Embibe Experts Solutions for Chapter: Thermodynamics, Exercise 1: BEGINNER'S BOX

One mole of an ideal gas is expanded isothermally from an initial volume of $1 \mathrm{litre}$ to $10 \mathrm{litres}$. The $∆\mathrm{E}$ for this process is: $\left(\mathrm{R} = 2 \mathrm{cal} {\mathrm{mol}}^{-1} {\mathrm{K}}^{-1}\right)$.

In an isochoric process, the increase in internal energy is:

When a solid melts, there is:

For the reaction, ${\mathrm{N}}_2\left(\mathrm{g}\right)+3{\mathrm{H}}_2\left(\mathrm{g}\right)⟶2{\mathrm{NH}}_3\left(\mathrm{g}\right)$. Which of the following is correct?

Latent heat of vaporisation of a liquid at $500 \mathrm{K}$ and $1 \mathrm{atm}$ pressure is $10.0 {\mathrm{kcal}}^{-1} {\mathrm{mol}}^{-1}$. What will be the change in internal energy $\left(\mathrm{\Delta E}\right)$ of $3 \mathrm{mol}$ of liquid at the same temperature?

$5$ mole of an ideal gas expand reversibly from a volume of $8 {\mathrm{dm}}^3$ to $80 {\mathrm{dm}}^3$ at a temperature of $27°\mathrm{C}$. The change in entropy is :

The latent heat of vaporization of water at $100°\mathrm{C}$ is $540 \mathrm{cal} {\mathrm{g}}^{-1}.$  Calculate the increase in entropy when one mole of water at $100°\mathrm{C}$ has evaporated.

If. ${\mathrm{\Delta G}}^{\circ }>0$ for a reaction then :