Standard Enthalpy
Standard enthalpy refers to the enthalpy change of a process measured under standard conditions: 298.15 K (25 °C) and 1 bar (approximately 1 atm) pressure, with all substances in their standard states. Standard enthalpy values provide a universal reference for comparing thermochemical data across different reactions and sources. The symbol ΔH° (read "delta H naught" or "delta H standard") denotes a standard enthalpy change.
Key Formula
ΔH° = ΔU° + Δng·RT
LaTeX: \Delta H^\circ = \Delta U^\circ + \Delta n_g RT
| Symbol | Meaning | Unit |
|---|---|---|
| ΔH° | Standard enthalpy change | kJ/mol |
| ΔU° | Standard internal energy change | kJ/mol |
| Δng | Change in moles of gas | mol |
| R | Universal gas constant (8.314 J/mol·K) | J/(mol·K) |
| T | Temperature | K |
Worked Example
Problem
For the reaction N₂(g) + 3H₂(g) → 2NH₃(g), ΔU° = −85.0 kJ/mol at 298 K. Calculate ΔH°.
Solution
Step 1: Determine Δng = moles of gaseous products − moles of gaseous reactants. Δng = 2 − (1 + 3) = 2 − 4 = −2 mol. Step 2: Apply the formula: ΔH° = ΔU° + Δng·RT. ΔH° = −85.0 + (−2 × 8.314 × 10⁻³ × 298) Step 3: Calculate the correction term: −2 × 0.008314 × 298 = −4.955 kJ/mol. Step 4: ΔH° = −85.0 + (−4.955) = −89.955 ≈ −90.0 kJ/mol.
Answer
ΔH° ≈ −90.0 kJ/mol
Common Standard Enthalpy Values for Key Reactions
| Reaction Type | Example | ΔH° (kJ/mol) | Exo/Endo |
|---|---|---|---|
| Combustion | CH₄ + 2O₂ → CO₂ + 2H₂O | −890.3 | Exothermic |
| Neutralization | HCl(aq) + NaOH(aq) → NaCl(aq) + H₂O | −57.3 | Exothermic |
| Vaporization | H₂O(l) → H₂O(g) | +44.0 | Endothermic |
| Fusion (melting) | H₂O(s) → H₂O(l) | +6.01 | Endothermic |
| Synthesis of NH₃ | N₂ + 3H₂ → 2NH₃ | −92.4 | Exothermic |
| Decomposition of CaCO₃ | CaCO₃ → CaO + CO₂ | +178.0 | Endothermic |
Interactive Tools
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Related Terms
Hess's Law
Hess's Law states that the total enthalpy change of a chemical reaction is independent of the pathway taken, depending only on the initial and final states of the system. This principle allows chemists to calculate enthalpy changes for reactions that are difficult or impossible to measure directly by combining known thermochemical equations. It is a direct consequence of the First Law of Thermodynamics, making enthalpy a state function.
Enthalpy of Formation
The enthalpy of formation (ΔH°f) is the heat change when one mole of a compound is formed from its constituent elements in their standard states at 298 K and 1 atm pressure. Standard enthalpies of formation for elements in their most stable form are defined as zero. These values are tabulated and used extensively with Hess's Law to calculate enthalpy changes for any chemical reaction.
Gibbs Free Energy
Gibbs free energy (G) is a thermodynamic potential that measures the maximum reversible work a system can perform at constant temperature and pressure, and determines the spontaneity of a process. A negative ΔG indicates a spontaneous reaction, a positive ΔG indicates a non-spontaneous reaction, and ΔG = 0 indicates a system at equilibrium. It combines enthalpy and entropy into a single criterion for spontaneity, making it one of the most powerful tools in chemical thermodynamics.
From Greek 'enthalpein' meaning 'to heat within' (en- 'in' + thalpein 'to heat'). The term was introduced by Heike Kamerlingh Onnes around 1909. The degree symbol (°) denotes standard-state conditions.