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Gen Chem (8)
| Question | Answer |
|---|---|
| What does the First Law of Thermodynamics state? | Energy cannot be created or destroyed; it can only be converted from one form to another. The total amount of energy remains constant. |
| Define system, surroundings, and universe in thermodynamics. | System = the chemical reaction (reactants and products). Surroundings = the environment around the system. Universe = system + surroundings. |
| What is enthalpy (H)? | The amount of heat energy contained within a system. |
| What is an endothermic process? What is the sign of ΔH? | Heat is transferred from the surroundings into the system. ΔH > 0 (positive). |
| What is an exothermic process? What is the sign of ΔH? | Heat is transferred from the system to the surroundings. ΔH < 0 (negative). |
| List the 3 endothermic phase changes. | Melting (solid → liquid), Vaporization (liquid → gas), Sublimation (solid → gas). |
| List the 3 exothermic phase changes. | Freezing (liquid → solid), Condensation (gas → liquid), Deposition (gas → solid). |
| In a reaction coordinate diagram, how do you distinguish endothermic from exothermic? | Endothermic: products are higher in energy than reactants. Exothermic: products are lower in energy than reactants. |
| What are the 3 types of heat transfer? | Conduction (direct contact), Convection (currents in liquids/gases; hot/less dense rises, cool/more dense sinks), Radiation (electromagnetic radiation). |
| When is work positive vs. negative? | Positive work: volume decreases (ΔV < 0), compression, work done onto the system. Negative work: volume increases (ΔV > 0), expansion, work done by the system. |
| Summarize the heat/work sign conventions. | +q = endothermic, −q = exothermic, +w = compression, −w = expansion. |
| What is specific heat capacity (C) and what are its units? | The amount of energy required to raise 1.0 gram of a substance by 1°C. Units: J·g⁻¹·°C⁻¹. |
| What happens to temperature during a phase change on a phase change diagram? | Temperature remains constant during the phase change. Heat energy goes toward changing the phase, not the temperature. |
| What is a bomb calorimeter? | A closed vessel used to measure heat emitted by a combustion reaction burned under an oxygen-rich atmosphere surrounded by water. |
| What are standard conditions and what symbol represents them? | 298 K, 1 atm, 1.0 M concentration. Represented by the symbol °. |
| What are the 3 methods to calculate standard enthalpy change (ΔH°)? | Bond enthalpies, enthalpies of formation, and Hess's Law. |
| Is breaking bonds endothermic or exothermic? What about forming bonds? | Breaking bonds = endothermic (ΔH > 0). Forming bonds = exothermic (ΔH < 0). |
| What is the standard enthalpy of formation for any element in its standard state? | Zero (0 kJ·mol⁻¹). |
| What are the 7 diatomic molecules? Give the mnemonic. | H₂, N₂, F₂, O₂, I₂, Cl₂, Br₂. Mnemonic: "Have No Fear Of Ice Cold Beer." |
| Which two elements are liquids at standard state? | Mercury (Hg) and Bromine (Br₂). |
| What does Hess's Law state? | The enthalpy change of a reaction is independent of the pathway taken. Total ΔH equals the sum of all intermediate ΔH values. |
| How do you manipulate ΔH when applying Hess's Law? | Reverse a reaction → multiply ΔH by −1. Multiply equation by a coefficient → multiply ΔH by the same coefficient. Divide equation → divide ΔH by the same coefficient. |
| What is entropy (S)? | A measure of the disorder or randomness of a system. All substances have some entropy; larger, more complex molecules have more entropy. |
| What does the 2nd Law of Thermodynamics state? | The entropy of the universe is always increasing. All spontaneous processes produce an increase in the entropy of the universe. |
| What does the 3rd Law of Thermodynamics state? | The entropy of a pure crystalline substance at absolute zero (0 K) is zero. |
| What do ΔS(rxn) > 0 and ΔS(rxn) < 0 indicate? | ΔS > 0: entropically favorable; products are more disordered. ΔS < 0: entropically unfavorable; products are more ordered. |
| List the 4 general entropy trends. | 1) Solid → liquid → gas increases entropy. 2) Dissolution increases entropy. 3) Higher temperature increases entropy. 4) More product gas molecules than reactant gas molecules increases entropy. |
| What is a spontaneous reaction? | Any process that occurs without the input of external energy. |
| What is Gibbs Free Energy (ΔG)? | The criterion for spontaneity of a reaction at constant temperature and pressure. It depends on enthalpy, entropy, and temperature. |
| What do the signs of ΔG indicate? | ΔG < 0 = spontaneous (exergonic). ΔG > 0 = non-spontaneous (endergonic). ΔG = 0 = equilibrium. |
| What is the difference between exergonic and endergonic reactions? | Exergonic: ΔG < 0, energy released, bonds formed are stronger than bonds broken. Endergonic: ΔG > 0, energy absorbed, bonds formed are weaker than bonds broken. |
| −ΔH, +ΔS → | spontaneous at all temperatures |
| −ΔH, −ΔS → | spontaneous at low T, non-spontaneous at high T |
| +ΔH, +ΔS → | non-spontaneous at low T, spontaneous at high T |
| +ΔH, −ΔS → | non-spontaneous at all temperatures |
| What is the relationship between ΔG° and the equilibrium constant (Keq)? | ΔG° < 0 → Keq > 1, products favored. ΔG° > 0 → Keq < 1, reactants favored. ΔG° = 0 → Keq = 1, both equally favored. |
| How do you find the temperature at which a reaction transitions between spontaneous and non-spontaneous? | Set ΔG = 0 and solve for T using ΔG = ΔH − TΔS. Make sure ΔS is converted to kJ·K⁻¹ to match units with ΔH. |
Created by:
smurtab