The negative value indicates an exothermic reaction (heat is released).
🌡️ Entropy Change in a Reversible Process
Problem: A system absorbs 6000 J of heat reversibly at 300 K. What is the entropy change?
Solution: Using ΔS = q_rev / T
ΔS = 6000 / 300 = 20 J/K
The positive entropy change indicates increased disorder in the system.
⚡ Gibbs Free Energy and Spontaneity
Problem: For a reaction, ΔH = −100 kJ and ΔS = 0.2 kJ/K at 298 K. Is the reaction spontaneous?
Solution: Using ΔG = ΔH − TΔS
ΔG = −100 − (298 × 0.2) = −100 − 59.6
ΔG = −159.6 kJ
Since ΔG is negative, the reaction is spontaneous at 298 K.
Thermochemistry Formula & Guide
q = m × c × ΔT
Enthalpy change (heat transfer)
Where q is heat energy, m is mass, c is specific heat capacity, and ΔT is the temperature change. A positive q indicates heat absorbed (endothermic), negative q indicates heat released (exothermic).
ΔH°rxn = ΣΔH°f(products) − ΣΔH°f(reactants)
Standard enthalpy of reaction
Where ΔH°f is the standard enthalpy of formation. This is Hess's law applied to standard formation enthalpies.
ΔS = qrev / T
Entropy change in a reversible process
Where ΔS is the entropy change, q_rev is the reversible heat transfer, and T is the absolute temperature in Kelvin.
ΔG = ΔH − TΔS
Gibbs free energy change
Where ΔG is Gibbs free energy change, ΔH is enthalpy change, T is absolute temperature, and ΔS is entropy change. A negative ΔG indicates a spontaneous process.
Key Concepts
📌 Spontaneity Rules
ΔG < 0: Spontaneous (exergonic). ΔG > 0: Non-spontaneous (endergonic). ΔG = 0: Equilibrium. A reaction can be spontaneous even with positive ΔH if ΔS is large enough and T is high.
📌 Endothermic vs Exothermic
Endothermic (ΔH > 0): System absorbs heat from surroundings. Exothermic (ΔH < 0): System releases heat to surroundings. Most combustion reactions are exothermic.
📌 Temperature Units
Always use Kelvin (K) for temperature in entropy and Gibbs free energy calculations. For enthalpy (q = mcΔT), the temperature change ΔT is the same in °C and K since the intervals are equal.
📌 Standard Conditions
Standard thermochemical conditions are 298.15 K (25°C) and 1 bar pressure. Standard enthalpies of formation (ΔH°f) are defined for forming 1 mole of a compound from its elements in their standard states.
🔥
Enthalpy Mode
Calculate heat transfer using q = mcΔT. Supports mass in grams and kilograms, specific heat in various units, and temperature in °C or K.
⚗️
Reaction Enthalpy Mode
Calculate standard reaction enthalpy ΔH°rxn from formation enthalpies using Hess's law. Enter sums of product and reactant formation enthalpies.
🌡️
Entropy Mode
Calculate entropy change using ΔS = q_rev / T. Supports energy units in J, kJ, and kcal with automatic temperature conversion to Kelvin.
⚡
Gibbs Free Energy Mode
Determine reaction spontaneity with ΔG = ΔH − TΔS. Automatic unit conversion and spontaneity classification (spontaneous/non-spontaneous/equilibrium).
⚠️ Important Note: Thermochemistry calculations assume ideal conditions and standard states unless otherwise specified. Real-world reactions may deviate due to non-ideal behavior, variable pressure conditions, side reactions, or incomplete conversion. Always verify results with experimental data for critical applications. For reactions involving gases, assume ideal gas behavior unless stated otherwise.
Frequently Asked Questions
What is enthalpy (ΔH) in thermochemistry?
Enthalpy (H) is a thermodynamic quantity that represents the total heat content of a system at constant pressure. The change in enthalpy ΔH measures the heat absorbed or released during a process. A positive ΔH means the system absorbs heat (endothermic), while a negative ΔH means the system releases heat (exothermic). Enthalpy change is calculated as q = mcΔT and is a state function, meaning it depends only on the initial and final states, not the path taken.
What is the difference between ΔH and ΔH°rxn?
ΔH is the general enthalpy change for any process or reaction under any conditions. ΔH°rxn (standard reaction enthalpy) specifically refers to the enthalpy change when a reaction occurs under standard conditions (298.15 K, 1 bar pressure, and all reactants and products in their standard states). The standard reaction enthalpy is calculated from standard enthalpies of formation using ΔH°rxn = ΣΔH°f(products) − ΣΔH°f(reactants).
What does Gibbs free energy (ΔG) tell us about a reaction?
Gibbs free energy (ΔG) determines whether a reaction is spontaneous (occurs without external intervention) under constant temperature and pressure. The sign of ΔG is the key indicator: ΔG < 0 means the reaction is spontaneous (exergonic), ΔG > 0 means non-spontaneous (endergonic), and ΔG = 0 means the system is at equilibrium. The Gibbs free energy combines enthalpy and entropy into a single spontaneity criterion: ΔG = ΔH − TΔS.
How do temperature and entropy affect spontaneity?
The equation ΔG = ΔH − TΔS shows that temperature (T) and entropy change (ΔS) together influence spontaneity. Even a reaction with positive ΔH (endothermic) can be spontaneous at high temperatures if the entropy change is positive (increased disorder). Conversely, a reaction with negative ΔH and negative ΔS may only be spontaneous at low temperatures. The crossover temperature where ΔG = 0 is given by T = ΔH/ΔS — above this temperature, the reaction becomes spontaneous (if ΔS > 0) or non-spontaneous (if ΔS < 0).
What is entropy (ΔS) and why is it important?
Entropy (S) is a measure of the disorder or randomness in a system. The change in entropy ΔS during a reversible process is calculated as ΔS = q_rev / T. According to the Second Law of Thermodynamics, the total entropy of an isolated system always increases over time. A positive ΔS indicates increased disorder (e.g., melting ice, dissolving salt), while a negative ΔS indicates decreased disorder (e.g., freezing water, gas condensation). Entropy plays a crucial role in determining reaction spontaneity through the Gibbs free energy equation.
What are standard conditions in thermochemistry?
Standard conditions in thermochemistry are defined as 298.15 K (25°C) and 1 bar (100 kPa) pressure. For substances in solution, the standard concentration is 1 M (molar). The standard state of an element is its most stable form at these conditions. Standard enthalpies of formation ΔH°f are defined as the enthalpy change when 1 mole of a compound is formed from its elements in their standard states. The superscript ° (degree symbol) always denotes standard conditions in thermochemical equations.