Thermodynamics & Thermochemistry flash cards

1. 1.Define a thermodynamic system and its surroundings.

   The system is the specific part of the universe under study (the reacting mixture). The surroundings is everything else in the universe outside the system, from which it may exchange matter and/or energy. Together they make up the universe.

   Hint: Part under study vs everything else.

2. 2.Distinguish open, closed and isolated systems.

   Open: exchanges both matter and energy with surroundings (open beaker). Closed: exchanges only energy, not matter (sealed flask). Isolated: exchanges neither matter nor energy (thermos flask, ideal).

   Hint: What can cross the boundary — matter, energy, both, neither?

3. 3.What is the difference between an extensive and an intensive property? Give examples.

   Extensive properties depend on the amount of matter: mass, volume, internal energy $U$, enthalpy $H$, entropy $S$, Gibbs energy $G$. Intensive properties are independent of amount: temperature, pressure, density, molar volume, refractive index.

   Hint: Does halving the sample halve it?

4. 4.What is a state function? Name several.

   A property whose value depends only on the present state of the system, not on the path taken to reach it. Examples: $U$, $H$, $S$, $G$, $T$, $P$, $V$. Their change depends only on initial and final states.

   Hint: Path-independent; depends only on state.

5. 5.Are heat $q$ and work $w$ state functions? Why?

   No. $q$ and $w$ are path functions — their magnitudes depend on the route taken between two states, not just on the endpoints. Only their sum, $\Delta U = q + w$, is path-independent (a state function).

   Hint: Same endpoints, different route, different values.

6. 6.State the sign conventions (IUPAC) for heat and work in $\Delta U = q + w$.

   Heat absorbed by the system: $q > 0$; heat released: $q < 0$. Work done on the system: $w > 0$; work done by the system (expansion): $w < 0$. Energy entering the system is positive.

   Hint: Energy into system = positive.

7. 7.Define internal energy $U$.

   The total energy stored within a system — the sum of all kinetic and potential energies of its molecules (translational, rotational, vibrational, electronic, nuclear, intermolecular). It is a state function; only changes $\Delta U$ are measurable, not absolute values.

   Hint: Total microscopic energy content.

8. 8.State the first law of thermodynamics and its equation.

   Energy can be neither created nor destroyed (conservation of energy). For a closed system: $\Delta U = q + w$, where the change in internal energy equals heat added plus work done on the system. For an isolated system $\Delta U = 0$.

   Hint: Conservation of energy for a system.

9. 9.Derive the expression for pressure–volume work during expansion.

   For expansion against external pressure $P_{ext}$, $w = -P_{ext}\,\Delta V = -P_{ext}(V_2 - V_1)$. The negative sign: on expansion $\Delta V>0$, the system does work, so $w<0$. This is the work done on the system.

   Hint: Force over area times distance → $-P_{ext}\Delta V$.

10. 10.Why is $\Delta U = q_V$ (heat at constant volume)?

    At constant volume $\Delta V = 0$, so PV-work $w = -P\Delta V = 0$. Then $\Delta U = q_V + 0 = q_V$. Thus the heat measured in a bomb calorimeter (rigid, constant $V$) equals $\Delta U$.

    Hint: No volume change → no work → all heat is $\Delta U$.

11. 11.Define enthalpy $H$ and show why $\Delta H = q_P$.

    $H = U + PV$. At constant pressure, $\Delta H = \Delta U + P\Delta V = (q_P - P\Delta V) + P\Delta V = q_P$. So enthalpy change equals heat absorbed at constant pressure — the usual lab condition.

    Hint: $H=U+PV$; constant-P heat.

12. 12.Give the relation between $\Delta H$ and $\Delta U$ for a reaction involving gases.

    $\Delta H = \Delta U + \Delta n_g RT$, where $\Delta n_g$ = (moles of gaseous products − moles of gaseous reactants). Derived from $\Delta(PV) = \Delta n_g RT$ for ideal gases at constant $T$.

    Hint: $\Delta n_g$ counts gas moles only.

13. 13.For which reactions is $\Delta H = \Delta U$?

    When $\Delta n_g = 0$ (no change in moles of gas), since $\Delta H = \Delta U + \Delta n_g RT$. Example: $H_2(g) + Cl_2(g) \to 2HCl(g)$, $\Delta n_g = 2-2 = 0$. Also true for reactions with only solids/liquids.

    Hint: Equal gas moles on both sides.

14. 14.Define heat capacity $C$, and molar and specific heat capacities.

    Heat capacity $C = q/\Delta T$ is the heat needed to raise a substance's temperature by 1 K. Molar heat capacity is per mole ($J\,K^{-1}\,mol^{-1}$); specific heat is per gram ($J\,K^{-1}\,g^{-1}$).

    Hint: Heat per degree; per mole or per gram.

15. 15.Define $C_V$ and $C_P$ in terms of $U$ and $H$.

    $C_V = \left(\dfrac{\partial U}{\partial T}\right)_V$ (constant volume) and $C_P = \left(\dfrac{\partial H}{\partial T}\right)_P$ (constant pressure). Thus $\Delta U = C_V\Delta T$ and $\Delta H = C_P\Delta T$ for given $n$ moles.

    Hint: Slopes of $U$–$T$ and $H$–$T$.

16. 16.State Mayer's relation for an ideal gas and why $C_P > C_V$.

    $C_P - C_V = R$ (per mole). $C_P > C_V$ because at constant pressure part of the supplied heat does expansion work, so more heat is needed for the same temperature rise than at constant volume.

    Hint: Extra heat goes into expansion work.

17. 17.Give $C_V$, $C_P$ and $\gamma$ for monatomic and diatomic ideal gases.

    Monatomic: $C_V = \tfrac{3}{2}R$, $C_P = \tfrac{5}{2}R$, $\gamma = 5/3 \approx 1.67$. Diatomic: $C_V = \tfrac{5}{2}R$, $C_P = \tfrac{7}{2}R$, $\gamma = 7/5 = 1.40$. Here $\gamma = C_P/C_V$.

    Hint: Degrees of freedom: 3 vs 5.

18. 18.Distinguish a reversible from an irreversible process.

    A reversible process proceeds infinitesimally slowly through a continuous series of equilibrium states; driving force is infinitesimal and it can be exactly reversed. An irreversible (real, spontaneous) process occurs at finite rate with a finite driving force and cannot be exactly reversed.

    Hint: Quasi-static equilibrium vs real finite-rate.

19. 19.Why does a reversible isothermal expansion do maximum work?

    At every step $P_{ext}$ is only infinitesimally less than the gas pressure, so the opposing pressure is as large as possible throughout. Work done by the gas $= \int P\,dV$ is therefore maximal. Any irreversible expansion against lower $P_{ext}$ does less work.

    Hint: Opposing pressure kept as high as possible.

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