1. Define Limiting Reagent. 1

2. State Aufbaus principle. 1

3. How many unpaired electrons are present in Cr+3 ion (atomic no. of Cr =24 ).


4. Which has higher ve electron gain enthalpy F or Cl . 1

5. Explain why does ice floats over water ? 1

6. Arrange the following hybrid orbitals in increasing order of electronegativity : sp , sp2 , sp3 . 1

7. Define absolute zero. 1

8. Predict the sign of entropy change when ammonium chloride is dissolved in water. 1

9. State and explain the law of multiple proportion with example. 2

10. The density of 3 M solution of NaCl is 1.25 g/ml. Calculate molality of the solution. 2

11. Calculate the number of molecules and number of atoms present in 5.60 L of O3 at NTP. 2

12. How many electrons in an atom can have the following set of quantum numbers ?

a. n=4 , ms= -1/2

b. n=3 , l = 0 2

13. What transition in the hydrogen spectrum would have the same wavelength as the Balmer transition n=4 to n=2 of He+ spectrum ? 2

14. [Rn] has atomic no. 86 and is a noble gas. What will be the atomic number of the next noble gas. Write its symbol and IUPAC name. 2

15. Which of the following pairs of elements would have a more negative electron gain enthalpy?

a. O or F

b. N or P 2

16. What are isoelectronic species? Arrange the following ions in increasing order of size :

K+ , Ca+2 , P-3 , S-2 ,Cl- 2

17. Which of the following pair of compounds have higher dipole moment? Explain with reason.

a. NH3 or NF3

b. CS2 or OCS 2

18. Use molecular orbital theory to explain why Be2 molecules does not exit.


19. A 5.2 molal aqueous solution of methyl alcohol, CH3OH, is supplied. Calculate the mole fraction of methyl alcohol in the solution? (3)

20. The electron energy in hydrogen atom is given by En = (-2.18 × 10-18)/n2 J. Calculate the energy required to remove an electron completely from the n = 2 orbit. What is the longest wavelength of light in cm that can be used to cause this transition? (3)

21. (i) Which has the larger atomic size and why? (1)

Al or Ga

(ii) Write the general electronic configuration of f- block elements. (1)

(iii) Explain why noble gases have positive electron gain enthalpy? (1)

22. (i) Draw the shapes of the following molecules on the basis of VSEPR theory. (2)

(a) ICl4 - (b) XeF2

(ii)Which has higher lattice enthalpy and why?

NaOH or Mg(OH)2 (1)


(i) Which is more ionic and why? (2)

MgCl2 or CaCl2

(ii) Describe the change in hybridisation (if any) of the Al atom in the following reaction.

AlCl3 + Cl → AlCl4 (1)

23. (i) What is meant by dispersion forces or London Forces? (1)

(ii) 'a' and 'b' are van der Waals' constants for gases.

Chlorine is more easily liquefied than ethane; explain on the basis of Van der Waals equation. (2)

24. (i) Spontaneous process for a gas on solid surface is an exothermic process. Explain. (2)

(ii) What do you mean by thermodynamic reversible process. (1)

25.(i) Calculate the difference between heat of reaction at constant pressure and constant volume for the reaction. (2)

2C6H6(l) + 15CO2(g) → 12 CO2(g) + 6 H20(l)

(ii) Give the relation between Gibbs free energy change and equilibrium constant.


26. Using the equation of state pV=nRT; show that at a given temperature density of a gas is proportional to gas pressure p. (3)

27. When 2 g of a gas A is introduced into an evacuated flask kept at 25o C , the pressure is found to be one atmosphere. If 3 g of another gas B is then added to the same flask, the total pressure becomes 1.5 atm. Assuming ideal gas behaviour, calculate the ratio of molecular weights. (3)

28. (i) Define Entropy. Explain that entropy change is a state function. (3)

(ii) Predict in which of the following, entropy increases/decreases: (2)

(a) A liquid crystallizes into a solid.

(b) Temperature of a crystalline solid is raised from 0 K to 115 K.


Calculate the standard enthalpy of formation of CH3OH (l) from the following data: (5)

CH3OH(l) + 3/2O2 → CO2 + 2H2O(g) ∆rH = -726 kj/mol

C(graphite) + O2 → CO2(g) ∆cH = -393kj/mol

H2(g) + ½ O2(g) → H2O ∆fH = -286kj/mol

29. (i) Draw the lewis structures of NO2- and calculate the formal charge on each atom. (2)

(ii) Expalin why the bond angle(H-N-H) in ammonia is not a normal tetrahedral angle although N is sp3 hybridised. (2)

(iii) Explain why the bond lengths (S-O) are same in SO2 molecule. (1)


(i) Describe the hybridisation in case of PCl5. Why are the axial bonds longer as

compared to equatorial bonds ?

(i) Define hydrogen bond. Is it weaker or stronger than the van der Waals forces?

(iii) Arrange the bonds in order of increasing ionic character in the molecules: LiF, K2O, N2, SO2 and ClF3.

30. (i) Which hydrogen like species will have the same radius as that of Bohrs first orbit of hydrogen atoms? (2)

n=2, Li2+ or n = 2 Be3+

(ii) Give the number of radial nodes of 3s and 2p orbital. (2)

(iii) What is the maximum number of electrons that may be present in all atomic orbitals with principal quantum number 3 and azimuthal quantum number 2. (1)


(i) What are the frequency and wavelength of a photon emitted during a transition from n = 5 state to the n = 2 state in the hydrogen atom? (3)

(ii) What is the maximum number of emission lines when the excited electron of a H atom in n = 6 drops to the ground state? (1)

(iii) What is Heisenbergs uncertainty principle?

In order to get the answers of all the queries please post them separately. However, we are providing the answers of some of your problems.
Kindly refer to the following links:
Limiting Reagent:
Aufbaus Principle:

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1 the reactant that is present in limiting amount is called limiting reagent. it affects the formation of compound.

2 aufbau's principle-- the electrons are added progressively to the various orbtals in the increasing order of the energy starting from the lowest energy.

3 one unpaired electron.

5 because density of ice is less than that of water.

12 a) 16

b) 6

16 the elements which have same no. of electrons are called isoelectronic species.


heisenberg's uncertainity principle-- the position and momentum of an electron cannot be found accurately simultaneously.

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