SRMJEEE Chemistry Flashcards

24 sample cards from the 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 ideal gas equation and identify each term.

   $$PV = nRT$$ where $P$ is pressure, $V$ is volume, $n$ is moles, $T$ is absolute temperature, and $R = 8.314\ \text{J K}^{-1}\text{mol}^{-1} = 0.0821\ \text{L atm K}^{-1}\text{mol}^{-1}$ is the gas constant.

2. State Dalton's law of partial pressures.

   The total pressure of a mixture of non-reacting gases equals the sum of the partial pressures of the individual gases: $$P_{\text{total}} = P_1 + P_2 + P_3 + \cdots$$

3. State Avogadro's law.

   Equal volumes of all gases at the same temperature and pressure contain equal numbers of molecules; equivalently, at fixed $T$ and $P$, $V \propto n$.

4. Write the expression for the root-mean-square speed of gas molecules from kinetic theory.

   $$u_{rms} = \sqrt{\frac{3RT}{M}} = \sqrt{\frac{3kT}{m}}$$ where $M$ is molar mass, $m$ is molecular mass, and $k$ is Boltzmann's constant.

5. Write the kinetic gas equation and the relation between average kinetic energy and temperature.

   $$PV = \frac{1}{3}mNu^{2}$$ and the average kinetic energy per mole is $E_k = \frac{3}{2}RT$, showing kinetic energy depends only on absolute temperature.

6. Write the van der Waals equation for real gases and explain the correction terms.

   $$\left(P + \frac{an^{2}}{V^{2}}\right)(V - nb) = nRT$$ The term $\frac{an^{2}}{V^{2}}$ corrects for intermolecular attractive forces and $nb$ corrects for the finite volume of the molecules.

7. Define compressibility factor $Z$ and state its value for an ideal gas.

   $$Z = \frac{PV}{nRT}$$ For an ideal gas $Z = 1$. $Z < 1$ indicates dominant attractive forces and $Z > 1$ indicates dominant repulsive forces.

8. Define surface tension and state how it varies with temperature.

   Surface tension is the force per unit length acting along the surface of a liquid (or energy per unit area) that minimizes surface area. It decreases with increasing temperature.

9. Define viscosity of a liquid and state how it changes with temperature.

   Viscosity is a measure of a liquid's resistance to flow, arising from internal friction between layers. Viscosity decreases with increasing temperature.

10. What is vapour pressure and when does a liquid boil?

    Vapour pressure is the pressure exerted by the vapour in equilibrium with its liquid at a given temperature. A liquid boils when its vapour pressure equals the external (atmospheric) pressure.

11. Define ionic and covalent bonding.

    Ionic bonding is the electrostatic attraction between oppositely charged ions formed by transfer of electrons (typically metal to non-metal). Covalent bonding is the sharing of electron pairs between atoms (typically between non-metals).

12. State the octet rule.

    Atoms tend to gain, lose, or share electrons so as to achieve a stable noble-gas configuration of eight electrons in their valence shell.

13. Define lattice energy and state Fajans' rules for covalent character in ionic bonds.

    Lattice energy is the energy released when gaseous ions combine to form one mole of an ionic solid. Fajans' rules: covalent character increases with small cation size, large anion size, and high charge on the ions (high polarizing power/polarizability).

14. What does VSEPR theory predict and what is the basic principle?

    VSEPR (Valence Shell Electron Pair Repulsion) theory predicts molecular geometry based on minimizing repulsion between electron pairs around the central atom. Repulsion order: lone pair–lone pair > lone pair–bond pair > bond pair–bond pair.

15. Give the hybridisation, shape and bond angle for $\ce{sp}$, $\ce{sp^2}$ and $\ce{sp^3}$ hybridisation.

    $\ce{sp}$: linear, $180^\circ$; $\ce{sp^2}$: trigonal planar, $120^\circ$; $\ce{sp^3}$: tetrahedral, $109.5^\circ$.

16. State the shape and bond angle of $\ce{H2O}$ and $\ce{NH3}$, and explain the deviation from $109.5^\circ$.

    $\ce{H2O}$ is bent with bond angle $\approx 104.5^\circ$ and $\ce{NH3}$ is trigonal pyramidal with $\approx 107^\circ$. Both are $\ce{sp^3}$ hybridised; lone pair–bond pair repulsion compresses the angle below the ideal $109.5^\circ$.

17. What is the hybridisation, shape and bond angle of $\ce{PCl5}$ and $\ce{SF6}$?

    $\ce{PCl5}$: $\ce{sp^3d}$ hybridisation, trigonal bipyramidal, bond angles $90^\circ$ and $120^\circ$. $\ce{SF6}$: $\ce{sp^3d^2}$ hybridisation, octahedral, bond angle $90^\circ$.

18. State the formula for bond order in molecular orbital theory.

    $$\text{Bond order} = \frac{1}{2}\left(N_b - N_a\right)$$ where $N_b$ is the number of electrons in bonding molecular orbitals and $N_a$ in antibonding molecular orbitals.

19. Use MO theory to explain why $\ce{O2}$ is paramagnetic.

    In $\ce{O2}$ the two highest electrons occupy the degenerate $\pi^{*}_{2p}$ antibonding orbitals singly with parallel spins (Hund's rule), giving two unpaired electrons; these unpaired electrons make $\ce{O2}$ paramagnetic. Its bond order is $2$.

20. Why does the $\ce{He2}$ molecule not exist according to MO theory?

    $\ce{He2}$ would have configuration $\sigma_{1s}^{2}\sigma_{1s}^{*2}$, giving bond order $= \frac{1}{2}(2-2) = 0$; with zero net bonding the molecule is unstable and does not exist.

21. State the first law of thermodynamics and its mathematical form.

    Energy can neither be created nor destroyed (conservation of energy). The change in internal energy is $$\Delta U = q + W$$ where $q$ is heat added to the system and $W$ is work done on the system.

22. Define enthalpy and give the relation between $\Delta H$ and $\Delta U$.

    Enthalpy is $H = U + PV$. For a reaction at constant pressure, $$\Delta H = \Delta U + \Delta(PV) = \Delta U + \Delta n_g RT$$ where $\Delta n_g$ is the change in moles of gas.

23. 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 the initial and final states. Thus enthalpy changes of individual steps can be added to obtain the overall $\Delta H$.

24. Distinguish between exothermic and endothermic reactions in terms of the sign of $\Delta H$.

    In an exothermic reaction heat is released and $\Delta H < 0$ (negative). In an endothermic reaction heat is absorbed and $\Delta H > 0$ (positive).

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