MHT CET · Chemistry · Thermodynamics (C)
Calculate Gibbs energy change for a reaction having \(\Delta \mathrm{H}=31400 \mathrm{~J}, \Delta \mathrm{~S}=32 \mathrm{JK}^{-1}\) at \(1000^{\circ} \mathrm{C}\) ?
- A -4668 J
- B -9336 J
- C -4073 J
- D -2334 J
Answer & Solution
Correct Answer
(B) -9336 J
Step-by-step Solution
Detailed explanation
\(\begin{aligned}
& \mathrm{T}=1000^{\circ} \mathrm{C}=1273 \mathrm{~K} \\
& \Delta \dot{G}=\Delta \mathrm{H}-\mathrm{T} \Delta \mathrm{~S} \\
& =31400 \mathrm{~J}-1273 \mathrm{~K}\left(32 \mathrm{~J} \mathrm{~K}^{-1}\right) \\
& =3.1400 \mathrm{~J}-40736 \mathrm{~J} \\
& \Delta \mathrm{G}=-9336 \mathrm{~J}
\end{aligned}\)
& \mathrm{T}=1000^{\circ} \mathrm{C}=1273 \mathrm{~K} \\
& \Delta \dot{G}=\Delta \mathrm{H}-\mathrm{T} \Delta \mathrm{~S} \\
& =31400 \mathrm{~J}-1273 \mathrm{~K}\left(32 \mathrm{~J} \mathrm{~K}^{-1}\right) \\
& =3.1400 \mathrm{~J}-40736 \mathrm{~J} \\
& \Delta \mathrm{G}=-9336 \mathrm{~J}
\end{aligned}\)
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