Chemistry
25 flashcards
~13 min
Thermodynamics and Thermochemistry
Understand the energy changes that drive chemical reactions with this comprehensive thermodynamics and thermochemistry flashcard deck. Thermodynamics answers one of chemistry's most fundamental questions: why do reactions occur spontaneously in one directi...
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Understand the energy changes that drive chemical reactions with this comprehensive thermodynamics and thermochemistry flashcard deck. Thermodynamics answers one of chemistry's most fundamental questions: why do reactions occur spontaneously in one direction? This deck covers the laws of thermodynamics, heat and work, enthalpy (ΔH), entropy (ΔS), Gibbs free energy (ΔG), Hess's law, standard enthalpies of formation, calorimetry calculations, specific heat capacity, and bond enthalpy calculations. Essential for AP Chemistry, A-Level Chemistry, IB Chemistry, and university general and physical chemistry courses.
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The study of energy and its transformations, particularly the relationship between heat, work, and chemical reactions. It determines whether a reaction is spontaneous.
Energy cannot be created or destroyed — only converted from one form to another. ΔU = q + w (change in internal energy = heat absorbed + work done on system). Conservation of energy.
The total entropy of an isolated system always increases in a spontaneous process. The universe tends toward disorder. No process can be 100% energy-efficient.
The entropy of a perfect crystalline substance is zero at absolute zero (0 K / −273.15°C). This sets the absolute reference point for entropy measurements.
A measure of the total heat content of a system at constant pressure. ΔH = heat absorbed or released in a reaction. ΔH < 0 = exothermic. ΔH > 0 = endothermic.
The enthalpy change when 1 mole of a compound is formed from its elements in their standard states at 298 K and 1 atm. ΔHf° of any pure element in its standard state = 0.
The total enthalpy change of a reaction is the same regardless of the route taken — it depends only on initial and final states. Allows calculation of ΔH for reactions that cannot be measured directly.
Add the enthalpy changes of individual steps that sum to the target reaction. If a step is reversed, change the sign of ΔH. If multiplied, multiply ΔH by the same factor.
A measure of the disorder or randomness of a system. ΔS > 0 = increase in disorder. ΔS < 0 = decrease in disorder. Gases have higher entropy than liquids, which have higher than solids.
1. Phase change from solid → liquid → gas. 2. Dissolving a solid in solution. 3. Increasing temperature. 4. Increasing moles of gas in a reaction. 5. Mixing different substances.
The energy available to do useful work. Determines spontaneity: ΔG = ΔH − TΔS. If ΔG < 0 → spontaneous. ΔG > 0 → non-spontaneous. ΔG = 0 → equilibrium.
ΔH− ΔS+: Always spontaneous (ΔG always negative). ΔH+ ΔS−: Never spontaneous (ΔG always positive). ΔH− ΔS−: Spontaneous at low temperature. ΔH+ ΔS+: Spontaneous at high temperature.
The amount of heat needed to raise the temperature of 1 gram of a substance by 1°C (or 1 K). Water has an unusually high specific heat: 4.18 J/g·°C.
q = mcΔT. q = heat energy (J), m = mass (g), c = specific heat capacity (J/g·°C), ΔT = temperature change (°C or K). Used in all calorimetry calculations.
Experimental measurement of heat flow in chemical reactions. A calorimeter measures temperature changes to calculate heat absorbed or released. Assumes no heat loss to surroundings.
A constant-volume calorimeter used to measure the heat of combustion reactions. The reaction occurs in a sealed steel container submerged in water. Measures ΔU (internal energy) not ΔH.
The energy required to break one mole of a specific bond in gaseous molecules. Breaking bonds = endothermic (absorbs energy). Forming bonds = exothermic (releases energy).
ΔH = Σ(bonds broken) − Σ(bonds formed). Sum the energies of bonds broken in reactants (positive) and subtract the sum of bonds formed in products (negative).
The energy required to separate one mole of an ionic solid into its gaseous ions. Large lattice enthalpy = strong ionic lattice = high melting point. Related to ion charge and size.
A thermochemical cycle applying Hess's Law to calculate lattice enthalpy of an ionic compound. Uses enthalpy of formation, atomization, ionization energies, and electron affinity values.
The enthalpy change when 1 mole of a substance is completely burned in excess oxygen under standard conditions. Always exothermic (negative ΔH). Example: CH₄: ΔHc° = −890 kJ/mol.
The heat released when an acid and base react to form 1 mole of water. For strong acid + strong base it is approximately −57 kJ/mol regardless of the specific acid/base used.
Exothermic: Products at lower energy than reactants. ΔH = negative. Energy is released as heat. Endothermic: Products at higher energy than reactants. ΔH = positive. Energy is absorbed as heat.
The minimum amount of energy required for a reaction to occur. A catalyst lowers the activation energy — increasing reaction rate without being consumed.
ΔG° = −RT ln K. If K > 1 → ΔG° < 0 → products favored. If K < 1 → ΔG° > 0 → reactants favored. If K = 1 → ΔG° = 0 → equilibrium.