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Thermochemistry studies the heat exchanged in chemical reactions. Enthalpy (H) is a state function; at constant pressure, ΔH equals heat flow: ΔH < 0 for exothermic reactions (heat released) and ΔH > 0 for endothermic. Hess's law states that ΔH for a reaction is path-independent — it equals the sum of ΔH values for any sequence of steps that add up to the overall reaction. Standard enthalpies of formation (ΔHf°) — enthalpy change for forming one mole of a compound from elements in standard state — provide a systematic table-based method: ΔH°rxn = ΣΔHf°(products) − ΣΔHf°(reactants).
Practice Hess's law by combining given equations algebraically (reversing equations changes sign of ΔH; multiplying changes magnitude proportionally). Use formation enthalpy tables with the products-minus-reactants formula. Connect to calorimetry through q = mcΔT.
Thermochemistry is the study of how energy flows when chemical reactions occur. The central concept is enthalpy (H), a thermodynamic state function designed to capture heat exchange at constant pressure. You don't need to know the absolute value of H for any substance — what matters is the change, ΔH, between products and reactants. When ΔH < 0, the system releases heat to the surroundings (exothermic — like combustion); when ΔH > 0, it absorbs heat from the surroundings (endothermic — like dissolving ammonium nitrate in water). The sign convention is always from the system's perspective: negative means energy leaves the system.
One key subtlety: enthalpy is not the total energy of a system. Formally, H = U + PV, where U is internal energy and PV is a pressure-volume correction. At constant pressure (most lab reactions open to the atmosphere), ΔH equals q, the heat transferred. This is why you can connect thermochemistry to calorimetry — q = mcΔT gives you the heat absorbed by the surroundings, which equals −ΔH for the reaction. At constant volume (a sealed bomb calorimeter), the situation is different: ΔH ≠ q because there is no PV work. Keeping track of conditions matters.
Hess's Law is the most practically useful tool in this topic: because enthalpy is a state function, ΔH depends only on initial and final states, not on the path. You can combine thermochemical equations algebraically. Reverse a reaction and its ΔH changes sign. Multiply a reaction by a scalar and its ΔH scales by the same factor. Add reactions step by step until they sum to your target reaction, and sum the ΔH values — the result is ΔH for the overall reaction. Hess's Law lets you calculate enthalpies for reactions that are difficult or impossible to measure directly.
Standard enthalpies of formation (ΔHf°) give you a systematic, table-based approach that is essentially Hess's Law pre-packaged. A formation enthalpy is the ΔH for forming one mole of a compound from its constituent elements in standard state. Elements in standard state have ΔHf° = 0 by definition. The formula ΔH°rxn = ΣΔHf°(products) − ΣΔHf°(reactants) works by conceptually decomposing reactants into elements and then assembling products from those elements, with the element steps canceling to zero.
A persistent sign-confusion error: exothermic reactions have negative ΔH, and students sometimes think the system is 'losing energy' in a bad or incomplete sense. Think of it via conservation of energy: the bonds in the products store less chemical energy than the bonds in the reactants, and the difference is released as heat. The system's energy decreases; the surroundings' energy increases by the same amount. Total energy is conserved — fully consistent with the first law of thermodynamics. The negative sign on ΔH is not a deficit; it is a direction indicator showing which way heat flows.