Thermodynamics first law

The first law of thermodynamics is a version of the law of conservation of energy, adapted for thermodynamic processes, distinguishing two kinds of transfer of energy, as heat and as thermodynamic work, and relating them to a function of a body's state, called internal energy.
The law of conservation of energy states that the total energy of an isolated system is constant; energy can be transformed from one form to another, but can be neither created nor destroyed.
For a thermodynamic process without transfer of matter, the first law is often formulated




Δ
U
=
Q
−
W


{\displaystyle \Delta U=Q-W}
,where



Δ
U


{\displaystyle \Delta U}
denotes the change in the internal energy of a closed system,



Q


{\displaystyle Q}
denotes the quantity of energy supplied to the system as heat, and



W


{\displaystyle W}
denotes the amount of thermodynamic work done by the system on its surroundings. An equivalent statement is that perpetual motion machines of the first kind are impossible.
For processes that include transfer of matter, a further statement is needed: 'With due account of the respective reference states of the systems, when two systems, which may be of different chemical compositions, initially separated only by an impermeable wall, and otherwise isolated, are combined into a new system by the thermodynamic operation of removal of the wall, then





U

0


=

U

1


+

U

2




{\displaystyle U_{0}=U_{1}+U_{2}}
,where




U

0




{\displaystyle U_{0}}
denotes the internal energy of the combined system, and




U

1




{\displaystyle U_{1}}
and




U

2




{\displaystyle U_{2}}
denote the internal energies of the respective separated systems.'

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