The First Law of Thermodynamics states that the change in internal energy of a system equals the heat absorbed by the system minus the work done by the system on its surroundings: (\Delta \cup = q - w). Here, the gas absorbs heat so (q = +450) J (positive, heat flows into the system). The gas does work on the surroundings during expansion, so (w = +120) J (work done BY the system is positive). Applying the first law: (\Delta \cup = 450 - 120 = 330) J. Option 570 J is incorrect because it adds heat and work instead of subtracting work done by the system. Option -330 J would result if the gas were releasing heat instead of absorbing it. Option -570 J arises from both incorrect sign conventions applied simultaneously. This is a core NCERT Chapter 6 application of the first law of thermodynamics. Plausibility check: since heat absorbed exceeds work done, internal energy must increase, confirming (\Delta \cup > 0).