Aim to answer every question before checking. Missed questions point you to the sections you need most.
1. (7.1) In thermodynamics, the "system" is:
B. Defining the system is the first step. System + surroundings = universe.
2. (7.1) An isolated system:
C. Open = exchanges both; closed = exchanges only energy; isolated = exchanges neither. A thermos is a decent approximation of an isolated system.
3. (7.2) State functions are characterized by:
D. State functions are written with capital letters (H, S, U, G). Their values are like elevation: only the start and end points matter.
4. (7.2) Work (w) and heat (q) are:
A. The amount of heat or work depends on how the change is carried out, so neither is a state function. Only the sum q + w (= ΔU) is.
5. (7.3) Enthalpy (H) is convenient for constant-pressure processes because:
B. Most lab and biological reactions happen at constant P, making ΔH a direct reading of heat flow.
6. (7.3) An exothermic reaction:
C. Combustion is exothermic; melting is endothermic. Sign conventions: negative ΔH = exothermic.
7. (7.4) Calorimetry relates heat exchanged to measurable quantities via:
D. Coffee-cup calorimeters assume constant pressure and that the heat absorbed by the solution equals -q of the reaction.
8. (7.4) A bomb calorimeter measures:
A. The rigid sealed chamber means no pressure-volume work, so all heat flow appears as internal-energy change.
9. (7.5) Hess's law states that:
B. A direct consequence of H being a state function. Great for computing ΔH of reactions that are hard to run directly.
10. (7.5) When applying Hess's law we can:
C. Treat chemical equations algebraically: reverse to flip the sign, scale to scale ΔH.
11. (7.6) The standard enthalpy of formation ΔH°_f is:
D. ΔH°_rxn = Σ ΔH°_f(products) - Σ ΔH°_f(reactants). Tabulated values let you compute almost any enthalpy change.
12. (7.6) The ΔH°_f of an element in its standard state is:
A. This sets a reference point. E.g., O₂(g), C(s, graphite), and Na(s) all have ΔH°_f = 0.
13. (7.7) Bond dissociation energy is:
B. Always positive; breaking bonds requires energy. Typical C-C is about 347 kJ/mol.
14. (7.7) For a reaction, ΔH can be estimated by:
C. Same idea as A, expressed operationally. Signs matter: breaking costs energy, forming releases it.
15. (7.8) Entropy (S) is best described as:
D. Boltzmann: S = k·ln(W). More microstates = higher entropy. Gases have more entropy than liquids than solids.
16. (7.8) The second law of thermodynamics states that for an isolated system:
A. Entropy of the universe tends to increase. Local decreases (a cell ordering itself) require an entropy increase somewhere else.
17. (7.9) The Gibbs free energy change is defined as:
B. At constant T and P, ΔG captures both enthalpic and entropic drivers of spontaneity.
18. (7.9) A reaction is spontaneous at constant T and P when:
C. Spontaneous means the reaction can occur, not that it does occur quickly. Kinetics controls rate.
19. (7.10) The relationship between ΔG° and Keq is:
D. Large K → very negative ΔG°; K = 1 → ΔG° = 0.
20. (7.10) If ΔG° < 0:
A. Negative ΔG° ↔ K > 1. The farther ΔG° sits below zero, the more products are favored.
21. (7.11) A coupled reaction is one in which:
B. Life depends on coupling. ATP hydrolysis (-30.5 kJ/mol under cellular conditions) pushes endergonic biosynthesis forward.
22. (7.11) The overall ΔG of a coupled reaction equals:
C. Just sum the ΔG values. If the sum is negative, the coupled process is spontaneous overall.
23. (7.12) The enthalpy of fusion (Δ
Hfus) is:
D. Positive (endothermic) by convention. Δ
Hvap > Δ
Hfus for most substances because more intermolecular forces are broken when going from liquid to gas.
24. (7.12) Melting (solid → liquid):
A. Heat input breaks intermolecular forces enough to let molecules slide past one another.
