(B) The entropy change of the surrounding $(\Delta S_{surr})$ is given by the formula: $\Delta S_{surr} = \frac{-q_{sys}}{T}$. Given that the reactio

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(B) The entropy change of the surrounding $(\Delta S_{surr})$ is given by the formula: $\Delta S_{surr} = \frac{-q_{sys}}{T}$. Given that the reaction releases heat,$q_{sys} = -525 \ kJ = -525000 \ J$. Therefore,the heat absorbed by the surrounding is $q_{surr} = +525000 \ J$. The temperature $T = 300 \ K$. Substituting the values: $\Delta S_{surr} = \frac{525000 \ J}{300 \ K} = 1750 \ J \ K^{-1}$.

Class 11 Chemistry Thermodynamics 2 nd Law of thermodynamics and Entropy

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Calculate the entropy change of the surrounding if $2 \ moles$ of $H_2$ and $1 \ mole$ of $O_2$ gas combine to form $2 \ moles$ of liquid water by releasing $525 \ kJ$ of heat to the surrounding at constant pressure and at $300 \ K$. (in $J \ K^{-1}$)

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A
$1700$
B
$1750$
C
$1800$
D
$1650$

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Thermodynamics

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2 nd Law of thermodynamics and Entropy

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Calculate the entropy change of the surrounding if $2 \ moles$ of $H_2$ and $1 \ mole$ of $O_2$ gas combine to form $2 \ moles$ of liquid water by releasing $525 \ kJ$ of heat to the surrounding at constant pressure and at $300 \ K$. (in $J \ K^{-1}$)

Similar Questions

Entropy is maximum in case of

In which of the following processes is $\Delta S$ positive?

Two moles of an ideal gas undergo free expansion from $10 \ L$ to $100 \ L$ at $300 \ K$. The values of $\Delta S_{system}$ and $\Delta S_{surroundings}$ are ($R$ is universal gas constant)

The enthalpy of vaporization of benzene is $30.779 \ kJ/mol$ and its boiling point is $353 \ K$. Find the change in entropy for the conversion of liquid benzene to its vapour at that temperature.

Will ice at $273 \ K$ placed in a surrounding at $298 \ K$ melt to form water at $273 \ K$? Prove this statement. The molar enthalpy of fusion of ice is $6.025 \ kJ \ mol^{-1}$.

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