What is the change in internal energy of a system that does 7.02 kj of work and absorbs 888 j of heat?
delta E = q + w
7.02 is negative since the system does work. q is + since the system absorbs heat.
To determine the change in internal energy of a system, you can use the first law of thermodynamics, which states that the change in internal energy (ΔU) is equal to the heat (Q) absorbed or released by the system minus the work (W) done by or on the system.
The formula is: ΔU = Q - W
Given:
Q = 888 J (heat absorbed by the system)
W = 7.02 kJ = 7.02 × 10^3 J (work done by the system)
Substituting the given values into the formula, we can calculate the change in internal energy:
ΔU = 888 J - 7.02 × 10^3 J
To perform the subtraction, we need to convert 7.02 kJ into J:
7.02 kJ = 7.02 × 10^3 J
Now we can calculate the change in internal energy:
ΔU = 888 J - 7.02 × 10^3 J
Simplifying the equation further:
ΔU = -6.13 × 10^3 J
Therefore, the change in internal energy of the system is -6.13 × 10^3 J.
To find the change in internal energy of a system, we can use the first law of thermodynamics, also known as the law of conservation of energy:
ΔU = Q - W
where ΔU is the change in internal energy, Q is the heat absorbed by the system, and W is the work done by the system.
Given:
Q = 888 J (heat absorbed by the system)
W = 7.02 kJ (work done by the system)
First, we need to convert the work value from kilojoules to joules since the heat is given in joules.
1 kJ = 1000 J
So, 7.02 kJ = 7.02 x 1000 J = 7020 J
Now, we can plug the values into the formula:
ΔU = Q - W
ΔU = 888 J - 7020 J
Calculating the difference:
ΔU = -6132 J
So, the change in internal energy of the system is -6132 J. The negative sign implies that the system has lost energy.