# 1st Law of Thermodynamics

1st law of thermodynamics state that whenever energy supplied to a system used for two purposes:

1. A part of energy use to raise the internal energy (temperature) of the body.
2. Rest of the energy used in doing external work. ## Relation in 1st law of thermodynamics

The first law gives the relation between heat, internal energy, and work. if

• ∆Q = heat energy absorbed by a system from the surroundings.
• ∆W = Work done(external) by the system on the surroundings.
• ∆U = Change in internal energy(I.E).
• Total Energy in the beginning: Heat absorb+I.E

If there is ∆Q heat energy absorbed by a system, then there is a rise in ∆U internal energy and do ∆W external work done by it, then all the quantities provided are measured in units of work. It means the first law is similar to the law of conservation of energy.

Mathematically,

ΔQ = ΔU + ΔW                      (Remember: W = –PΔV) ## Sign Convention

1.When heat is supplied to a system, ΔQ is considered as positive. When the heat is taken out from the system, then ΔQ is taken as negative.

2.If work is done by the surroundings(compression of a gas), W is taken as positive

ΔQ = ΔU+W.

During the expansion of a gas, W is taken as negative so that

ΔQ = ΔU – ΔW.

3.When the temperature of the gas increases, its I.E energy increases, ΔU is taken as positive.When the temperature of the gas decreases, its I.E decreases, ΔU is considered negative.

Heat and work are two different methods of energy transfer in a system that results in changes in internal energy.

(A) heat energy transfer due to the temperature difference between the system and the surroundings.

(B) The energy is transmitted through the work.For example, the piston moves (some weight associated with it).

Must Read: Zeroth law.

## Limitations

1.The 1st law does not tell the direction in which the change can occur.

2.The first law does not give any idea about the limit of change.

3.The 1st rule of thermodynamics does not give any information about the source of heat.

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