What is the change in internal energy of a system that takes in 2.40 ✕ 106 J of heat, and does 4.50 ✕ 105 J of work while dissipating 9.00 ✕ 106 J of heat?

To find the change in internal energy of a system, we need to apply the First Law of Thermodynamics, which states that the change in internal energy is equal to the heat added to the system minus the work done by the system.

In this case, the heat added to the system is 2.40 ✕ 10^6 J, the work done by the system is -4.50 ✕ 10^5 J (since work done by the system is negative), and the heat dissipated by the system is -9.00 ✕ 10^6 J (since heat dissipated is negative).

Therefore, we can calculate the change in internal energy (ΔU) as follows:

ΔU = Heat added - Work done - Heat dissipated
= 2.40 ✕ 10^6 J - (-4.50 ✕ 10^5 J) - (-9.00 ✕ 10^6 J)
= 2.40 ✕ 10^6 J + 4.50 ✕ 10^5 J + 9.00 ✕ 10^6 J

Adding these values together gives us the change in internal energy of the system.

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