Concept:The standard enthalpy of atomization (\( \Delta H_{at}^\circ \)) is the enthalpy change when one mole of gaseous atoms is formed from an element in its standard state.
Formula:$$ \text{Element}_{\text{(standard state)}} \rightarrow \text{Gaseous Atoms} \quad (\Delta H = +\text{ve}) $$
Solution:- To turn a solid (like Na) or a diatomic gas (like \( Cl_2 \)) into separated gaseous atoms, chemical bonds or strong intermolecular forces must be broken.
- Bond breaking strictly requires an input of energy, making atomization an endothermic process (\( \Delta H = +\text{ve} \)).
- In an endothermic process, the reacting system actively absorbs heat from the immediate surroundings, which causes the heat of the surroundings to decrease.
Why other options are incorrect:The heat of the surroundings would only increase if the reaction were exothermic (releasing heat). It does not remain the same because energy transfer is required to break the bonds.
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