions holding bromine molecules together in the liquid phase compared to chlorine (which is already a gas), the intermolecular forces among $\text{Br}_2$ molecules must be stronger than those among $\text{Cl}_2$ molecules.
* *(Note: Both are nonpolar diatomic halogens experiencing London dispersion forces. Because $\text{Br}_2$ has a larger, more polarizable electron cloud than $\text{Cl}_2$, its dispersion forces are significantly stronger).*
4. Relate the high vapor pressure of bromine to its liquid state:
* Although liquid bromine has a relatively high vapor pressure (meaning it evaporates easily because its IMFs are moderate), the fact that it is still a liquid at room temperature provides definitive evidence that its IMFs are stronger than those of gaseous chlorine under the same conditions.
5. Conclude:
* These phase observations directly compare the attractions *between* molecules. Therefore, they prove that the forces among $\text{Br}_2$ molecules are stronger than those among $\text{Cl}_2$ molecules, confirming Option A is correct.

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WHY_OTHERS_WRONG:

• Option B is incorrect: Intramolecular covalent bonds (which involve the sharing of valence electrons between atoms) are chemistry's primary chemical bonds and are orders of magnitude stronger than intermolecular attractions between neutral molecules. The forces among $\text{Cl}_2$ molecules (weak dispersion forces) are much weaker than the covalent $\text{Cl}-\text{Cl}$ bond.
• Options C and D are incorrect: These options compare the strength of the intramolecular covalent bonds ($\text{Br}-\text{Br}$ vs. $\text{Cl}-\text{Cl}$). Because a phase change (liquid to gas) and vapor pressure only involve overcoming attractions *between* separate molecules without breaking any chemical bonds, macroscopic phase observations do not provide evidence regarding the strength of the internal covalent bonds within the diatomic molecules.