First Law of Thermodynamics
Definition and meaning of First Law of Thermodynamics in chemistry.
The first law of thermodynamics says energy cannot be created or destroyed, only changed from one form to another. For a closed system, the change in internal energy equals the heat added to the system plus the work done on it.
In more detail
Written as an equation, this is ΔU = q + w. Here, q stands for heat transferred into the system, and w stands for work done on the system. This law is really just the law of conservation of energy, applied to heat and chemical processes.
It is a starting assumption in chemistry, not something proven from a simpler rule, and it forms the basis for calorimetry and reaction energy calculations. The first law also leads to the idea of enthalpy, written as H = U + pV. This idea simplifies energy bookkeeping for reactions run at constant pressure.
One thing the first law cannot do is predict which direction a process will naturally go. It only tracks energy amounts. Figuring out spontaneous direction is a job for the second law instead.
A common misconception is that the first law only applies to gases in a piston. It applies to any closed system, including chemical reactions in a flask, biological processes in a cell, and heat flowing through a solid object. The law also underlies calorimetry experiments, where chemists measure the heat released or absorbed by a reaction to find its energy change directly.
Key facts
| Field | Physical Chemistry |
|---|---|
| Equation | ΔU = q + w |
| Also Known As | Law of conservation of energy (thermodynamic form) |
| Status | Fundamental postulate, not derived from other laws |
| Related Concept | Enthalpy, H = U + pV |
A gas trapped in a piston absorbs 500 joules of heat while expanding and pushing 200 joules of work onto its surroundings. Using the first law, ΔU = q + w = 500 J + (-200 J) = 300 J. So the gas's internal energy goes up by 300 joules overall.
Frequently asked questions
How does the first law connect to enthalpy?
Enthalpy is defined as H = U + pV. At constant pressure, the heat exchanged in a process equals ΔH, making enthalpy a practical version of the first law for reactions run in open containers.
Does the first law explain why reactions proceed in a particular direction?
No. The first law only accounts for energy conservation. It cannot predict spontaneity. That requires the second law of thermodynamics and the concept of entropy.
What is the sign convention in ΔU = q + w?
In this common convention, heat added to the system counts as positive and work done on the system counts as positive too. Heat released or work done by the system on its surroundings is written as negative.