endothermic
Summary
An endothermic process absorbs heat from its surroundings, so the system's enthalpy rises (ΔH > 0) and the surroundings cool. Most testable point: raising the temperature shifts an endothermic equilibrium toward products, and K increases.
Detail
Meaning: In an endothermic process the system takes in heat, so the products have higher enthalpy than the reactants. Overall, more energy is needed to break bonds than is released when new bonds form. Melting, vaporization, sublimation and the dissolving of salts such as NH₄NO₃ in water are examples. Heat can be treated like a reactant: reactants + heat ⇌ products.
Equation: ΔH = Hproducts − Hreactants > 0, where ΔH is the enthalpy change in J/mol. At constant pressure, ΔH equals the heat q absorbed by the system, in J. From bond energies, ΔH ≈ Σ(energy of bonds broken) − Σ(energy of bonds formed).
On the MCAT: Expect an energy diagram with products higher than reactants, a heating curve with plateaus at phase changes, or a calorimetry setup in which the surroundings (the water) get colder, as in an instant cold pack. For equilibrium questions, apply Le Châtelier's principle: adding heat favors the endothermic direction, and K for an endothermic reaction increases as T increases.
Common confusion: Endothermic does not mean nonspontaneous. Spontaneity depends on ΔG = ΔH − TΔS, so an endothermic process can proceed when ΔS is positive and TΔS is larger than ΔH, as when ice melts above 0 °C. Also, the sign is from the system's point of view: the system gains heat while the surroundings lose it.
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.