Q.What amount of heat must be supplied to $2.0 \times 10^{-2}\ \text{kg}$ of nitrogen (at room temperature) to raise its temperature by $45\,^{\circ}\text{C}$ at constant pressure? (Molecular mass of $\text{N}_2 = 28$; $R = 8.3\ \text{J mol}^{-1}\ \text{K}^{-1}$.)

Heat Capacity at Constant Pressure — From Intuition to Precision

Imagine you have a pot of water on a stove. You turn the burner on, and the water gets hotter. How much heat does it take to raise its temperature by, say, 10°C? That depends on two things: how much water you have, and whether the pot is open to the air or sealed tight.

If the pot is open (constant pressure — the air above it is always at atmospheric pressure), the water can expand as it heats. Some of the energy you supply goes into pushing the atmosphere aside — doing work against the outside air. So you need to put in more heat than if the pot were sealed (constant volume), where no expansion work is possible.

That extra heat is the key idea behind heat capacity at constant pressure, denoted $C_p$.

The Intuition First

Heat capacity tells you: "How much heat must I add to raise the temperature of this substance by 1°C (or 1 K)?"

• At constant volume ($C_v$): All the heat goes into increasing the internal energy (the kinetic and potential energy of the molecules). No work is done because the volume doesn't change.
• At constant pressure ($C_p$): Some heat goes into internal energy, but some also goes into the work of expansion against the constant external pressure. So $C_p$ is always larger than $C_v$ for gases (and for most solids/liquids, the difference is tiny because they barely expand).

The Precise Statement

Heat capacity at constant pressure is defined as the amount of heat required to raise the temperature of a substance by 1 K (or 1°C) while keeping the pressure constant.

Mathematically:

The subscript $p$ means "at constant pressure." The $\delta Q$ (not $dQ$) reminds us that heat is a path-dependent quantity, not a state function.

But we can rewrite this in terms of a state function — enthalpy ($H$). At constant pressure, the heat added equals the change in enthalpy:

Therefore:

This is the working definition you'll use in problems: $C_p$ is the partial derivative of enthalpy with respect to temperature at constant pressure.

Molar vs. Specific Heat Capacity

You'll encounter two common forms:

• Molar heat capacity at constant pressure ($C_{p,m}$): heat capacity per mole (units: J mol⁻¹ K⁻¹)
• Specific heat capacity at constant pressure ($c_p$): heat capacity per unit mass (units: J kg⁻¹ K⁻¹)

The total heat capacity of a sample is:

Why It Matters

In most chemical reactions and physical processes, the system is open to the atmosphere — constant pressure. So $C_p$ is the relevant quantity for:

• Calculating enthalpy changes ($\Delta H = n C_{p,m} \Delta T$)
• Designing calorimeters (like coffee-cup calorimeters that operate at constant pressure)
• Understanding why gases heat up when compressed (and cool when expanded)

A Quick Example

How much heat is needed to raise the temperature of 2 moles of an ideal gas from 300 K to 400 K at constant pressure? (Given $C_{p,m} = 29.1 \text{ J mol}^{-1} \text{K}^{-1}$)

If the same gas were heated at constant volume, you'd need less heat — about $5820 - nR\Delta T = 5820 - (2)(8.314)(100) = 4157 \text{ J}$ — because no expansion work is done.

Final takeaway: $C_p$ is the heat capacity you measure when the system is free to expand against a constant external pressure. It's always larger than $C_v$ for gases, and the difference comes from the work of expansion.

Students searching for "Heat Capacity at Constant Pressure: Definition, Formula & Real-World Examples" or "Heat Capacity at Constant Pressure 11 physics" will find this explanation directly aligned with the Class 11 Physics curriculum prescribed under NCERT/CBSE. It is also a recurring theme in JEE Main, NEET and state engineering/medical entrance exams, so working through it carefully pays off well beyond board exams.