What's the Difference Between Isovolumetric, Isobaric, Isothermal, and Adiabatic Processes?

These four thermodynamic processes differ in what remains constant or what energy transfer is prevented. In IB Physics B.4 Thermodynamics, isovolumetric means constant volume, isobaric means constant pressure, isothermal means constant temperature, and adiabatic means no thermal energy transfer.

How the Processes Differ

Use the IB sign convention for the first law of thermodynamics:

$$Q=\Delta U+W$$

Here, $Q$ is thermal energy transferred to the gas, $\Delta U$ is its change in internal energy, and $W$ is work done by the gas.

For an ideal gas undergoing an isothermal process, $PV=\text{constant}$. For an adiabatic process involving a monatomic ideal gas, the IB data booklet gives

$$PV^{5/3}=\text{constant}.$$

An adiabatic curve on a $P$-$V$ diagram is therefore steeper than an isothermal curve because the gas temperature changes during adiabatic expansion or compression.

For example, suppose a gas expands isobarically at $2.0\times10^5\ \text{Pa}$ from $3.0\times10^{-3}\ \text{m}^3$ to . The work done is . If , then . This shows how supplied energy is divided.

A common misconception is that adiabatic means constant temperature. It does not: adiabatic means $Q=0$, while isothermal means $\Delta T=0$. An isothermal expansion requires thermal energy to enter the gas to replace the energy transferred as work.

Exam Technique

This is HL-only content. In an IB response, first identify the defining condition, then apply $Q=\Delta U+W$. For calculations, show the condition explicitly, such as $W=0$, $Q=0$, or $$, before substituting values.