enthalpy
Summary
Enthalpy (H) is the heat content of a system, H = U + PV. At constant pressure, the heat absorbed or released equals ΔH: negative is exothermic, positive is endothermic. It is a state function, so ΔH depends only on initial and final states, which is why Hess's law works.
Detail
Meaning: Enthalpy (H) is defined as H = U + PV, where U is internal energy. It is a state function, so only the initial and final states matter, not the path. At constant pressure, the heat exchanged equals ΔH. ΔH < 0 is exothermic (system releases heat, surroundings warm); ΔH > 0 is endothermic (system absorbs heat, surroundings cool). Breaking bonds absorbs energy and forming bonds releases it.
Equation: ΔH°rxn = ΣnΔH°f(products) − ΣmΔH°f(reactants), where ΔH°rxn is the standard enthalpy change of reaction and ΔH°f is the standard enthalpy of formation of each species (both in J/mol, often written in kJ/mol), and n and m are the stoichiometric coefficients.
On the MCAT: Expect Hess's law problems: add step equations, flip the sign of ΔH when a reaction is reversed, and multiply ΔH when the equation is multiplied. Bond-energy problems use ΔH ≈ energy of bonds broken − energy of bonds formed. Calorimetry (q = mcΔT of the surrounding water) gives ΔH at constant pressure, with the opposite sign. Energy diagrams show ΔH as product level minus reactant level.
Common confusion: A negative ΔH does not by itself make a reaction spontaneous; ΔG = ΔH − TΔS decides, so entropy and temperature matter. ΔH also says nothing about rate, which is set by activation energy.
Sources
- Chemistry: The Central Science
- Chemistry (Zumdahl)
- General Chemistry (Petrucci)
Written for the MCAT and checked before it was published. For study reference only. Spotted an error? Let us know.