What Is Enthalpy
Why Enthalpy Exists
Most processes in the real world occur at constant pressure rather than constant volume. Chemical reactions in open containers, biological processes in living organisms, and atmospheric phenomena all happen at or near atmospheric pressure. Under these conditions, the first law becomes Q{sub}p{/sub} = dU + PdV. The heat absorbed at constant pressure accounts for both the change in internal energy and the expansion work done against the atmosphere.
Rather than carrying the expression dU + PdV through every calculation, it is convenient to define a new quantity H = U + PV so that dH = Q{sub}p{/sub}. This simplification is the entire reason enthalpy exists. It is not a new form of energy but a bookkeeping tool that automatically includes the PV work that accompanies constant-pressure processes. For constant-volume processes, internal energy U remains the more natural variable, since dU = Q{sub}v{/sub}.
The PV term in enthalpy can be thought of as the energy needed to make room for the system against its surroundings at pressure P. When a gas is created in a chemical reaction at constant pressure, the atmosphere must be pushed back to accommodate the new gas. The PV work required for this is automatically included in the enthalpy change, so the heat measured in a constant-pressure calorimeter directly equals the enthalpy change.
Enthalpy Changes in Chemical Reactions
The standard enthalpy of reaction (delta H) is the enthalpy change when reactants in their standard states convert to products in their standard states at a specified temperature (usually 298 K). A negative delta H indicates an exothermic reaction (heat is released), while a positive delta H indicates an endothermic reaction (heat is absorbed). The combustion of methane, for example, has delta H = -890 kJ/mol, meaning 890 kJ of heat is released per mole of methane burned.
Hess law states that the enthalpy change for a reaction is the same whether it occurs in one step or in a series of steps. This is a direct consequence of enthalpy being a state function. You can calculate the enthalpy of a reaction you cannot measure directly by combining reactions whose enthalpies are known. This principle underlies the use of standard enthalpies of formation, which tabulate the enthalpy change when one mole of a compound is formed from its elements in their standard states.
Bond enthalpy analysis provides another route to estimating reaction enthalpies. Breaking bonds requires energy (endothermic), and forming bonds releases energy (exothermic). The overall enthalpy change is approximately the sum of bond enthalpies broken minus the sum of bond enthalpies formed. This method is less accurate than using formation enthalpies because bond energies vary depending on the molecular environment, but it provides useful estimates when tabulated formation data is unavailable.
Enthalpy and Phase Transitions
Phase transitions involve large enthalpy changes at constant temperature. The enthalpy of fusion (melting) is the energy required to convert a solid to a liquid at the melting point. For water, this is 6.01 kJ/mol (334 J/g). The enthalpy of vaporization is the energy required to convert a liquid to a gas at the boiling point. For water, this is 40.7 kJ/mol (2260 J/g), much larger than the fusion enthalpy because vaporization requires completely separating molecules from each other.
The large enthalpy of vaporization of water has profound consequences for climate and weather. When water evaporates from the ocean surface, it absorbs a large amount of energy from the surrounding water, cooling the surface. When that water vapor condenses in clouds, it releases that same energy as heat, warming the atmosphere and powering storms. Hurricanes are essentially heat engines driven by the enthalpy of vaporization of ocean water.
Sublimation (solid directly to gas) combines the enthalpies of fusion and vaporization. The enthalpy of sublimation equals the sum of the fusion and vaporization enthalpies, by Hess law. Freeze-drying exploits sublimation to remove water from food and pharmaceuticals at low temperatures, preserving delicate structures that would be damaged by liquid-phase processes.
Enthalpy simplifies energy accounting at constant pressure by automatically including expansion work. Every heat value you see on a food label or in a chemistry textbook is an enthalpy change.
Enthalpy in Engineering Applications
In engineering, enthalpy is the primary energy variable for analyzing open systems (systems with mass flowing in and out). The steady-state energy equation for an open system involves the enthalpy of entering and leaving streams rather than internal energy, because the flow work (PV work needed to push fluid through the system boundaries) is already included in enthalpy. Turbines, compressors, heat exchangers, and nozzles are all analyzed using enthalpy balances.
Steam tables and refrigerant tables list enthalpy as a function of temperature and pressure for working fluids used in power generation and refrigeration. Engineers use these tables to calculate the work output of turbines, the heat input to boilers, and the cooling capacity of refrigeration systems. The Mollier diagram (an enthalpy-entropy chart) provides a graphical tool for visualizing and analyzing thermodynamic processes in power and refrigeration cycles.
In atmospheric science, moist enthalpy (also called moist static energy) combines sensible heat (cpT), gravitational potential energy (gz), and latent heat (Lq, where L is the latent heat of vaporization and q is the specific humidity). This quantity is approximately conserved in adiabatic atmospheric processes, making it a powerful diagnostic tool for analyzing tropical convection, monsoons, and other large-scale weather patterns.