Why does the temperature remain constant during a phase transition?

During a phase transition, such as melting or boiling, the energy supplied to the substance is not used to increase the kinetic energy of its particles, which would lead to a temperature rise. Instead, this energy is entirely utilized to overcome the intermolecular forces holding the particles in their current state. For instance, during melting, energy breaks the rigid lattice structure of a solid. During boiling, energy helps molecules escape the liquid phase into the gas phase. Once these forces are overcome, the particles gain more freedom of movement, but their average kinetic energy, and thus temperature, remains constant until the entire sample has completed the transition.

What is latent heat, and how is it related to enthalpy of phase transition?

Latent heat literally means 'hidden heat' because it's the heat absorbed or released during a phase change that doesn't cause a change in temperature. It's directly synonymous with the enthalpy of phase transition. Specifically, latent heat is the amount of heat absorbed or released per unit mass (or per mole) during a phase transition at constant temperature and pressure. For example, the latent heat of fusion is the enthalpy of fusion per unit mass or mole, representing the energy required to melt a substance.

How do intermolecular forces influence the magnitude of enthalpy of vaporization?

Intermolecular forces (IMFs) play a crucial role in determining the magnitude of the enthalpy of vaporization ($Delta H_{vap}$). Stronger intermolecular forces, such as hydrogen bonding or strong dipole-dipole interactions, require more energy to overcome. When a liquid vaporizes, its molecules must completely separate from each other, effectively breaking these IMFs. Therefore, substances with stronger IMFs will have higher enthalpies of vaporization because more energy is needed to pull their molecules apart and allow them to enter the gaseous phase.

Can enthalpy of phase transition be negative?

Yes, enthalpy of phase transition can certainly be negative. While processes like melting, vaporization, and sublimation are endothermic (absorb heat, $Delta H > 0$), their reverse processes are exothermic (release heat, $Delta H < 0$). For example, freezing (liquid to solid), condensation (gas to liquid), and deposition (gas to solid) all involve the formation of stronger intermolecular bonds and the release of energy to the surroundings. Thus, the enthalpy of freezing is $-Delta H_{fus}$, and the enthalpy of condensation is $-Delta H_{vap}$.

What is the difference between boiling and evaporation?

Both boiling and evaporation are processes where a liquid turns into a gas, but they differ significantly. Evaporation is a surface phenomenon that can occur at any temperature below the boiling point. Only molecules with sufficient kinetic energy at the liquid's surface escape into the gas phase. Boiling, on the other hand, is a bulk phenomenon that occurs at a specific temperature called the boiling point, where vapor bubbles form throughout the entire liquid. Boiling requires a continuous supply of heat at a constant temperature, whereas evaporation can occur slowly without external heating, drawing energy from the surroundings.

Enthalpy of Phase Transition

Q: What is the difference between boiling and evaporation?

Both boiling and evaporation are processes where a liquid turns into a gas, but they differ significantly. Evaporation is a surface phenomenon that can occur at any temperature below the boiling point. Only molecules with sufficient kinetic energy at the liquid's surface escape into the gas phase. Boiling, on the other hand, is a bulk phenomenon that occurs at a specific temperature called the boiling point, where vapor bubbles form throughout the entire liquid. Boiling requires a continuous supply of heat at a constant temperature, whereas evaporation can occur slowly without external heating, drawing energy from the surroundings.

Chemistry

NEET UG

Version 1Updated 22 Mar 2026

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Enthalpy of phase transition refers to the change in enthalpy that accompanies a phase change of a substance at constant temperature and pressure. These transitions involve a physical change in the state of matter, such as melting (solid to liquid), vaporization (liquid to gas), or sublimation (solid to gas). During such a process, the energy supplied or released is used to overcome or establish i…

Quick Summary

Enthalpy of phase transition, also known as latent heat, is the energy absorbed or released when a substance changes its physical state (solid, liquid, gas) at constant temperature and pressure. This energy is used to overcome or establish intermolecular forces, not to change the kinetic energy of particles, hence the constant temperature.

Key transitions include fusion (melting, solid to liquid, $Delta H_{fus} > 0$ ), vaporization (boiling, liquid to gas, $Delta H_{vap} > 0$ ), and sublimation (solid to gas, $Delta H_{sub} > 0$ ). Their reverse processes (freezing, condensation, deposition) are exothermic, having negative enthalpy values.

The magnitude of these enthalpy changes is primarily determined by the strength of intermolecular forces. Hess's Law relates these, stating $Delta H_{sub} = Delta H_{fus} + Delta H_{vap}$ . These concepts are vital for understanding energy changes in various natural and industrial processes, and for solving numerical problems in NEET that often combine specific heat calculations with phase change enthalpies.

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Key Concepts

Enthalpy of Fusion (

Delta H_{fus}

)

The enthalpy of fusion quantifies the energy required to disrupt the ordered structure of a solid and allow…

Enthalpy of Vaporization (

Delta H_{vap}

)

The enthalpy of vaporization represents the energy needed to completely separate liquid molecules from each…

Enthalpy of Sublimation (

Delta H_{sub}

)

Enthalpy of sublimation describes the direct transition of a solid to a gas without passing through the…

Key Formulas & Concepts:

Phase Change: — Constant T, energy for IMFs.

Heating/Cooling: —

q = m cdot c cdot Delta T

(temperature change).

Fusion (Melting): — Solid

ightarrow

Liquid,

Delta H_{fus} > 0

q = n cdot Delta H_{fus}

(or

m cdot L_{fus}

Vaporization (Boiling): — Liquid

ightarrow

Gas,

Delta H_{vap} > 0

q = n cdot Delta H_{vap}

(or

m cdot L_{vap}

Sublimation: — Solid

ightarrow

Gas,

Delta H_{sub} > 0

Delta H_{sub} = Delta H_{fus} + Delta H_{vap}

Reverse Processes: — Freezing, Condensation, Deposition are exothermic (

Delta H < 0

* Freezing: $-Delta H_{fus}$ * Condensation: $-Delta H_{vap}$ * Deposition: $-Delta H_{sub}$

Units: — Be careful with J vs. kJ, g vs. mol.

To remember the endothermic phase changes: My Very Solid Substance Melts, Vaporizes, Sublimes.

Melts (Fusion)

Vaporizes (Vaporization)

Sublimes (Sublimation)

All these processes require energy input (endothermic, $Delta H > 0$ ). Their opposites (Freezing, Condensation, Deposition) release energy (exothermic, $Delta H < 0$ ).

For the relationship: Sublimation Is Fusion Plus Vaporization (SIFPV) -> $Delta H_{sub} = Delta H_{fus} + Delta H_{vap}$ .