| Term | Definition |
|---|
| First Law of Thermodynamics | Energy is constant - can be neither created nor destroyed (only transferred / converted) |
| Second Law of Thermodynamics | If a process is spontaneous in one direction, it cannot be in the opposite direction |
| Standard State Conditions | For ΔH°, ΔS°, ΔG°: gases @ 1 atm; liquids & solids are pure; solutions are 1M; energy of formation for an element is ZERO; temperature used is 28°C (298K) |
| Enthalpy | ΔH: when bonds formed, energy released; bonds broken, absorbed |
| ΔH | ΔH = H(products) - H(reactants); if negative, rxn is exothermic; positive, endothermic |
| ΔH°f | Δ energy when 1 mol compound formed from pure element; negative value = more stable |
| ΔH° | ΔH°= ΣΔH°f (products) - ΣΔH°f (reactants) |
| Bond Energy | Energy required to break a bond (endothermic); ΔH° = Σ (bond energies of bonds broken) - Σ (bond energies of bonds formed) |
| Hess's Law | If a rxn is multistep, ΔH for overall rxn is sum of ΔH for all steps |
| Heat Capacity (Cp) | measure of how much temperature of an object is raised when it absorbs heat |
| Cp = | Cp = ΔH / ΔT; small cp = large ↑ in temp w/ increase in energy |
| Specific Heat (c) | Amount of heat required to raise one gram substance by 1°C; q = mcΔT |
| Entropy (S) | Measure of disorder; solid crystal = ZERO; solid < liquid < gas; solid < particles in sln; 1 mol < 2 mol |
| ΔS° | ΔS° = ΣΔS° (products) - ΣΔS° (reactants) |
| Gibbs Free Energy (G) | Measure of spontaneity; negative = spontaneous; if zero = @ equilibrium |
| ΔG° | ΔG° = ΣΔG° (products) - ΣΔG° (reactants) |
| Relationship Between ΔG, ΔS, and ΔH | ΔG° = ΔH - TΔS° [high temp, entropy dominant; low temp, enthalpy dominant] |
| ΔG and ΔG° | Used when sln not @ 1M: ΔG = ΔG° + RT ln (Q) |
| ΔG° and k (equilibrium constant) | ΔG° = -RT ln (k) [IF ΔG° negative THEN products favored; IF ΔG° positive THEN reactants favored] |
Combine with other sets
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