ΔH

Meaning: Enthalpy (H) is a state function equal to the system's internal energy plus PV. ΔH is the change in enthalpy, and at constant pressure it equals the heat q exchanged with the surroundings. A negative ΔH means the products are lower in enthalpy than the reactants, so heat flows out (exothermic). A positive ΔH means heat flows in (endothermic). Because H is a state function, ΔH depends only on initial and final states, which is why Hess's law works: ΔH values of steps add up to the overall ΔH.

Equation: ΔH = Hproducts − Hreactants = qp, where H is enthalpy (J) and qp is heat at constant pressure (J). From bond energies: ΔH ≈ Σ(bonds broken) − Σ(bonds formed), with each term in J/mol. For calorimetry: q = mcΔT, with m in kg, c in J/(kg·K), and ΔT in K.

On the MCAT: Expect Hess's law sums, reading energy diagrams (products below reactants means ΔH < 0), bond-energy estimates, and coffee-cup calorimetry. Reversing a reaction flips the sign of ΔH, and multiplying coefficients scales it. ΔH also enters ΔG = ΔH − TΔS.

Common confusion: Breaking bonds requires energy and forming bonds releases it. Students reverse this. A reaction is exothermic when the bonds formed are stronger overall than the bonds broken. Also, ΔH does not by itself decide spontaneity, since ΔS and T matter through ΔG, and it is unrelated to activation energy.