GATE Life Sciences Thermodynamics Syllabus

Thermodynamics syllabus — full chapter and topic list

Expand any chapter to see its topics and sub-topics. This is the whole examinable outline for Thermodynamics in GATE Life Sciences, not a summary of it.

1. 1

   Qualitative treatment of state and path functions

   overview

   Examined as a single unit within Thermodynamics — no further topic split in the official outline.

2. 2

   First law

   7 topics

   • Reversible and irreversible processes
   • Internal energy
   • Enthalpy
   • Kirchoff equation
   • Heat of reaction
   • Hess's law
   • Heat of formation
3. 3

   Second law

   2 topics

   • Entropy
   • Free energy
4. 4

   Gibbs-Helmholtz equation

   overview

   Examined as a single unit within Thermodynamics — no further topic split in the official outline.

5. 5

   Free energy change and spontaneity

   overview

   Examined as a single unit within Thermodynamics — no further topic split in the official outline.

6. 6

   Free energy changes from equilibrium constant

   overview

   Examined as a single unit within Thermodynamics — no further topic split in the official outline.

Thermodynamics flashcards for GATE Life Sciences

18 of 50 cards from the Thermodynamics deck — real questions with worked answers.

1. Define a reversible process in thermodynamics.

   A process carried out infinitesimally slowly through a continuous series of equilibrium states, such that the system and surroundings can be restored to their initial states by an infinitesimal reversal of conditions. The driving force exceeds the opposing force by an infinitesimal amount.

2. Define an irreversible process and give a common example.

   A process that occurs at a finite rate with a finite difference between driving and opposing forces, passing through non-equilibrium states; the system cannot be returned to its initial state without leaving permanent changes in the surroundings. Example: free expansion of a gas into vacuum, or heat flow across a finite temperature difference.

3. Which gives more work: reversible or irreversible expansion of a gas between the same two states?

   The reversible expansion delivers the maximum possible work, $w_{rev} > w_{irrev}$ (in magnitude), because at every step the system pushes against the largest possible opposing pressure.

4. Compare reversible and irreversible processes with respect to total entropy change of the universe.

   For a reversible process $\Delta S_{universe} = 0$; for an irreversible (spontaneous) process $\Delta S_{universe} > 0$. No real process gives $\Delta S_{universe} < 0$.

5. Define internal energy ($U$) of a system.

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

6. State the first law of thermodynamics in terms of internal energy.

   $$\Delta U = q + w$$ where $q$ is heat absorbed by the system and $w$ is work done on the system (IUPAC sign convention). Energy can be neither created nor destroyed.

7. For a process at constant volume, how is the heat exchanged related to internal energy?

   At constant volume no pressure-volume work is done ($w = 0$), so $q_{V} = \Delta U$. The heat measured in a bomb calorimeter equals $\Delta U$.

8. Why is internal energy a state function?

   Because its value depends only on the present state of the system (defined by state variables like $T$, $P$, $V$) and not on the path taken to reach that state; hence $\oint dU = 0$ over a cycle.

9. Define enthalpy ($H$) and give its defining equation.

   Enthalpy is the heat content of a system at constant pressure, defined as $$H = U + PV.$$ It is a state function and an extensive property.

10. How is enthalpy change related to heat at constant pressure?

    At constant pressure, $q_{P} = \Delta H$. The heat absorbed or released by a reaction carried out at constant pressure equals its enthalpy change.

11. Derive the relation between $\Delta H$ and $\Delta U$ for a process at constant pressure.

    From $H = U + PV$, at constant $P$: $$\Delta H = \Delta U + P\Delta V.$$ For ideal gases this becomes $\Delta H = \Delta U + \Delta n_{g}RT$.

12. What is $\Delta n_{g}$ in the equation $\Delta H = \Delta U + \Delta n_{g}RT$?

    $\Delta n_{g}$ is the change in the number of moles of gas in the reaction: $\Delta n_{g} = n_{products(g)} - n_{reactants(g)}$. Only gaseous species are counted.

13. For the reaction $\ce{N2(g) + 3H2(g) -> 2NH3(g)}$, what is $\Delta n_{g}$ and the relation between $\Delta H$ and $\Delta U$?

    $\Delta n_{g} = 2 - (1+3) = -2$, so $\Delta H = \Delta U - 2RT$. Here $\Delta H < \Delta U$.

14. When are $\Delta H$ and $\Delta U$ equal for a reaction?

    When $\Delta n_{g} = 0$ (no change in moles of gas), e.g. $\ce{H2(g) + I2(g) -> 2HI(g)}$, or for reactions involving only solids and liquids where $P\Delta V \approx 0$.

15. Define an exothermic reaction in terms of $\Delta H$.

    A reaction that releases heat to the surroundings, having $\Delta H < 0$ (negative). The products have lower enthalpy than the reactants.

16. Define an endothermic reaction in terms of $\Delta H$.

    A reaction that absorbs heat from the surroundings, having $\Delta H > 0$ (positive). The products have higher enthalpy than the reactants.

17. State Kirchhoff's equation (differential form) for the temperature dependence of reaction enthalpy.

    $$\left(\frac{\partial (\Delta H)}{\partial T}\right)_{P} = \Delta C_{P}$$ where $\Delta C_{P}$ is the difference in heat capacities at constant pressure between products and reactants.

18. State Kirchhoff's equation in integrated form for reaction enthalpy between two temperatures.

    $$\Delta H_{T_{2}} = \Delta H_{T_{1}} + \Delta C_{P}\,(T_{2} - T_{1})$$ assuming $\Delta C_{P}$ is constant over the temperature range.

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