JEE Advanced Physical Chemistry Flashcards

24 sample cards from the Physical Chemistry deck

Sampled from the end of the deck, so these are different cards from the ones shown on the syllabus page.

1. Write the equation for osmotic pressure of a dilute solution.

   π = (n/V)RT = CRT, where C is molar concentration. For electrolytes π = iCRT.

2. State the first law of thermodynamics and its mathematical expression.

   Energy can neither be created nor destroyed; the internal energy change of a system equals heat added minus work done by the system: ΔU = q + w (w = work done ON system).

3. Define enthalpy (H) and relate ΔH to ΔU.

   H = U + PV (heat content at constant pressure). ΔH = ΔU + Δ(PV); for ideal gases ΔH = ΔU + Δn_g·RT, where Δn_g is change in moles of gas.

4. State Hess's law of constant heat summation.

   The total enthalpy change of a reaction is the same whether it occurs in one step or several steps, depending only on initial and final states (consequence of H being a state function).

5. Define entropy and give the sign convention for a spontaneous process in an isolated system.

   Entropy (S) is a measure of disorder/randomness. For a spontaneous process in an isolated system, total entropy increases: ΔS_total > 0 (second law of thermodynamics).

6. Write the Gibbs free energy equation and the criterion for spontaneity.

   ΔG = ΔH - TΔS. A process is spontaneous when ΔG < 0, at equilibrium when ΔG = 0, and non-spontaneous when ΔG > 0.

7. Relate standard Gibbs free energy change to the equilibrium constant.

   ΔG° = -RT ln K = -2.303 RT log K. A negative ΔG° (K > 1) favours products.

8. State Le Chatelier's principle.

   If a system at equilibrium is disturbed by a change in concentration, pressure, or temperature, the equilibrium shifts in the direction that counteracts (minimises) the disturbance.

9. Write the relationship between Kp and Kc for a gaseous equilibrium.

   Kp = Kc(RT)^Δn_g, where Δn_g = (moles of gaseous products) - (moles of gaseous reactants).

10. How does an increase in pressure affect a gaseous equilibrium, and when has it no effect?

    Increasing pressure (decreasing volume) shifts equilibrium toward the side with fewer moles of gas. It has no effect when Δn_g = 0 (equal moles of gas on both sides).

11. Define the reaction quotient Q and state how comparing Q with K predicts direction.

    Q has the same form as K but uses non-equilibrium concentrations. If Q < K reaction proceeds forward; Q > K reaction proceeds backward; Q = K system is at equilibrium.

12. State the Bronsted-Lowry definition of acids and bases.

    An acid is a proton (H+) donor and a base is a proton acceptor. A conjugate acid-base pair differs by one proton.

13. Write the relation between Ka and Kb of a conjugate acid-base pair (in water at 25°C).

    Ka × Kb = Kw = 1.0 × 10^-14, so pKa + pKb = 14.

14. Write the Henderson-Hasselbalch equation for an acidic buffer.

    pH = pKa + log([salt]/[acid]) = pKa + log([conjugate base]/[acid]). For a basic buffer: pOH = pKb + log([salt]/[base]).

15. Define solubility product (Ksp) and the condition for precipitation.

    Ksp is the equilibrium constant for the dissolution of a sparingly soluble salt into its ions. Precipitation occurs when the ionic product (Q_sp) exceeds Ksp; no precipitate forms if Q_sp < Ksp.

16. What is the common ion effect?

    The suppression of ionisation (or solubility) of a weak electrolyte/sparingly soluble salt by adding a strong electrolyte that has an ion in common with it, shifting equilibrium per Le Chatelier.

17. State Faraday's first law of electrolysis.

    The mass of substance deposited/liberated at an electrode is directly proportional to the quantity of electricity passed: m = Z·Q = Z·I·t, where Z is electrochemical equivalent.

18. Write the Nernst equation for an electrode/cell at 298 K.

    E = E° - (0.0591/n) log Q (at 298 K), where n is number of electrons transferred and Q is the reaction quotient. For a cell, E_cell = E°_cell - (0.0591/n) log Q.

19. How is the standard EMF of a cell related to ΔG° and to K?

    ΔG° = -nFE°_cell and E°_cell = (0.0591/n) log K at 298 K. A positive E°_cell means a spontaneous (galvanic) reaction.

20. Define molar conductivity and state how it varies with dilution for strong and weak electrolytes.

    Molar conductivity Λm = κ × 1000 / C. On dilution it increases for both; for strong electrolytes it rises gradually to Λm°, while for weak electrolytes it rises steeply near infinite dilution due to increased dissociation.

21. Write the integrated rate law and half-life expression for a first-order reaction.

    k = (2.303/t) log([A]0/[A]); t_1/2 = 0.693/k. The half-life of a first-order reaction is independent of initial concentration.

22. State the Arrhenius equation and define activation energy.

    k = A·e^(-Ea/RT), where A is the frequency/pre-exponential factor and Ea is the activation energy (minimum energy needed for an effective collision). ln k vs 1/T gives a straight line of slope -Ea/R.

23. Distinguish between order and molecularity of a reaction.

    Order is the experimentally determined sum of powers of concentration terms in the rate law (can be zero/fractional). Molecularity is the number of reacting species in an elementary step (always a positive whole number, max ~3).

24. Distinguish between physisorption and chemisorption, and define an adsorption isotherm.

    Physisorption: weak van der Waals forces, low enthalpy (20-40 kJ/mol), reversible, decreases with temperature, multilayer. Chemisorption: chemical bonds, high enthalpy (80-240 kJ/mol), specific, irreversible, monolayer. An adsorption isotherm (e.g., Freundlich x/m = k·P^(1/n)) relates amount adsorbed to pressure at constant temperature.

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