Final answer:
The change in internal energy of the system is -11.3 kJ, calculated using the first law of thermodynamics formula ΔU = Q - W, accounting for the heat released and work done on the system.
Step-by-step explanation:
The question pertains to the first law of thermodynamics, which is used to calculate the change in internal energy of a system. According to the first law, the change in internal energy (denoted as ΔU) is the difference between the heat added to the system (Q) and the work done by the system (W). However, in this scenario, work is done on the system, which should be considered as negative work for the calculation.
The formula to find the change in internal energy is: ΔU = Q - W. In the given problem, the system releases heat, so Q is negative, Q = -7.7 kJ, and the work done on the system is given as 3.6 kJ, thus W is positive, W = 3.6 kJ. Substituting into the equation, the change in internal energy is: ΔU = (-7.7 kJ) - (3.6 kJ) = -11.3 kJ.
Therefore, the internal energy of the system decreases by 11.3 kJ as a result of the processes described.