Enthalpy is the thermodynamic quantity H used in Physical Chemistry II to track heat flow at constant pressure. You usually work with ΔH to tell whether a process releases or absorbs heat.
Enthalpy is the thermodynamic state function H, defined as H = U + pV, where U is internal energy. In Physical Chemistry II, you use it because many chemical processes happen at roughly constant pressure, so the heat transferred is directly tied to the enthalpy change, ΔH.
That connection is what makes enthalpy so practical. If a reaction occurs in an open flask or in solution, pressure usually stays near atmospheric pressure. Under those conditions, q_p, the heat at constant pressure, equals ΔH. That means a calorimetry experiment can tell you the enthalpy change of a reaction without having to measure every microscopic energy transfer inside the system.
A positive ΔH means the system absorbs heat from the surroundings, so the process is endothermic. A negative ΔH means the system releases heat, so the process is exothermic. The sign comes from the system’s point of view, not the surroundings. That sign convention matters a lot when you are reading energy diagrams, balancing thermochemical equations, or interpreting whether a reaction warms or cools its environment.
Enthalpy is a state function, so only the initial and final states matter. The path does not. That is why Hess’s law works: if you can break a reaction into steps, the total ΔH is just the sum of the step enthalpies. In practice, this lets you build unknown reaction enthalpies from standard enthalpies of formation or from other tabulated data.
In irreversible thermodynamics, enthalpy is still part of the picture, but it is not the whole story. A real process can dissipate energy, produce entropy, and move away from equilibrium while still having a well-defined ΔH between initial and final states. So enthalpy tells you the heat bookkeeping, while entropy production tells you how much irreversibility came along with it.
Enthalpy is one of the main ways Physical Chemistry II tracks energy changes in reactions, phase changes, and calorimetry. If you know ΔH, you can tell whether a process is heating the surroundings or cooling them, and you can connect that result to measurable lab data instead of just abstract theory.
It also shows up when you compare reaction pathways. Because enthalpy is a state function, you can add stepwise reactions with Hess’s law and still get the correct total energy change. That makes enthalpy useful for building thermochemical cycles, estimating reaction heats from formation data, and checking whether a proposed pathway is energetically plausible.
The concept also sets up later ideas in the course. Internal energy, Gibbs free energy, and heat capacity all connect to enthalpy in different ways, so if you are shaky on H, the rest of thermodynamics gets harder fast. In irreversible thermodynamics, ΔH alone does not tell you how efficient a process is, but it is still part of the energy balance you need before you can talk about entropy production or dissipative losses.
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view galleryInternal Energy
Internal energy U is the energy stored inside the system from molecular motion and interactions. Enthalpy adds the pV term, which is why H is often the more convenient quantity at constant pressure. If you know U but forget the pressure-volume piece, your heat bookkeeping will be off for processes that involve expansion or compression.
Gibbs Free Energy
Gibbs free energy combines enthalpy and entropy through G = H - TS. That means enthalpy gives the heat side of the story, while Gibbs free energy tells you whether a process is thermodynamically favorable under constant temperature and pressure. In later problems, you often need both to decide if a reaction can proceed spontaneously.
Heat Capacity
Heat capacity links temperature change to heat flow, and that makes it one of the main tools for measuring enthalpy changes in the lab. If you know the heat absorbed by a solution or calorimeter, you can convert that into ΔH for a reaction. The better you handle heat capacity, the better your enthalpy calculations will be.
dissipative processes
Dissipative processes turn organized energy into dispersed energy, often with entropy production. Enthalpy still tracks the heat exchanged, but it does not tell you how much energy was irreversibly spread out during the process. That distinction matters in irreversible thermodynamics, where a process can have a clean ΔH and still be inefficient.
A problem set question might give you calorimetry data, a reaction equation, or a thermochemical cycle and ask for ΔH. Your job is to use sign conventions correctly, decide whether the process is endothermic or exothermic, and combine step enthalpies with Hess’s law when needed.
In a lab report, you may calculate enthalpy from temperature change and heat capacity, then explain whether the system or surroundings gained heat. On a quiz, you might also compare enthalpy to internal energy or identify why constant-pressure conditions let q_p equal ΔH. If the topic is irreversible thermodynamics, expect to describe the heat flow while also noting that enthalpy does not measure entropy production.
Internal energy, U, is the total microscopic energy inside the system. Enthalpy, H = U + pV, includes that energy plus the pressure-volume term. They are not interchangeable, especially when a process involves expansion, compression, or constant-pressure heat flow. If a question asks for heat at constant pressure, enthalpy is usually the better quantity.
Enthalpy is the state function H = U + pV, and in Physical Chemistry II it is the go-to quantity for constant-pressure heat flow.
The sign of ΔH tells you the direction of heat transfer, with negative values for exothermic processes and positive values for endothermic ones.
Because enthalpy is a state function, you can add reaction steps with Hess’s law and still get the same overall ΔH.
Calorimetry problems often convert measured heat into enthalpy change, so heat capacity and sign conventions matter a lot.
Enthalpy tells you the heat bookkeeping, but it does not tell you everything about irreversibility or entropy production.
Enthalpy is the thermodynamic quantity H = U + pV. In Physical Chemistry II, you use ΔH to track heat flow for reactions and phase changes, especially when the pressure stays constant. That is why it shows up so often in calorimetry and thermochemical calculations.
Not exactly. Heat is energy transferred because of a temperature difference, while enthalpy is a state function that helps you account for that heat under constant pressure. For many chemistry problems at constant pressure, q_p equals ΔH, which is why the two get linked so often.
If ΔH is negative, the system releases heat and the process is exothermic. If ΔH is positive, the system absorbs heat and the process is endothermic. The sign is from the system’s perspective, so always check whether the system or the surroundings are gaining heat.
You break the target reaction into smaller reactions whose enthalpy changes you already know, then add the ΔH values together. Because enthalpy is a state function, the total depends only on the starting and ending states, not the path. That makes Hess’s law especially useful for reaction cycles and formation data.