The first law of thermodynamics is based on:

When a system is moved from state \(a\) to state \(b\) along the path \(acb\), it is discovered that the system absorbs \(200~\text{J}\) of heat and performs \(80~\text{J}\) of work.  Along the path \(adb\), heat absorbed \(Q =144~\text{J}\). The work done along the path \(adb\) is:

The latent heat of vaporisation of water is \(2240~\text{J/gm}\). If the work done in the process of expansion of \(1~\text{g}\) is \(168~\text{J}\), then the increase in internal energy is:
1. \(2408~\text{J}\)
2. \(2240~\text{J}\)
3. \(2072~\text{J}\)
4. \(1904~\text{J}\)

An ideal gas goes from state \(A\) to state \(B\) via three different processes, as indicated in the \(P\text-V\) diagram. If \(Q_1,Q_2,Q_3\)${\mathrm{}}_{}$ indicates the heat absorbed by the gas along the three processes and \(\Delta U_1, \Delta U_2, \Delta U_3\) indicates the change in internal energy along the three processes respectively, then:

An ideal monoatomic gas \(\left(\gamma = \frac{5}{3}\right )\) absorbs 50 cal in an isochoric process. The increase in internal energy of the gas is:

1 kg of gas does 20 kJ of work and receives 16 kJ of heat when it is expanded between two states. The second kind of expansion can be found between the same initial and final states, which requires a heat input of 9 kJ. The work done by the gas in the second expansion will be:

If a gas undergoes the change in its thermodynamic state from A to B via two different paths, as shown in the given pressure (P) versus volume (V) graph, then: