To solve this problem, we need to understand the first law of thermodynamics, which states that the change in internal energy of a system is equal to the heat added to the system minus the work done by the system. In this case, the system is releasing heat, so the heat is negative, and work is being done on the system, so the work is positive.

The formula for the first law of thermodynamics is:

ΔU = Q - W

where:
ΔU is the change in internal energy,
Q is the heat added to the system, and
W is the work done by the system.

In this problem, we are given:

Q = -675 J (because the system is releasing heat)
W = 4.90 x 10^2 cal (work done on the system)

However, the work is given in calories, and we need to convert it to joules because the heat is given in joules and we want the answer in joules. We can use the conversion factor 1 cal = 4.184 J.

So, W = 4.90 x 10^2 cal x 4.184 J/cal = 2050.16 J

Now we can substitute Q and W into the formula:

ΔU = Q - W
ΔU = -675 J - 2050.16 J
ΔU = -2725.16 J

So, the change in internal energy of the system is -2725 J. The answer is negative because the system is losing energy. The answer has four significant figures because the given values have at most four significant figures.

Question

To solve this problem, we need to understand the first law of thermodynamics, which states that the change in internal energy of a system is equal to the heat added to the system minus the work done by the system. In this case, the system is releasing heat, so the heat is negative, and work is being done on the system, so the work is positive.

The formula for the first law of thermodynamics is:

ΔU = Q - W

where:
ΔU is the change in internal energy,
Q is the heat added to the system, and
W is the work done by the system.

In this problem, we are given:

Q = -675 J (because the system is releasing heat)
W = 4.90 x 10^2 cal (work done on the system)

However, the work is given in calories, and we need to convert it to joules because the heat is given in joules and we want the answer in joules. We can use the conversion factor 1 cal = 4.184 J.

So, W = 4.90 x 10^2 cal x 4.184 J/cal = 2050.16 J

Now we can substitute Q and W into the formula:

ΔU = Q - W
ΔU = -675 J - 2050.16 J
ΔU = -2725.16 J

So, the change in internal energy of the system is -2725 J. The answer is negative because the system is losing energy. The answer has four significant figures because the given values have at most four significant figures.

Knowee AI · Accepted Answer

To solve this problem, we need to understand the first law of thermodynamics, which states that the change in internal energy of a system is equal to the heat added to the system minus the work done by the system. In this case, the system is releasing heat, so the heat is negative, and work is being done on the system, so the work is positive.

The formula for the first law of thermodynamics is:

ΔU = Q - W

where:
ΔU is the change in internal energy,
Q is the heat added to the system, and
W is the work done by the system.

In this problem, we are given:

Q = -675 J (because the system is releasing heat)
W = 4.90 x 10^2 cal (work done on the system)

However, the work is given in calories, and we need to convert it to joules because the heat is given in joules and we want the answer in joules. We can use the conversion factor 1 cal = 4.184 J.

So, W = 4.90 x 10^2 cal x 4.184 J/cal = 2050.16 J

Now we can substitute Q and W into the formula:

ΔU = Q - W
ΔU = -675 J - 2050.16 J
ΔU = -2725.16 J

So, the change in internal energy of the system is -2725 J. The answer is negative because the system is losing energy. The answer has four significant figures because the given values have at most four significant figures.

What is the change in internal energy (in J) of a system that releases 675J of thermal energy to its surroundings and has ×4.90102cal of work done on it? Be sure your answer has the correct number of significant figures.

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