First Law of Thermodynamics

Energy can neither be created nor destroyed. Energy only changes from one form to another. This is the observation made by many scientists over the years. This observation has taken the form of first law of thermodynamics.

First law of thermodynamics states that the energy can neither be created nor destroyed. The total energy of the universe or an isolated system is constant.

Mathematically, the first law of thermodynamics is stated as:

ΔU = q + w

where ΔU is change in internal energy, q is heat absorbed by the system, and w is work done on the system.

Internal Energy (U)

The internal energy is defined as the sum of the energies of all the atoms, molecules or ions contained in the system. It is a state variable. It is not possible to measure the absolute values of internal energy. However, you can calculate the change in internal energy.

If the internal energy of the system in the initial state is U1 and that in the final state is U2, then change in internal energy ΔU is independent of the path taken from the initial to the final state.

ΔU = U2 – U1

The internal energy of the system can be changed in two ways:

  1. either by allowing heat to flow into the system or out of the system
  2. by work done on the system or by the system

Heat (q) and Work (w)

Heat and work are not state functions. This is because the values of both q and w depend upon the way in which the change is carried out. 

Any thing which increases the internal energy of a system is given a positive sign. Heat given to the system (q) and work done on the system (w) are given positive signs.

For example, if a certain change is accompanied by absorption of 50 kJ of heat and expenditure of 30 kJ of work,

q= +50 kJ

w= –30 kJ

Change in internal energy ΔU = (+50 kJ) + (–30 kJ) = +20 kJ

Work of Expansion

If the pressure p is constant and the volume of the system changes from V1 to V2. The work done by a system is given as:

w = – p(V2 – V1) = – pΔV

Minus sign is taken because the work is done by the system.

ΔU = q – pΔV

If the process is carried out at constant volume, then ΔV = 0

ΔU = qv

You can determine internal energy change if you measure the heat gained or lost by the system at constant volume. However, in chemistry, the chemical reactions are generally carried out at constant pressure (atmospheric pressure). So, another state function, enthalpy, is used.