Chemistry Notes (MCQs,Notes) -1st year and 2nd year.







Chemistry MCQs intermidate XI MCQs



1. The process in which a solid directly changes to vapours without melting is called __________.

(Evaporation, Condensation, Sublimation)

2. The oxidation number of P in PO3-4 is __________.

(3+, 5+, 3-)

3. The pH of 0.001 M HCl is __________.

(2, 4, 3)

4. K ( rate constant) is dependent on __________.

(temperature, concentration, volume)

5. The universal indicator in water shows the colour __________.

(red, green, blue)

6. The pH of blood is __________.

(7.3, 8.4, 5.6)

7. The oxidation potential of hydrogen electrode is __________.

(0.0 volt, +0.76volt, -0.36volt)

8. __________ quantum number describes the shape of a molecule.

(Pricipal, Azimuthal, Spin)

9. An orbital can have the maximum number of two electrons but with opposite spin, it is called __________.

(Pauli’s Exclusion Principle, Hund’s Rule, Aufbau Principle)

10. When a-particle is emitted from the nucleus of radioactive element, the mass number of the atom __________.

(Increases, Decreases, Does not change)

11. Dissociation of KclO3 is a __________ process.

(Reversible, Irreversible)

12. The e/m ratio of cathode rays is the __________ when Hydrogen is taken in the discharge tube.

(Lowest, Highest)

13. The negative ion tends to expand with the __________ of negative change on it.

(Decreases, Increases)

14. Ionic compounds have __________ melting points.

(Low, High)

15. The allotropic forms of an element are called __________.

(Polymorphs, Isomorphs)

16. Absolute Zero is equal to __________.

(273.16°C, -273.16°C)

17. The compounds having hydrogen bond generally have __________ boiling points.

(High, Low)

18. Surface tension __________ with the rise of temperature.

(Increases, Decreases)

19. Mercury forms __________ meniscus in a glass tube.

(concave, convex)

20. The reactions with the high value of energy of activation are __________.

(slow, fast)

21. 2.000 has/have __________ significant figure(s).

(1, 4)

22. E + PV is called __________.

(Entropy, Enthalpy)

23. The shorter the bond length in a molecule, the __________ will be bond energy.

(Lesser, Greater)

24. Positive rays are produced from __________.

(anode, Cathode, Ionization of gas in a discharge tube)

25. __________ of the following contains the fewer number of molecules.

(1 gm of hydrogen, 4 gm of oxygen, 2 gm of nitrogen)

26. the true statement about the average speed of the molecules of hydrogen, oxygen and nitrogen confined in a container is __________.

(Hydrogen is quicker, Oxygen is quicker, The molecules of all the gases have the same average speed)

27. The correct statement about the glass is __________.

(It is crystalline solid, Its atoms are arranged in an orderly fashion, It is a super cooled liquid)

28. When a substance that has absorbed energy emits it in the form of radiation the spectrum obtained is __________.

(Continuous Spectrum, Line Spectrum, Emission Spectrum)

29. __________ of the overlap forms strong bond.

(S-S, P-S, P-P)

30. __________ compound has a greater angle between a covalent bond.

(H2O, NH3, CO2)

31. When sodium chloride is mixed in water then __________.

(pH is changed, NaOH and HCl are formed, Sodium and chloride ions become hydrated)

32. The boiling point of a liquid __________ with an increase in pressure.

(Decreases, Increases, remains constant)

33. An Azimuthal Quantum Number describes the __________.

(size of an atom, shape of an orbital, spin of orbital)

34. The rate of the backward reaction is directly proportional to the product of the molar concentration of __________.

(Reactants, Products, None of them)

Chapter 1

Introduction To Fundamental Concepts


1. The formula, which gives the simple ratio of each kind of atoms present in the molecule of compound, is called __________.

(Molecular Formula, Empirical Formula, Structural Formula)

2. The formula, which expresses the actual number of each kind of atom present in the molecule of a compound, is called __________.

(Empirical Formula, Molecular Formula, Structural Formula)

3. Mole is a quantity, which has __________ particles of the substance.

(One billion, 6.02 x 1023, 1.013 x 105)

4. The simplest formula of a compound that contain 81.8% carbon and 18.2% hydrogen is __________.

(CH3, CH, C2H6)

5. The empirical Formula of a compound __________.

(is always the same as the molecular formula, Indicates the exact composition, Indicates the simplest ratio of the atoms)

6. Very small and very large quantities are expressed in terms of __________.

(significant figures, Exponential Notation, Logarithm)

7. Two moles of water contains __________ molecules.

(6.02 x 1023, 1.204 x 1024, 3.01 x 1023)

8. One mole of Cl- ions contains __________ ions.

(6.02 x 1023, 1.204 x 1024, 3.01 x 1023)

9. 220 gms of CO2 contains __________ moles of CO2.

(One, Five, Ten)

10. In rounding off __________ figure is dropped.

(First, Last, No)

11. Precision is linked with __________.

(Individual measurements, Actual results, Accepted Value)

12. Accuracy refers to how closely a measured value agrees with __________.

(Individual result, Actual result, Average value)

13. 6600 contains __________ significant figures.

(2, 3, 4)

14. 3.7 x 104 contains __________ significant figures.

(2, 3, 5)

15. 9.40 x 10-19 contains __________ significant figures.

(2, 3, 5)

16. The figure 39.45 will be rounded off to __________.

(39.4, 39.5, 39)

17. __________ means that the result obtained in different experiments are very close to the accepted values.

(Accuracy, Precision, Significant Figure)

18. The average weight of atoms of an element as compared to the weight of one atom of carbon taken as __________ is called the atomic weight.

(12, 13, 14)

19. 58.5 is __________ of NaCl.

(Atomic weight, Formula Weight, Molecular Weight)

20. 18.0 a.m.u is the __________ weight of water.

(Atomic, Formula, Molecular)

21. 28 gms of nitrogen will have __________ molecules.

(6.02 x 1023, 12.04 x 1023, 3.01 x 1023)

22. 22.4 dm3 of CO2 is __________ 22.4 dm3 of SO2.

(Heavier than, Lighter than, Equal to)

23. 100 gms of water is equal to __________ moles.

(5.56, 27.78, 6.25)

24. The reactions, which proceed in both the directions are called __________ reactions.

(Reversible, Irreversible, Neutrilization)

25. The reactions, which proceed in forward direction only are called __________ reactions.

(Reversible, Irreversible, Ionic)

26. Molecular weight is used for __________ substances.

(Ionic, Non ionic, Neutral)

27. Formula weight is used for __________ substances.

(Ionic, Non ionic, Neutral)

28. The modern system of measurement is called __________ system.

(SI, Metric, F.P.S)

29. The S.I unit of mass is __________.

(kilogram, gram, pound)

30. One mole of glucose contains __________ gms.

(100, 180, 342)

Chapter 2

The Three States of Matter

1. __________ was the first scientist who expressed a relation between pressure and the volume of a gas.

(Charles, Boyle, Avogadro)

2. If the pressure upon a gas confined in a vessel varies, the temperature remaining same, the volume will __________.

(Vary directly as the pressure, Vary inversely as the temperature, Vary inversely as the pressure)

3. The statement concerning the relation of temperature to the volume of a gas under fixed pressure was first synthesized by __________.

(Boyle, Charles, Avogadro)

4. Absolute Zero is __________.

(273°C, -273°C, -273°K)

5. Gases intermix to form __________.

(Homoge\= ous mixture, Heterogenous mixture, compound)

6. Water can exists in __________ physical states at a certain condition of temperature pressure.

(One, Two, three)

7. The temperature at which the volume of a gas theoretically becomes zero is called __________.

(Transition temperature, Critical Temperature, Absolute Zero)

8. Gases deviate from ideal behaviour at __________ pressure and __________ temperature.

(Low, High, Normal)

9. Very low temperature can by produced by the __________ of gases.

(Expansionn, Contraction, Compression)

10. Boiling point of a liquid __________ with increase in pressure.

(increases, decreases, remains same)

11. 273°K = __________

(100°C, 273°C, 0°C)

12. -273°C is equal to __________.

(0°K, 273°K, 100°K)

13. Evaporation takes place at __________.

(All temperatures, At constant temperature, at 100°C)

14. __________ is the temperature at which the vapour pressure of a liquid becomes equal to atmospheric pressure.

15. The freezing point of water in Fahrenheit scale is __________.

(0°F, 32°F, 212°F)

16. All gases change to solid before reaching to __________.

(-100°C, 0°C, -273°C)

17. Pressure of the gas is due __________ of the molecules on the wall of the vessel.

(Collisionns, Attraction, Repulsion)

18. Boiling point of water in absolute scale is __________.

(212°K, 100°K, 373°K)

19. Boyle’s Law relates __________.

(Pressure and volume, Temperature and volume, Pressure and temperature)

20. Charles Law deals with __________ relationship.

(temperature and volume, pressure and volume, temperature and pressure)

21. Effusion is the escape of gas through __________.

(A small pin hole, Semi permeable membrane, porous container)

22. The expression P = P1 + P2 + P3 represents __________ mathematically.

(Graham’s Law, Avogadro’s Law, Dalton’s law of partial Pressure)

23. According to __________ equal volumes of all gases at the same temperature and pressure contain equal number of molecules.

(Graham’s Law, Avogadro’s Law, Dalton’s Law)

24. The boiling point of pure water is __________.

(32°C, 100°F, 373°K)

25. The internal resistance of a liquid to flow is called __________.

(Surface tension, Capillary action, Viscosity)

26. The existence of different crystals forms of the same substance is called __________.

(Isomorphism, Polymorphism, Isotopes)

27. Rate of Evaporation __________ on increasing temperature.

(Increases, Decreases, Remains same)

28. The temperature at which more than one crystalline forms of a substance coexist is called the __________.

(Critical Temperature, Transition Temperature, Absolute Temperature)

29. The gases which strictly obey the gas laws are called __________.

(Ideal gases, Permanent gases, Absolute gases)

30. Lighter gas diffuse __________ than the heavier gases.

(More readily, Less readily, Very slowly)

Chapter 3

Structure of Atom

1. The charge on an electron is __________.

(-2.46 x 104 coulombs, -1.6 x 10-19 coulombs, 1.6 x 10-9coulombs)

2. The maximum number of electrons that can accommodated by a p-orbital is __________.

(2, 6, 10)

3. A proton is __________.

(a helium ion, a positively charged particle of mass 1.67 x 10-27 kg, a positively charged particle of mass 1/1837 that of Hydrogen atom)

4. Most penetrating radiation of a radioactive element is __________.

(a-rays, b-rays, g-rays)

5. The fundamental particles of an atom are __________.

(Electrons and protons, electrons and neutrons, Electrons, Protons, Neutrons)

6. The fundamental particles of an atoms are __________.

(the number of protons, The number of neutrons, The sum of protons and neutrons)

7. “No two electrons in the same atoms can have identical set of four quantum numbers.” This statement is known as __________.

(Pauli’s Exclusion Principle, Hund’s rule, Aufbau Rule)

8. __________ has the highest electronegativity value.

(Fluorine, Chlorine, Bromine)

9. Principle Quantum number describes __________.

(Shape of orbital , size of the orbital, Spin of electron in the orbital)

10. Canal rays are produced from __________.

(Anode, Cathode, Ionization of gas in the discharge tube)

11. Electromagnetic radiation produce from nuclear reactions are known as __________.

(a-rays, b-rays, g-rays)

12. Cathode rays consist of __________.

(Electorns, Protons, Positrons)

13. The properties of cathode rays __________ upon the nature of the gas inside the tube.

(depend, partially depend, do not depend)

14. Anode rays consists of __________ particles.

(Negative, Positive, Neutral)

15. Atomic mass of an element is equal to the sum of __________.

(electrons and protons, protons and neutrons, electrons and neutrons)

16. Neutrons were discovered by __________.

(Faraday, Dalton, Chadwick)

17. The value of Plank’s constant is __________.

(6.626 x 10-34, 6.023 x 1024, 1.667 x 10-28)

18. P-orbitals are __________ in shape.

(spherical, diagonal, dumb bell)

19. The removal of an electron from an atom in gaseous state is called __________.

(Ionization energy, Electron Affinity, Electronegativity)

20. The energy released when an electron is added to an atom in the gaseous state is called __________.

(Ionization Potential, electron Affinity, Electronegativity)

21. The power of an atom to attract a shared pair of electrons is called __________.

(Ionization Potential, Electron Affinity, Electronegativity)

22. Electronegativity of Fluorine is arbitrarily fixed as __________.

(2, 3, 4)

23. The energy difference between the shells go on __________ when moved away from the nucleus.

(Increasing, decreasing, equalizing)

24. __________ discovered that the nucleus of an atom is positively charged.

(William Crooke’s, Rutherford, Dalton)

25. Isotopes are atoms having same __________ but different __________.

(Atomic weight, Atomic number, Avogadro’s Number)

26. __________ consists of Helium Nuclei or Helium ion (He++).

27. The angular momentum of an electron revolving around the nucleus of atom is __________.

(nh/2p, n2h2/2p, nh3/3p)

28. The wavelengths of X-rays are mathematically related to the __________ of anticathode element.

(atomic weight, atomic number, Avogadro’s number)

29. Lyman Series of spectral lines appear in the __________ portion of spectrum.

(Ultraviolet, Infra red, Visible)

30. According to __________ electrons are always filled in order of increasing energy.

(Pauli’s Exclusion Principle, Uncertainty Principle, Aufbau Principle)

Chapter 4

Chemical Bonding

1. The energy required to break a chemical bond to form neutral atoms is called __________.

(Ionization Potential, Electron Affinity, Bond Energy)

2. The chemical bond present in H-Cl is __________.

(Non Polar, Polar Covalent, Electrovalent)

3. A polar covalent bond is formed between two atoms when the difference between their E.N values is __________.

(Equal to 1.7, less than 1.7, More than 1.7)

4. The most polar covalent bond out of the following is __________.

(H-Cl, H-F, H-I)

5. __________ bond is one in which an electron has been completely transferred from one atom to another.

(Ionic, Covalent, co-ordinate)

6. __________ bond is one in which an electron pair is shared equally between the two atoms.

(Ionic, Covalent, Co-ordinate)

7. Bond angle in the molecule of CH4 is of __________.

(120°, 109.5°, 180°)

8. A molecule of CO2 has __________ structure.

9. The sigma bond is __________ than pi bond.

(Weaker, Stronger, Unstable)

10. The sp3 orbitals are __________ in shape.

(Tetrahedral, Trigonal, Diagonal)

11. The shape of CH4 molecule is __________.

(Tetrahedral, Trigonal, Diagonal)

12. The bond in Cl2 is __________.

(Non polar, Polar, Electrovalent)

13. Water is __________ molecule.

(None polar, Polar, Electrovalent)

14. Covalent bonds in which electron pair are shared equally between the two atoms is called __________ covalent bond.

(Non polar, Polar, Co-ordinate)

15. Each carbon atom in CH4 is __________ hybridized.

(Sp3, Sp2, Sp)

16. Each carbon atom in C2H4 is __________ hybridized.

(Sp3, Sp2, Sp)

17. Each carbon atom in C2H2 is __________ hybridized.

(Sp3, Sp2, Sp)

18. Oxygen atom in H2O has __________ unshared electron pair.

(One, two , three)

19. Nitrogen atom in NH3 has __________ unshared electron pair.

(One, two, three)

20. The cloud of charge that surrounds two or more nuclei is called __________ orbital.

(Atomic, Molecular, Hybrid)

21. A substance, which is highly attracted by a magnetic field, is called __________.

(Electromagnetic, Paramagnetic, Diamagnetic)

22. HF exists in liquid due to __________.

(Vander Waal Forces, Hydrogen bond, covalent Bond)

23. Best hydrogen bonding is found in __________

(HF, HCl, HI)

24. Shape of CCl4 molecule is __________.

(tetrahedral, Trigonal, Diagonal)

25. __________ bond is formed due to linear overlap.

(Sigma bond, Pi bond, Hydrogen bond)

26. __________ is defined as the quantity of energy required to break one mole of covalent in gaseous state.

(Bond energy, Ionization energy, Energy of Activation)

27. Repulsive force between electron pair in a molecule is maximum when it has an angle of __________.

(120°, 109.5°, 180°)

28. Repulsive force between electron pair in a molecule is maximum when it has an angle of __________.

(120°, 109.5°, 180°)

29. The sum of total number of electrons pairs (bonding and lone pairs) is called __________.

(Atomic Number, Avogadro’s Number, Steric Number)

30. Shape of __________ molecule is tetrahedral.

(BaCl2, BF3, NH3)

Chapter 5

Energetics of Chemical Reaction

1. The quantity of heat evolved or absorbed during a chemical reaction is called __________.

(Heat or Reaction, Heat of Formation, Heat of Combination)

2. An endothermic reaction is one, which occurs __________.

(With evolution of heat, With absorption of Heat, In forward Direction)

3. An exothermic reaction is one during which __________.

(Heat is liberated, Heat is absorbed, no change of heat occurs)

4. The equation C + O2 ® CO2 DH = -408KJ represents __________ reaction.

(Endothermic, Exothermic, Reversible)

5. The equation N2 + O2 ® 2NO DH = 180KJ represents __________ reaction.

(Endothermic, Exothermic, Irreversible)

6. Thermo-chemistry deals with __________.

(Thermal Chemistry, Mechanical Energy, Potential Energy)

7. Enthalpy is __________.

(Heat content, Internal energy, Potential Energy)

8. Hess’s Law is also known as __________.

(Law of conservation of Mass, Law of conservation of Energy, Law of Mass Action)

9. Any thing under examination in the Laboratory is called __________.

(Reactant, System, Electrolyte)

10. The environment in which the system is studied in the laboratory is called __________.

(Conditions, Surroundings, State)

11. When the bonds being broken are more than those being formed in a chemical reaction, then DH will be __________.

(Positive, Negative, Zero)

12. When the bond being formed are more than those being broken in a chemical reaction, then the DH will be __________.

(Positive, Negative, Zero)

13. The enthalpy change when a reaction is completed in single step will be __________ as compared to that when it is completed in more than one steps.

(Equal to, Partially different from, Entirely different from)

14. The enthalpy of a system is represent by __________.

(H, DH, DE)

15. The factor E + PV is known as __________.

(Heat content, Change in Enthalpy, Work done)

16. Heat of formation is represented by __________.

(Df, DHf, Hf)

17. The heat absorbed by the system at constant __________ is completely utilize to increase the internal energy of the system.

(Volume, Pressure, Temperature)

18. Heat change at constant __________ from initial to final state is simply equal to the change in enthalpy.

(Volume, Pressure, Temperature)

19. A system, which exchange both energy and energy with the surrounding, is __________ system.

(Open, Closed, Isolated)

20. A system, which only exchange energy with the surrounding but not the matter, is __________ system.

(Open, Closed, Isolated)

21. A system, which neither exchanges energy nor matter with the surroundings is __________ system.

(Open, Closed, Isolated)

22. __________ property of a system is independent of the amount of material concerned.

(Intensive, Extensive, Physical)

23. __________ property of a system depends upon the amount of substance present in the system.

(Intensive, Extensive, Physical)

24. DE = q – w represents __________.

(First Law of Thermodynamics, Hess’s Law, Enthalpy Change)

25. __________ is defined as the change in enthalpy when one gram mole of a compound is produced from its elements.

(Heat of Reaction, heat of Formation, Heat of Neutrilization)

Chapter 6

Chemical Equilibrium

1. At equilibrium the rate of forward reaction and the rate of reverse reaction are __________.

(Equal, Changing, Different)

2. Such reactions, which proceed to forward direction only and are completed after sometime are called __________ reaction.

(Irreversible, Reversible, Molecular)

3. Such reactions, which proceed to both the direction and are never completed, are called __________ reaction.

(Irreversible, Reversible, Molecular)

4. The rate of chemical reaction is directly proportional to the product of the molar concentration of __________.

(Reactants, Products, Both reactants and products)

5. “If a system in equilibrium is subjected to a stress, the equilibrium shifts in a direction to minimize or undo the effect of this stress. This principle is known as __________.

(Le-Chatelier’s Principle, Gay Lussac’s Principle, Avogadro’s Principle)

6. A very large value of Kc indicates that reactants are __________.

(very stable, unstable, moderately stable)

7. A very low value of Kc indicates that reactants are __________.

(very stable, very unstable, moderately stable)

8. The equilibrium in which reactants are products are in single phase is called __________.

(Homogenous Equilibrium, Heterogenous Equilibrium, Dynamic Equilibrium)

9. The equilibrium in which reactants and products are in more than one phases are called __________.

(Homogenious Equilibrium, Heterogenious Equilibrium, Dynamic Equilibrium)

10. Chemical Equilibrium is __________ equilibrium.

(Dunamic, Static, Heterogeneous)

11. In exothermic reaction, lowering of temperature will shift the equilibrium to __________.

(right, left, equally on both the direction)

12. In endothermic reaction, lowering of temperature will shift the equilibrium to __________.

(right, left, equally on both the direction)

13. A catalyst __________ the energy of activation.

(increases, decreases, has no effect on)

14. At equilibrium point __________.

(forward reaction is increased, backward reaction is increased, forward and backward reactions become equal)

15. NH3 is prepared by the reaction N2 + 3H2 Û 2NH3 DH = -21.9 Kcal. The maximum yield of NH3 is obtained __________.

(At low temperature and high pressure, at high temperature and low pressure, at high temperature and high pressure)

16. When a high pressure is applied to the following reversible process: N2 + O2 Û 2NO The equilibrium will __________

(shift to the forward direction, shift to the backward direction, not change)

17. The value of Kc __________ upon the initial concentration of the reaction.

(depends, partially depends, does not depend)

18. While writing the Kc expression, the concentration of __________ are taken in the numerator.

19. Solubility product constant is denoted by __________.

(Kc, Ksp, Kr)

20. “The degree of ionization of an electrolyte is suppressed by the addition of another electrolyte containing a common ion.” This phenomenon is called __________.

(Solubility Product, Common Ion Effect, Le-Chatelier’s Principle)

Chapter 7

Solutions and Electrolytes

1. Molarity is the number of moles of a solute dissolved per __________.

(dm3 of a solution, dm3 of solvent, Kg of solvent)

2. Molality is defined as the number of moles of solute dissolved per __________.

(dm3 of solution, kg of solvent, kg of solute)

3. The solubility of a solute __________ with the increase of temperature.

(increases, decreases, does not alter)

4. The loss of electron during a chemical reaction is known as __________.

(Oxidation, Reduction, Neutralization)

5. The gain of electron during a chemical reaction is known as __________.

(Oxidation, Reduction, Neutralization)

6. The ions, which are attracted towards the anode, are known as __________.

(Anins, Cations, Positron.

7. The pH of a neutral solution is __________.

(1.7, 7, 14)

8. A current of one ampere flowing for one minute is equal to __________.

(One coulomb, 60 coulomb, one Faraday)

9. A substance, which does not allow electricity to pass through, is known as __________.

(Insulator, Conductor, Electrolyte)

10. Such substances, which allow electricity to pass through them and are chemically decomposed, are called __________.

(Electrolytes, Insulators, Metallic conductors)

11. __________ is an example of strong acid.

(Acetic Acid, Carbonic Acid, Hydrochloric Acid)

12. __________ is an example of weak acid.

(Hydrochloric Acid, Acetic Acid, Sulphuric Acid)

13. When NH4Cl is hydrolyzed, the solution will be __________.

(Acidic, Basic, Neutral)

14. When Na2CO3 is hydrolyzed, the solution will be __________.

(Acidic, Basic, Neutral)

15. When blue hydrated copper sulphate is heated __________.

(It changes into white, it turns black, it remains blue)

16. Sulphur has the highest oxidation number in __________.

(SO2, H2SO4, H2SO3)

17. The reaction between an acid and a base to form a salt and water is called __________.

(Hydration, Hydrolysis, Neutralization)

18. __________ is opposite of Neutralization.

(Hydration, Hydrolysis, Ionization)

19. The substance having pH value 7 is __________.

(Basic, Acidic, Neutral)

20. An aqueous solution whose pH is zero is __________.

(Alkaline, Neutral, Strongly Acidic)

21. Solubility product of slightly soluble salt is denoted by __________.

(Kc, Kp, Ksp)

22. The increase of oxidation number is known as __________.

(Oxidation, Reduction, Hydrolysis)

23. The decrease of Oxidation number is known as __________.

(Oxidation, Reduction, Electrolysis)

24. One molar solution of glucose contains __________ gms of glucose per dm3 of solution.

* 180, 100, 342)

25. The number of moles of solute present per dm3 of solution is called __________.

(Molality, Molarity, Normality)

26. ‘M’ is the symbol used for representing __________.

(Molality, Molarity, Normality)

27. 1 mole of H2SO4 is equal to __________.

(98gms, 49gms, 180gms)

28. Buffer solution tends to __________ pH.

(Change, Increase, maintain)

29. The logarithm of reciprocal of hydroxide ion is represented as __________.

(pH, pOH, POH)

30. In __________ water molecules surround solute particles.

(Hydration, Hydrolysis, Neutralization)

Chapter 8

Introduction to Chemical Kinetics

1. The rate of chemical reaction __________ with increase in concentration of the reactants.

(Increases, Decreases, Does not alter)

2. Ionic reactions of inorganic compounds are __________.

(very slow, moderately slow, very fast)

3. The rate of __________ reactions can be determined.

(Very Slow, Moderately Slow, Very fast)

4. The sum of exponents of the concentrations of reactants is called __________.

(Order of reaction, Molecularity, Equilibrium Constant)

5. The rate of reaction generally __________ in the presence of a suitable catalyst.

(Increases, Decreases, remains constant)

6. The rate of a reaction __________ upon the temperature.

(depends, slightly depends, does not depends)

7. The minimum energy required to bring about a chemical reaction is called __________.

(Bond energy, Ionization energy, Energy of Activation)

8. Oxidation of SO2 in the presence of V2O5 in Sulphuric Acid industry is an example of __________.

(Homogenous catalyst, Heterogeneous catalyst, Negative catalyst)

9. Hydrolyses of ester in the presence of acid is an example of __________.

(Homogenous catalyst, Heterogeneous catalyst, Negative catalyst)

10. Concentration of the reactants __________ with the passage of time during a chemical reaction.

(Increases, Decreases, Does not alter)

11. Concentration of the products __________ with the passage of time during a chemical reaction.

(Increases, Decreases, Does not alter)

12. The rate constant __________ with temperature for a single reaction.

(Varies, Slightly Varies, Does not vary)

13. The rate of reaction at a particular time is called __________.

(Average Rate of reaction, Absolute rate of reaction, Instantaneous rate of reaction)

14. The specific rate constant K has __________ value for all concentrations of the reactant.

(Fixed, Variable, negligible value)

15. By increasing the surface area the rate of reaction can be __________.

(Increased, Decreased, Doubled)

16. MnO2 when heated with KClO3 __________.

(Gives up its own oxygen, Produces ozone O3, Acts as catalyst)

17. Reactions with high energy of activation proceed with __________.

(High speed, Moderately slow speed, slow speed)

18. The minimum amount of energy required to bring about a chemical reaction is called __________.

(Energy of ionization, Energy of Activation, Energy of Collision)

19. An inhibitor is a catalyst which __________ rate of reaction.

(Increases, Decreases, Does not alter)

20. __________ is the change of the concentration of reactant divided by the time.

(Rate of reaction, Velocity Constant, Molecularity)

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Numerical Chemistry.




1. Simplify according to the rule of significant figure .

2. The atomic mass of Zn is 65.4 a.m.u. Calculate (i) the number of moles and also the number of atoms in 10.9 gm of Zn. (ii) The mass of 1.204 x 1024 atoms of Zn in gm.

3. Adipic acid is used in the manufacture of Nylon. The acid contains 49.3%C, 6.9%H and 43.6%O by mass. The molecular mass of the acid is 146 a.m.u. Find the molecular formula of the Adipic Acid.

4. Calculate the value of R (Gas constant) with the help of Gas Equation when (i) the pressure is in atmosphere and the volume in dm3 or litre. (ii) the pressure is in Nm-2 and the volume is in cubic metre.

5. 400cm3 of helium gas effuse from a porous container in 20 seconds. How long will SO2 gas take to effuse from the same container? (Atomic Weight = S = 32, He = 4).

6. A system absorbs 200J of heat from the surroundings and does 120 J of work on the surroundings by expansions. Find the internal energy change of the system.

7. 1.2 gm of acetic acid (CH3COOH) is dissolved in water to make 200cm3 of the solution. Find the concentration of the solution in Molarity.

8. The solubility of calcium oxalate (CaC2O4) is 0.0016 g/dm3 at 25°C. Find the solubility product of calcium oxalate: CaC2O4 ® Ca2+ + C2O42-

9. Calculate H+ ion concentration of a solution whose pH = 5.6.

10. The rate constant (k) for the decomposition of nitrogen dioxide 2NO2(g) ® 2NO(g) + O2(g) is 1.8 x 103- dm3mole1-sec1-. Write down the rate expression and (i) find the initial rate when the initial concentration of NO2 is 0.75 M. (ii) Find the rate constant (k) when the initial concentration of NO2 is doubled.

11. Calculate the volume of nitrogen gas produced by heating 800 gm of ammonia at 21°C and 823 torr pressure. 2NH3 ® N2 + 3H2 (Atomic Weight = N = 14, H = 1)

12. In collection of 24 x 1025 molecules of C2H5OH. What is the number of moles. ( Atomic weight = C = 12, O = 16, H = 1)

13. Simplify using exponential notation: 43100 + 3900 + 2100.

14. A given compound contains 75. 2% carbon, 10.75% hydrogen and 14.05% oxygen. Calculate the empirical formula of the compound. (Atomic weight: C = 12, O = 16, H = 1)

15. Calculate the wave number of spectral line of hydrogen gas when an electron jumps from n = 4 to n= 2. (RH = 109678 cm-1)

16. 13.2 gm of gas occupies a volume of 0.918 dm3 at 25°C and 8 atm pressure. Calculate the molecular mass of the gas.

17. Calculate the heat of formation of benzene at 25°C when the heat of formation of CO2 and water and heat of combustion of benzene are given:

(i)


6C + 3H2 ® C6H6


DHf = ?

(ii)


C + O2 ® CO2


DH = -286KJ/mole

(iii)


H2 + ½O2 ® H2O


DH = -286KJ/mole

(iv)


C6H6 + 7.5O2 ® 6CO2 + 3H2O


DH = -3267 KJ/mole

18. The rate constant for the decomposition of nitrogen dioxide is 1.8 x 10-8 dm3 mole-1s-1. What is the initial rate when the initial concentration of NO2 is 0.50M? 2NO2 ® 2NO + O2.

19. Should AgCl precipitate from a solution prepared by mixing 400cm3 of 0.1M NaCl and 600cm3 of 0.03 M of solution of AgNO3? (Ksp for AgCl = 1.6 x 10-10 mole/dm3)

20. A sample of chlorine gas at S.T.P has a volume of 800cm3 calculate The number of moles of chlorine, the mass of the sample and the number of chlorine molecules in the sample.

21. How many atoms of carbon are present in 10 gm of coke?

22. The volume of the oxygen gas, collected over water at 24°C and 762mm pressure, is 128 ml. Calculate the mass in gm of oxygen gas obtained. The pressure of water vapour at 24°C is 22 mm.

23. Calculate the radius of orbit n = 3 for a Hydrogen atom in Armstrong unit. (h = 6.625 x 10-27 erg-sec, p = 3.14, m = 9.11 x 10-28gm, e = 4.8 x 10-10 esu)

24. For the reaction H2 + I2 ® 2HI. Kc is 49. Calculate the concentration of HI at equilibrium when initially one mole of H2 is mixed with one mole of I2 in one litre flask.

25. Determine the mass of HCl required to prepare 400 ml of 0.85M HCl solution.

26. Calculate pH value of 0.004M NaOH solution.

27. Kc for the reaction is 0.0194 and the calculated ratio of the concentration of the reactants and the product is 0.0116. Predict the direction of the reaction.

28. For the decomposition of ethyl chlorocarbonate ClCOOC2H5 ® CO2 + Cl.C2H5. Find the value of rate constant when initial concentration of Ethyl Chlorocarbonate is 0.25 M and the initial rate of the reaction is 3.25 x 10-4 mole/dm3/sec.

29. 1.0 gm of a sample of an organic substance was burnt in excess of oxygen yield 3.03 gm of CO2 and 1.55 gm of H2O. If the molecular mass of the compound is 58. Find the molecular formula.

30. Calculate the volume of the oxygen at S.T.P that may be obtained by complete decomposition of 51.3 gm of KClO3 on heating in presence of MnO2 as a catalyst. 2KClO3 ® 2KCl + 3O2. (Atomic mass of K = 39, Cl = 35.5, O = 16, Mn = 55)

31. Calculate the wave number of the Line in Lyman Series when an electron jumps from orbit 3 to orbit 1.

32. Calculate the heat of formation of ethane (C2H6) at 25°C from the following data:

(i)


2C + 3H2 ® C2H6


DHf = ?

(ii)


C + O2 ® CO2


DH = -394KJ/mole

(iii)


H2 + ½O2 ® H2O


DH = -286KJ/mole

(iv)


C2H6 + ½O2 ® 2CO2 + 3H2O


DH = -1560.632KJ/mole

33. At the equilibrium a 12 litre flask contains 0.21 mole of PCl5, 0.32 mole of Cl2 at 250°C. Find the value of Kc for the reaction. PCl5 Û PCl3 + Cl2.

34. A given compound contains C = 60%, H = 13.0% and O = 27%. Calculate its Empirical Formula.

35. How many grams of chlorine are required to prepare 7.75 dm3 of chloro benzene? The equation of the reaction is C6H6 + Cl2 ® C6H5Cl + HCl. (Atomic Number of C = 12, H = 1 and Cl = 35.5)

36. A mixture of helium and hydrogen is confined in a 12 dm3 flask at 30°C. If 0.2 mole of the helium is present, find out the partial pressure of each gas whereas the pressure of the mixture of gases is 2atm.

37. Calculate the radius by hydrogen atom by applying Bohr’s Theory. (h = 6.625 x 10-27 erg-sec, p = 3.14, m = 9.11 x 10-28gm, e = 4.8 x 10-10 esu)

38. Calculate the heat of formation of C2H2 from carbon and hydrogen from the following data:

(i)


2C + H2 ® C2H2


DHf = ?

(ii)


C + O2 ® CO2


DH = -94.05Kcal/mole

(iii)


H2 + ½O2 ® H2O


DH = -68.32Kcal/mole

(iv)


C2H2 + 5/2O2 ® 2CO2 + H2O


DH = -310Kcal/mole

39. Calculate the pH of a 2.356 x 10-3m HCl solution.

40. For the reaction N2 + 3H2 Û 2NH3. The equilibrium mixture contains 0.25 M nitrogen, 0.15M hydrogen gas at 25°C. Calculate the concentration of NH3 gas when Kc = 9.6. the volume of the container is 1dm3.

41. Determine the initial rate of the following reaction at 303°C in which its rate constant is 8.5 x 10-5 litre-mol-1 sec-1. Initial concentration of the reaction is 9.8 x 10-2 mole/litre. 2NO2 ® 2NO + O2.

Extra Numericals

1. 4.6gm of ethyl alcohol and 6.0gm of acetic acid kept at constant temperature until equilibrium was established. 2 gm of acid were present unused. Calculate Kc.

2. Kc for the dissociation of HI at 350°C is 0.01. If 0.2 mole of H2, 1.3 moles of I2 and 4 moles of HI are present. Predict the direction of reaction.

3. What is the solubility of PbCrO4 at 30°C when Ksp is 1.8 x 10-14.

4. 1.06m of an organic compound on combustion gave 1.49 gm of CO2 and 0.763gm H2O. It also has 23.73% N. Find its compercial formula.

5. 500 dm3 of moist O2 gas was collected over water at 27°C and 726torr pressure. Find the mass in gm. Of dry O2 gas at S.T.P. When the vapour pressure of water 27°C is 26 torr.

6. Atomic mass of phosphorus is 31. Calculate the mass of 45 atoms in a.m.u.

7. Methane burn in steam according the following reaction: CH4 + 2O2 ® CO2 + 2H2O. If 100 gm of each CH4 and O2 is taken, then what amount of CO2 liberated?

8. An organic compound containing C = 65.45%, H = 5.45% and O = 29.09%. If molecular weight of compound is 110, calculate molecular formula.

9. What mass of CO2 is produced by the complete combustion of 100g pentane. C5H12 + 3O2 ® 2CO2 + 2H2O.

10. One atom of an unknown element is found to have a mass of 67.8 x 10-23g. What is the atomic weight of the element?

11. The heat of combustion of glucose and alcohol is given below.

(i)


C6H12O6 + 6O2 ® 6CO2 + 6H2O


DH = -673Kcal/mole

(ii)


C2H5OH+ 3O2 ® 2CO2 + 3H2O


DH = -328Kcal/mole

Find DH for the fermentation given below:



C6H12O6 ® 2C2H5OH + 3CO2



12. At certain temperature, the equilibrium mixture contain 0.4 mole of H2, 0.4mole I2 and 1 mole of HI. If addition 2 mole of H2 are added. How many moles of HI will be present when the new equilibrium established. H2 + I2 ® 2HI.

13. A solution has pH of 8.4. Find concentration of H+ and OH-.

14. 180cm3 of a known gas diffuse in 15minutes, when 120 cm3 of SO2 diffuses in 20 minutes. What is the molecular mass of the unknown gas.

Chapter 1

Introduction to Fundamental Concepts

1. Calculate the moles of the following in 500gm, NH3, HCl, Na2CO3, H2SO4, MgBr2, CaCO3, Xe and C.

2. How many moles of Na are present in 5gm of Na?

3. Calculate the number of atoms in 12 gms of Mg.

4. 2gm diamond is studded in a ring. Diamond is a pure carbon. How many atoms of carbon are present in the ring?

5. Calculate the number of molecules in 9gms of H2O.

6. How many molecules are present in 25 gms of CaCO3?

7. Calculate the weight in gram of 3.01 x 1020 molecules of glucose (C6H12O6)

8. How many atoms of hydrogen are there in 2.57 x 10-6 gram of hydrogen?

9. A sample of oxygen contains 1.87 x 1027 atoms of oxygen. What would be the weight of the oxygen?

10. Find the weight of oxygen obtained from 49gm of KClO3.

2KClO3 ® 2KCl + 3O2

11. What weight of CO2 and CaO can be obtained by heating 12.5gm of Limestone (CaCO2)?

CaCO3 ® CaO + CO2

12. Calculate the weight of sodium chloride required to produce 142 gm of chlorine.

2NaCl ® 2Na + Cl2

13. Calculate the weight of carbon, required to produce 88gm of CO2.

C + O2 ® CO2

14. The action of CO on Fe2O3 can be represented by the following equation.

Fe2O3 + 3CO ® 2Fe + 3CO2

15. What weight of NH3 will be required to produce 100 gm of NO?

4NH3 + 5O2 ® 4NO + 6H2O

16. Find out the moles of CuSO4 which are obtained from 31.75 gm of Cu.

Cu + H2SO2 ® CuSO2 + H2

17. Calculate the number of N2 and H2 molecules, which are obtained from 8.5 gm of NH3.

N2 + 3H2 ® 2NH3

18. Find out the number of Cu and H2O molecules obtained from 7.95gm of CuO.

CuO + H2 ® Cu + H2O

19. 400gm of H2 was made to combine with 14200gm of Cl2. How much HCl will be produced?

20. 1kg of Limestone was heated 500gm of CaO was obtained. How much CO2 gas produced into air.

21. Find the weight of O2 obtained from 49 gm of KClO3.

2KClO3 ® 2KCl + 3O2

22. Chlorine is produced on the large scale by the electrolysis of NaCl aqueous solution. Chlorine the weight of NaCl required to produce 142 gm of Cl2.

2NaCl + 2H2O ® Cl2 + H2 + 2NaOH

23. How many grams of O2 are required to completely burn 18.0gm of C? How many grams of CO2 will be formed?

24. Calculate the weight of NH3, required to produce 100 gms of NO.

4NH3 + 5O2 ® 4NO + 6H2O

25. Find out the moles of H2 and N2 required producing 17gm of NH3.

26. Calculate the volume of H2 at S.T.P, which is obtained by the reaction of 120 gm Mg with MgSO4.

Mg + H2SO4 ® MgSO2 + H2

27. NH3 gas can be produced from ammonium chloride (NH4Cl) as follows:

CaO + 2NH4Cl ® CaCl2 + H2O + NH3

Calculate the volume of NH3 obtained at S.T.P by the reaction of 100 gm of NH4Cl.

28. 500gm of C2H4 on combustion in air gave CO2 and H2O. Calculate the volume of O2 and CO2 at S.T.P.

29. Find out the volume of O2, CO2 and SO2 gases at S.T.P react and obtained from 2 moles of CS2.

CS2 + 3O2 ® CO2 + 2SO2

30. Calculate the volume of CO2 gas at S.T.P obtained by the combustion of 20gm of CH4.

CH4 + 2O2 ® CO2 + 2H2O

31. Calculate the volume of O2 gas at S.T.P required to burn 600dm3 of H2S, also find the volume of SO2 gas produced at S.T.P.

32. Calculate the volume of O2 gas at S.T.P required to burn 50 gm of CH4.

33. What volume of H2 at S.T.P can be produced by the reaction of 6.54gm Zn with HCl?

Zn + 2HCl ® ZnCl2 + 2H2

34. Calculate the volume of O2 and H2 gases at S.T.P obtained from 9gm of H2O.

35. 0.264gm of Mg was burnt in pure O2. How much MgO will be formed?

2Mg + O2 ® 2MgO

36. How much H2 can be generated by passing 200gm of steam over hot iron.

4H2O + 3Fe ® Fe3O4 + 4H2

37. If 112dm3 of N2 react with 336 dm3 of H2, both at S.T.P. How many grams of NH3 would be obtained?

N2 + 3H2 ® 2NH3

38. An organic compound contains 12.8%C, 2.1% and 85.1% Br. If the mass of the compound is 188, find the molecular formula.

39. An organic compound contains 66.70%C, 7.41% H and 25.90% N2. The molecular mass of the compound is 108. Find out its molecular formula.

40. A compound contains 19.8%C, 2.5%H, 66.1%O and 11.6%N. Find out empirical formula of the compound.

41. 0.2475gm of a compound, containing C, H and O gave 0.4950gm CO2 and 0.2025gm H2O. If the molecular mass of the compound is 88. Find out the molecular formula.

42. An organic compound contains 32%C, 6.67%H, 18.66%N and 42.67%O. Its molecular mass is 75. Find out the molecular formula of the compound.

43. 1.367gm of a compound containing C, H and O on heating gave 3.002gm CO2 and 1.640gm H2O. Find out its molecular formula, when the molecular mass is 120.

44. A compound was found to contain 40%C and 6.7%H. Its molecular mass was 60. Find out its molecular formula.

45. An organic compound contains 75.2%C, 10.15%H and oxygen. Its molecular mass is 115. Find its molecular formula.

46. The empirical formula of a compound is CH2O. If the molecular mass 180. Find out the molecular formula.

47. An organic compound composed of C, H and O. On combustion of 0.94gm of this compound, 1.32gm CO2 and 0.568gm H2O were obtained. Its molecular mass is 180. Find its molecular formula.

48. An organic compound composed of C, H and O. 4.2gm of the compound on heating gave 6.21gm CO2 and 2.54gmH2O. Its molecular mass is 60. Find its molecular formula.

49. An organic compound contains C,H and 6.38gm of compound on combustion gave 9.06gm CO2 and 5.58gm H2O. Its molecular mass is 62. Find out its molecular formula.

50. 1gm of a hydrocarbon on combustion gave 3.03gm of CO2 and 1.55gm of H2O. If the molecular mass is 58, find its molecular formula.

51. 1.434gm of a compound on combustion gave 4.444gm CO2 and 2.0 gm H2O. Find out its empirical formula.

52. An organic compound composed of C, H and N. 0.225gm of compound on combustion gave 0.44gm CO2 and 0.315gm H2O. If the molecular mass of a compound is 90, find out its molecular formula.

53. An organic compound contains 40.68%C, 8.47%H, 23.73%N and 27.12%O. Find its empirical formula.

54. An organic compound composed of C, H and N. 0.419 gm of compound on combustion gave 0.88gm CO2 and 0.27gm H2O. Find out its empirical formula.

55. The analysis of a compound shows, C = 24.24%, H = 4.04% and Cl = 71.71%. If the molecular mass of the compound is 49.5, find its molecular formula.

56. An organic compound of molecular mass 90 has the empirical formula CH2O. What is its molecular formula?

57. The empirical formula of an organic compound is CH3NO2. If it’s molecular mass is 61. What is its molecular formula?

58. 0.638gm of an organic compound on combustion gave 0.594gm H2O and 1.452gm CO2.The compound is composed of C, H and O atoms. If the molecular mass is 116, find out its molecular formula.

59. The molecular formula of ethyl acetate is CH3COOC2H5. What is its empirical formula.

60. Find the empirical formulae of the following compounds from their percentage composition by mass:

· N = 26.17% H = 7.48% Cl = 66.35%

· Ca = 71.43% O = 28.57%

· Ag = 63.53% N = 8.23% O = 28.24%

· Na = 32.40% H = 45.07% Cl = 22.53%

61. A certain compound on analysis yielded 2.00gm C, 0.34gm H and 2.67gm O. If the relative molecular mass of the compound is 60, calculate its molecular formula.

62. What is the empirical formula of a compound, which contains 42.5% chlorine and 57.5 oxygen. If it’s formula mass is 167. What is its molecular formula?

63. What will be the weight of 5 moles of water in grams?

64. What is the mass of each of the following:

· 1.25 mole of NaCl

· 2.42 mole of NaNO3

· 1.5 mole of HCl

· 3.0 mole of NaOH

65. A piece of Aluminium metal weighs 70.0g. How many atoms are present in the piece.

66. How many atoms of carbon are present in 20-carat Diamond? (1 carat = 0.2g)

67. How many grams of oxygen have the same number of atoms as 16gm of sulphur?

68. A sample of oxygen gas at STP has a mass of 16gm. Calculate:

· The number of moles of oxygen

· The volume of the sample

· The number o molecules in the sample

69. Calculate the volume of CH4 gas at STP having a mass 32g.

70. What mass of zinc sulphate can be obtained from the reaction of 10.0gm of Zinc with an excess of dilute H2SO4?

Zn + H2SO4 ® ZnSO4 + H2*

71. Calculate what mass of sodium hydroxide you would need to neutralize a solution containing 7.3g hydrogen chloride by the reaction:

NaOH + HCl ® NaCl + H2O

72. Calculate how much sodium nitrate you need to give 126g of nitric acid by the reaction:

NaNO3 + H2SO4 ® HNO3 + NaHSO4

73. What volume of hydrogen at STP is evolved when 0.325g of zinc reacts will dilute hydrochloric acid.

Zn + 2HCl ® ZnCl2 + H2

74. What mass of oxygen is formed by the decomposition of a solution containing 120cm3 of H2O2 at STP?

2H2O2 ® 2H2O + O2

75. What is the mass of one molecule of water in grams?

76. 100cm3 of butane are burned in an excess of oxygen. Calculate:

· The volume of oxygen used

· The mass and volume of CO2 produced (assume all gases at STP)

2C4H10 + 13O2 ® 8CO2 + 10H2O

77. A cook is making a small cake. It needs 500cm3 at STP of CO2 to make the cake rise. The cook decides to add baking powder, which contains sodium bicarbonate. This generates CO2 by thermal decomposition.

2NaHCO3 ® CO2 + Na2CO3 + H2O

What mass of baking powder must the cook add to cake mixture?

78. What volume of ammonia at STP can be obtained by heating 0.25 mole of ammonium sulphate with calcium hydroxide?

(NH4)SO4 + Ca(OH)2 ® 2NH3 + CaSO4 + 2H2O

79. How many grams of SO2 are produced when 100g of H2S is reacted with 50g of oxygen.

2H2S + 3O2 ® 2H2O + 2SO2

80. How many grams of chlorobenzene will be produced when 100gm of each reactant is reacted?

C6H6 + Cl2 ® C6H5Cl + HCl

81. A car releases about 5g of NO into the air for each mile driven. How many molecules of NO are emitted per mile?

82. Simplify according to the rule of significant figures.

· 2.60 x 3.05

· 0.009 ¸ 0.3

·

·

Chapter 2

The Three States of Matter

1. 540cm3 of N2 at 400mm pressure are compressed to 300cm3 without changing the temperature. What will be the pressure of the gas?

2. A gas occupies 6dm3 at 1atm pressure keeping the temperature constant. If the pressure reduces to 600mm, what volume does the gas occupy?

3. At a certain temperature and 800mm pressure, the volume of H2 is 700cm3. If the pressure is increased to 1000mm at the same temperature, find the new volume of the gas.

4. 150ml of a gas at 27°C is heated to 77°C at constant pressure. Find the new volume of the gas.

5. 300ml of N2 are at 50° and the pressure is kept constant. If the temperature is doubled, what will be the volume of the gas?

6. A gas measures 5dm3 at 5°C under 0.5atm pressure. Calculate its volume at 25° and 5000mm pressure.

7. 2060ml of a gas is at 7°C and 860mm pressure. Find its volume at S.T,P.

8. 350ml of H2 was collected over water at 26°C. The pressure of the gas was 900mm. What volume will dry gas have at 30°C and 750mm pressure? The vapour pressure at 26°C is 25mm.

9. The volume of oxygen collected over water at 20°C and 1200mm pressure, is 200cm3. If aqueous at 20°C is 17.4mm, what will be the volume of the gas under S.T.P.

10. A 20dm3 flask contains H2 at 22°C under pressure of 1.2 atm. How many moles of H2 are present.

11. A gaseous mixture is at the pressure of 3000mm. The mixture contains 6 moles of N2, 0.5mole of CO2 and 2.5 moles of O2. Find the partial pressure of each gas.

12. A 5dm3 vessel contains 1.2 moles of H2 and 0.8 mole of N2 at 27°C. Find the total pressure of the mixture.

13. Composition of a sample of air by volume is, N2 = 76%, O2 = 20%, H2O = 2.5%, CO2 = 1.4% and He = 0.1%. If the pressure of the air is 760 mm, Calculate the partial pressure of these gases.

14. A 10dm3 container contains a mixture of He and Ne gases at 17°C. There are two moles of He gas and 3 moles of Ne gas. What is the partial pressure of the gases?

15. 10gm of H2, 96gm of O2 and 196gm of N2 are mixed together. The partial pressure of H2 is 0.6 atm. What is the partial pressure of O2 and N2?

16. A cylinder contains 1 mole of H2, 3 mole of He and 6 moles of N2. The total pressure in the cylinder is 15 atm. Calculate the partial pressure of H2, He and N2.

Chapter 5

Energetics of Chemical Reaction

1. Calculate the heat of formation of Acetic Acid from the following data:

(i)


2C + 2H2+ O2 ® CH3COOH


DHf = ?

(ii)


C + O2 ® CO2


DH = -394KJ/mole

(iii)


H2 + ½O2 ® H2O


DH = -286 KJ/mole

(iv)


CH3COOH + 2O2 ® 2CO2 + 2H2O


DH = -870KJ/mole

2. Calculate the heat of formation of Ethane from the following data:

(i)


2C + 3H2 ® C2H6


DHf = ?

(ii)


C + O2 ® CO2


DH = -394KJ/mole

(iii)


H2 + ½O2 ® H2O


DH = -286 KJ/mole

(iv)


C2H6 + 7/2O2 ® 2CO2 + 3H2O


DH = -1560KJ/mole

(v)


C2H5OH + 3O2 ® 2CO2 + 3H2O


DH = -327 KJ/mole

3. Calculate the heat of formation of Methane from the following data:

(i)


C + 2H2 ® CH4


DHf = ?

(ii)


C + O2 ® CO2


DH = -394KJ/mole

(iii)


H2 + ½O2 ® H2O


DH = -286 KJ/mole

(iv)


CH4 + 2O2 ® CO2 + 2H2O


DH = -890.3KJ/mole

4. Calculate the heat of formation of Ethyl Alcohol from the following data:

(i)


2C + 3H2 ½ O2® C2H5OH


DHf = ?

(ii)


C + O2 ® CO2


DH = -394KJ/mole

(iii)


H2 + ½O2 ® H2O


DH = -286 KJ/mole

(iv)


C2H5OH+ 3O2 ® 2CO2 + 3H2O


DH = -1369KJ/mole

5. Calculate the heat of formation of Ethane from the following data:

(i)


C2H6 + 7/2O2 ® 2CO2 + 3H2O


DHf = ?

(ii)


C + O2 ® CO2


DH = -394KJ/mole

(iii)


H2 + ½O2 ® H2O


DH = -286 KJ/mole

(iv)


C2H6 ® 2C + 3H2


DH = -84.68KJ/mole

6. Calculate the heat of formation of Methane from the following data:

(i)


C + 2H2 ® CH4


DHf = ?

(ii)


C + O2 ® CO2


DH = -94.1cal

(iii)


H2 + ½O2 ® H2O


DH = -68.3 cal

(iv)


CH4 + 2O2 ® CO2 + 2H2O


DH = -212.8 cal

7. Calculate the heat of formation of Ethene from the following data:

(i)


2C + 2H2 ® C2H4


DHf = ?

(ii)


C + O2 ® CO2


DH = -97kcal

(iii)


H2 + ½O2 ® H2O


DH = -65 kcal

(iv)


C2H4 + 3O2® 2CO2 + 2H2O


DH = 340 kcal

8. Calculate the heat of formation from the following data:

(i)


2C + 3H2 +1/2O2 ® C2H5O


DHf = ?

(ii)


C + O2 ® CO2


DH = -94.2Kcal/mole

(iii)


H2 + ½O2 ® H2O


DH = -68.5 Kcal/mole

9. Calculate the heat of formation of from the following data:

(i)


C + 2H2 + O2® CH3OH


DHf = ?

(ii)


C + O2 ® CO2


DH = -94.2Kcal/mole

(iii)


H2 + ½O2 ® H2O


DH = -68.32 Kcal/mole

(iv)


CH3OH + O2 ® CO2 + 2H2O


DH = -347.6Kcal/mole

10. Calculate the heat of formation of from the following data:

(i)


3C + 4H2 ® C3H8


DHf = ?

(ii)


C + O2 ® CO2


DH = -94.1Kcal/mole

(iii)


H2 + ½O2 ® H2O


DH = -68.3 Kcal/mole

(iv)


C3H8 + 5O2 ® 3CO2 + 4H2O


DH = -530.7Kcal/mole

11. Calculate the heat of formation of from the following data:

(i)


H2 + O2® H2O2


DHf = ?

(ii)


H2 + ½O2 ® H2O


DH = -68.32Kcal

(iii)


H2O + ½ O2 ® H2O2


DH = -23.48Kcal

12. Given:

(i)


NH3 + HCl ® NH4Cl


DH1 = 42.100Kcal

(ii)


H2O + ½ O2 ® H2O2


DH2 = 3.900cal

Find DH for the reaction,



NH3 + HCl ® NH4Cl


DHf = ?

Chapter 6

Chemical Equilibrium

1. 1.5 moles of acetic acid and 1.5 moles of ethyl alcohol were reacted at a certain temperature. At equilibrium, 1 mole of ethyl acetate was present in 1 litre of the equilibrium mixture. Calculate the equilibrium constant Kc.

CH3COOH + C2H5OH Û CH3COOC2H5 + H2O

2. 6.0 gm of hydrogen and 1016gm of iodine were heated in a sealed tube at a temperature, at which Kc is 50. The volume of the tube is 1 dm3. Calculate the concentration of HI.

H2 + I2 Û 2HI

3. At a certain temperature, an equilibrium mixture contains 0.4 mole H2, 0.4 mole I2 and 1 mole of HI. The volume of the reacting vessel is 4 dm3. Find out the equilibrium constant kc.

H2 + I2 Û 2HI

4. 3 moles of A and 2 moles of B are mixed in a 4dm3 flask, at a certain temperature. The following reaction occurs.

3A + 2B Û 4C

At equilibrium the flask contains 1 mole of B. Find the equilibrium constant kc.

5. At a certain temperature, 0.205 mole of H2 and 0.319 mole of I2 were reacted. The equilibrium mixture contains 0.314 mole of I2. Calculate the kc.

H2 + I2 Û 2HI

6. The kc for the reaction A + B Û C + D is 1/3. How many moles of A must be mixed with 3 moles of B to yield at equilibrium, 2 moles of C and D each. The volume of the vessel is 2 litre.

7. At a certain temperature the equilibrium mixture for the reaction A + B Û 2C, contains 2 moles A, 3 moles of B and 5 moles of C. Find the Kc for the reaction.

8. For the reaction 2A Û B + C, equilibrium constant kc is 1. If we start with 6 moles of A, how many moles of B will be formed.

9. 20 moles of SO2 and 10 moles of O2 are taken in a 20 litre flask. If at equilibrium 5 moles of SO3 are formed, Calculate kc.

2SO2 + O2 Û 2SO3

10. A quantity of PCl5 was heated in a 12 dm3 vessel at 250°C.

PCl5 Û PCl3 + Cl2

11. 2 moles of HI was introduced in a vessel held at constant temperature. When equilibrium was reached, it was found that 0.1 mole of I2 have been formed. Calculate the equilibrium constant.

H2 + I2 Û 2HI

12. When 1 mole of pure C2H5OH is mixed with 1 mole of CH3COOH at room temperature, the equilibrium mixture contains 2/3 moles of ester and water each.

· What will be the kc?

· How many moles of ester are formed at equilibrium when 3 moles of C2H5OH are mixed with 1 mole of CH3COOH?

CH3COOH + C2H5OH Û CH3COOC2H5 + H2O

13. PCl5 Û PCl3 + Cl2. Calculate the number of moles of Cl2 produced at equilibrium when 1 mole of PCl5 is heated at 250°C in a vessel having capacity of 10dm3. At 250°C, Kc is 0.041.

14. When 2.94 moles of iodine and 8.1 moles of Hydrogen were mixed and heated at 444°C and at constant volume, until the equilibrium was established. 5.64 moles of HI were formed. Calculate the value of kc.

H2 + I2 Û 2HI

15. What is the solubility of lead chromate in moles/dm3 at 25°C. The solubility product is 1.8 x 10-14.

PbCrO4 Û Pb++ + CrO4--

16. The solubility of Mg(OH)2 at 25°C is 0.00764 gm/dm3. What is the solubility product of Mg(OH)2?

Mg(OH)2 Û Mg++ + 2OH-

17. Find the solubility of AgCl in gm/dm3, when the solubility product is 1.25 x 10-10.

18. Calculate the solubility product of BaSO4. The solubility of the salt is 1.0 x 10-5 moles/dm3.

19. Calculate the solubility product of BaSO4 is 9.0 x 10-3 gm/dm3. Find its solubility product.

20. Predict whether there will be any precipitate formation by mixing 30cm3 of 0.01M NaCl with 60cm3 of 0.01M AgNO3 solution. Ksp of AgCl is 1.5 x 10-10.

21. A saturated solution of calcium fluoride was found to contain 0.0168 gm/dm3 of solute at 25°C. Calculate the ksp for CaF2.

22. A saturated solution of BaF2 at 25°C is 0.006M. Calculate Ksp of the salt.


Fill in the Blank

 1. The property of a crystal, which is different in different directions, is called __________.

2. 0.00051 contains __________ significant figures.

3. The oxidation number of oxygen in OF2 is __________.

4. The volume of 1 gm of hydrogen gas at S.T.P is __________.

5. The oxidation number Mn in KMnO4 is __________.

6. The product of ionic concentration in a saturated solution is called __________.

7. 16 gm of oxygen at S.T.P occupies a volume of __________ dm3.

8. The shape of the orbital for which l = 0 is __________.

9. The radius of Cl-1 is __________ than the radius of Cl0.

10. Sp2 hybridization is also known as __________.

11. The value of 1 Debye is __________.

12. The reactions catalyzed by sunlight are called __________.

13. The blue colour of CuSO4 is due to the presence of __________.

14. The force of attraction between the liquid molecules and the surface of container is called __________.

15. The heat of neutralization of a strong acid and a strong base is __________.

16. C º C triple bond is __________. C = C double bond length.

17. The ions having the same electronic configuration are called iso electronic.

18. On heating, if a solid changes directly into vapours without changing into the liquid state, the phenomenon is called __________.

19. Each orbital in an atom can be completely described by __________.

20. In a molecule of alkene, __________ restricts the rotation of the group of atoms at either end of the molecule.

21. Density, refractive index and vapour pressure are __________ properties.

22. The addition of HCl to H2 solution __________ the ionization of H2S.

23. The reaction of cation or anion (or both) with water so as to change its __________ is known as Hydrolysis.

24. A reaction with higher activation energy will start at __________ temperature.

25. 6.02 x 1023 has __________ significant figures.

26. The internal resistance in the flow of liquid is called __________.

27. A catalyst increases the velocity of a reaction but decreases the __________.

Chapter 1

Introduction to Fundamental Concepts

1. 1 mole of a gas at S.T.P occupies a volume of __________.

2. A gas occupying a volume of 22.4 dm3 at S.T.P contains __________ molecules.

3. A formula, which gives the relative number of atoms in the molecule of a compound, is called __________.

4. A formula which gives the actual number of all kinds of atoms present in the molecule of compound is termed as __________.

5. The chemical formula that not only gives the actual number of atoms but also shows the arrangement of different atoms present in the molecule is called __________.

6. Atomic weight or molecular weight expressed in grams is known as __________.

7. 2 moles of H2O contain __________ grams and __________ number of molecules.

8. Any thing that occupies space and has __________ is called matter.

9. Volume of one __________ mole of a gas at S.T.P is 22.4 cubic feet.

10. A ton mole of iron is equal to __________ tons.

11. The force with which the earth attracts a body is called the __________ of the body.

12. A pure substance contains __________ kind of molecules.

13. The smallest indivisible particle of matter is called __________.

14. The atomic number is equal to the number of __________ in nucleus.

15. The atomic mass is the total number of protons and __________ in an atom of the element.

16. The average weight of atoms of an element as compared to the weight of one atom of __________ is called the atomic mass.

17. 1.0007 contains __________ significant figures.

18. The figure 24.75 will be rounded off to __________.

19. __________ means that the readings and measurements obtained in different experiments are very close to each other.

20. __________ means that the results obtained in different experiments are very close to the accepted values.

21. The degree of a measured quantity __________ with increasing number of significant figures in it.

22. The atomic mass of sodium is __________.

23. The symbolic representation of a molecule of a compound is called __________.

24. Molecular formula of CHCl3 and its Empirical formula is __________.

25. Molecular formula of benzene is C6H6 and its empirical formula is __________.

26. 58.5 is the __________ of NaCl.

27. 4.5 gms of nitrogen will have __________ molecules.

28. 28 gms of nitrogen will have __________ molecules.

29. 2 moles of SO2 is equal to __________ gms.

30. 1000 gms of H2O is equal to __________ moles.

31. The reactions, which proceed in both directions, are called __________.

32. The reactions, which proceed in forward directions only, are called __________ reactions.

33. The __________ reactions are completed after some time.

34. 0.0006 has __________ significant figures

35. 7.40 x 108 has __________ significant figures.

36. 7 x 108 has __________ significant figures.

37. Usually Molecular formula is simple multiple of the __________.

38. 0.1 mole of H2O contains __________ molecules of H2O.

39. Mass of 3.01 x 1022 molecules of CO2 is __________.

40. __________ is the branch of science which deals with the properties, composition and structure of matter.

41. None zero digits are all __________.

42. The integer part of logarithm is called __________.

43. The decimal fraction of logarithm is called __________.

44. __________ is the amount of substance, which contains as many number of particles as there are in 12 gms of Carbon.

45. 6.02 x 1023 is called the __________.

46. The accuracy of measurement depends on the number of __________.

47. __________ is the branch of chemistry that deals with quantitative relationships among the substances undergoing chemical changes.

48. The sum of atomic weights of all the elements present in molecular formula is called the __________.

49. __________ is the sum of atomic weights of the elements represented by the Empirical formula of the compound.

50. Very small and very large quantities are expressed in terms of __________.

51. In rounding off __________ figure is dropped.

52. Mole is the quantity, which has __________ particle of the substance.

53. For three significant figures, 25.55 is rounded off to __________.

54. The S.I unit of a mass is __________.

55. Mass of 6.02 x 1023 molecules of NaCl is __________ gm.

56. 1 mole of NaOH is __________ gm of NaOH.

57. Formula weight is used for __________ substances.

58. The word S.I stands for __________.

59. 4.5 gms of water will have __________ molecules.

60. 0.0087 has __________ significant figure.

Chapter 2

The Three States of Matter

1. The intermixing of gases or liquids in a container irrespective of their densities, is called __________.

2. At constant temperature, if the pressure of a given mass of a gas is decreased, its volume will __________.

3. A volume of __________ dm3 will hold 128 gms of SO2.

4. At constant temperature of a given mass of a gas, the product of its __________ and __________ is constant.

5. The rates of diffusion of gases are __________ proportional to the square root of their densities.

6. Gases deviate from ideal behaviour more markedly at high __________.

7. Liquid diffuse __________ than gases.

8. An imaginary line passing through the centre of a crystal is called __________.

9. The temperature at which more than one crystalline forms of a substance coexist in equilibrium is called __________.

10. Two or more substances crystallizing in the same form is called __________.

11. The existence of solid substances in more than one crystalline form is known as __________.

12. Rate of diffusion of gases is __________ as compared to liquids.

13. Boiling point of a liquid __________ with the pressure.

14. Mercury in a glass tube forms __________ curvature.

15. Gases can be compressed to __________ extent.

16. Viscosity of a liquid __________ with the increase of temperature.

17. Surface tension of water __________ by adding soap solution into it.

18. The internal resistance to the flow of a liquid is called __________.

19. The rise or the fall of a liquid in a capillary tube is called __________.

20. Matter exists in __________ states.

21. The freezing point of water in Fahrenheit scale is __________.

22. Boiling point of water is __________ °K.

23. SI unit for measurement of pressure is __________.

24. The value of gas law constant R = __________ dm3 atm/°K/mole.

25. The absolute Zero is equal to __________.

26. If P is plotted against 1/V at constant temperature a __________ is obtained.

27. Gases __________ in heating.

28. The pressure of air __________ at higher altitude.

29. Standard temperature means __________.

30. Standard pressure means __________.

31. Cooling is caused by __________ of gases.

32. Rate of diffusion of O2 is __________ times more than H2.

33. H2O has __________ viscosity than CH3OH.

34. Mercury does not wet the glass surface due to its higher __________.

35. Surface tension of mercury is __________ than water.

36. Viscosity can be easily measured by an instrument called __________.

37. The pressure exerted by the vapours when these vapours are in equilibrium with the liquid is called __________.

38. Vapour pressure __________ at high temperature.

39. Boyle’s Law and Charles Law can be combined into the mathematical expression __________.

40. Equal volumes of all gases at the same temperature and pressure contain __________ number of molecules.

41. The average Kinetic energy of a gas is proportional to its __________ temperature.

42. Kinetic equation may be mathematically written as __________.

43. The temperature at which two crystalline forms of a substance can coexist in equilibrium is called __________.

44. Lighter gases diffuse __________ than heavier gases.

45. Rain drops are __________ in shape.

46. Due to surface tension, the surface area of the liquid is __________.

47. Water __________ in the capillary tube.

48. Viscosity of a solution at 10°C is __________ than at 20°C.

49. Shape of NaCl crystal is __________.

50. Gases intermix to form a mixture.

51. Pressure of a dry gas is __________ than the pressure of a moist gas.

52. 22.4 dm3 of nitrogen at S.T.P will weigh equal to __________ gm.

53. 1 mole of any gas at S.T.P is equal to __________ dm3.

54. At -273°C, volume of all gases becomes __________.

55. The gases, which strictly follows the gas Laws are called __________ gas.

56. __________ is the property that determines the direction of flow of heat.

57. __________ is defined as force per unit area.

58. __________ viscosity is defined as the viscosity of a liquid as compared to the water.

Chapter 3

Structure of Atom

1. The maximum number of electrons in 2p orbital is __________.

2. 3d orbital has __________ energy than 4s orbital.

3. __________ rays are non-material in nature.

4. Charge to mass ratio of cathode rays resembles to that of __________.

5. __________ rays are most penetrating.

6. Neutrons have mass equal to that of __________.

7. Energy is __________ when an electron jumps from higher to lower orbit.

8. Second Ionization Potential has __________ value than the First Ionization Potential.

9. Electronegativity __________ from left to right in a period of Periodic Table.

10. __________ was discovered during the course of Artificial Radioactivity.

11. The velocity of alpha rays is nearly __________ of velocity of light.

12. Natural Radioactivity is confined in __________ elements.

13. The isotopes of an element differ in their __________.

14. Two electrons with the __________ spin, can never occupy the same atomic orbital.

15. ‘Al’ has electronic configuration, 1s2, 2s2, __________.

16. In a group of Periodic Table, the ionization potential __________ from top to bottom as the size of atom increases.

17. Ionization potential values __________ from left to right in a period.

18. The energy required to remove the most loosely bond electron from an atom in gaseous state is called __________.

19. The SI unit of Ionization Potential is __________.

20. An atom of sodium possesses 11 protons and __________ neutrons.

21. The particles of Cathode rays possess __________ charge.

22. The negatively charged particles found in Cathode rays are named as __________.

23. Positive rays are emitted from __________.

24. __________ rays are also known as Canal rays.

25. __________ consists of helium ions and are doubly positively charged.

26. __________ rays consists of negatively charged particles.

27. __________ rays are light waves of very short wavelength.

28. The phenomenon in which a stable element is made radioactive by artificial disintegration is called __________.

29. The electron move around the nucleus in different circular paths called __________.

30. The maximum number of electron in a shell is determined by the formula __________.

31. A particle whose mass is equal to that of electron but carries a positive charge is called __________.

32. 2p electrons are __________ in energy that 2s electrons in the same atom.

33. Number of protons of an element also indicates its __________.

34. According to __________ Principle electrons are fed in the order of increasing orbital energy.

35. According to __________ electrons are distributed among the orbitals of a sub shell to give maximum number of unpaired electron and have same spin.

36. The specific way in which the orbitals of an atom are occupied by electrons is called __________.

37. __________ rays are stream of doubly positively charged particles.

38. Electron in the outer most shell of an atom is called __________.

39. Protons are found in the __________ of an atom and bear __________ charge.

40. The atomic number of an atom is the sum of __________ inside the nucleus.

41. __________ limits the number of electron to different shell or orbits.

42. Sir William Crookes in 1878, discovered that the cathode in high vacuum tube emit radiations what he called __________.

43. X-rays were discovered in 1895 by __________.

44. The discovery of proton was done in 1886 by __________.

45. Neutrons were discovered by __________ in 1932 by the bombardment of beryllium with alpha particles.

46. Each atom has a __________, which contains all the positive charge and practically all the mass of atom.

47. Complete the reaction: 4Be9 + 2H4 ® __________ + __________.

48. __________ have higher ionization power as compared to b-rays.

49. No dark spaces between the colours are present in __________.

50. The symbol e+ represents __________.

51. p-orbitals are __________ shaped.

52. The energy released when an electron is added to an atom in the gaseous state is called __________.

53. The power of an atom to attract a shared pair of electrons towards itself is called __________.

54. Fluorine is __________ electronegative than chlorine.

55. Lyman series of spherical lines appear in the __________ portion of spectrum.

56. The electrons with __________ spin occupy the same orbital.

57. 3d orbital has __________ energy than 4s orbital.

58. Energy and frequency are __________ proportional to each other.

59. Ionic radii of cations are __________ than the atoms from which they are formed.

60. Ionic radii of anions are __________ than the atoms from which they are formed.

Chapter 4

Chemical Bonding

1. A bond formed due to transference of electron is called __________.

2. A bond formed due to sharing of electron is called __________.

3. Sigma bond is __________ than pi bond.

4. The shape of methane molecule is __________.

5. One s and 3p orbitals overlap to produce four __________ hybrid orbitals.

6. Ethene, C2H4 is an example of __________ hybridization.

7. Water molecule has __________ structure.

8. Water molecules are inter-linked with one another due to __________.

9. Polarity of the molecule is due to the difference of __________ between the two bonded atoms.

10. A chemical bond formed between to different atoms by mutual sharing of electron is termed as __________.

11. A chemical bond formed between two similar atoms by mutual sharing of electrons is known as __________.

12. The difference between the Electronegativity values of the two atoms forming covalent bond must be __________ than 1.7.

13. When two orbitals of different atoms by hybridize with each other having their axes in the same straight lines, the bond formed is termed as __________.

14. __________ bond is formed when p-orbitals of the two atoms with their axes parallel to each other overlap with each other.

15. Melting and boiling point of ionic compounds are usually __________ than that of covalent compounds.

16. Non polar compounds are usually __________ in non polar solvent.

17. The nitrogen in NH3 is __________ hybridized.

18. A hybrid orbital is called __________ orbital.

19. Since dipole moment of CS2 is zero, it is a __________ molecule.

20. A bond formed due to the electrostatic forces of attraction between the oppositely charged ions is called __________ bond.

21. The ionic bond is formed between the atoms with low ionization potential and high __________.

22. A bond formed by the sharing of an electron pair contributed by one atom only is called a __________ bond.

23. A co-ordinate covalent bond is also known as __________ bond.

24. Polar covalent bond is __________ than a non polar covalent bond.

25. H-F bond is __________ than H-Br bond.

26. The SI unit of dipole moment is __________.

27. Commonly used unit of dipole moment is __________.

28. Dipole moment of non-polar compound is __________ D.

29. The reactions of ionic compounds are usually very __________.

30. Covalent compounds are generally __________ in nature.

31. Ionic compounds are generally __________ in nature.

32. A covalent bond is represented by a __________.

33. A co-ordinate covalent bond is represented by an __________.

34. The covalent bond between H-F is called __________ covalent bond.

35. The power of an atom to attract a shared pair of electron itself is called __________ of that atom.

36. m = d x e represents __________.

37. CO2 and SO2 molecules have __________ polar bonds.

38. NH3 molecule has __________ polar bonds.

39. A double bond has __________ bond energy than a single bond.

40. An orbital which surrounds a single nucleus is called __________ orbital.

41. An orbital which surrounds two or more atomic nuclei is called __________ orbital.

42. A molecular orbital, which is of lower energy than the atomic orbitals from which it is derived, is known as __________ orbital.

43. A molecular orbital, which has higher energy than the atomic orbitals from which it is derived, is known as __________ orbital.

44. Orbitals formed after hybridization are called __________ orbitals.

45. Bond angle in Sp3 hybridization is of __________.

46. Bond angle in Sp2 hybridization is of __________.

47. Bond angle in Sp hybridization is of __________.

48. Sp3 hybridization is also known as __________.

49. Sp2 hybridization is also known as __________.

50. Sp hybridization is also known as __________.

51. A pair of electrons residing on the central atom and which is not used in bonding is called a __________.

52. The sum of total number of electron pairs (bonding and lone pairs) is called __________ number.

53. __________ bond is usually expressed by dotted line.

54. Water molecule has dipole moment because of its __________ structure.

55. CO2 is non polar because of its __________ structure.

56. Overlapping in __________ bond is perfect.

57. Overlapping in __________ bond is not perfect.

58. H-H bond is __________ than H-Cl bond.

59. __________ hybrid orbitals are not co-planar.

60. Covalent bond is Cl2 molecule is __________.

Chapter 5

Energetics of Chemical Reaction

1. The branch of Chemistry, which deals with the heat changes that take place during chemical reaction, is called __________.

2. The branch of science which deals with energy changes accompanying physical and chemical transformation is called __________.

3. The amount of heat evolved or absorbed in a chemical reaction is called __________.

4. Such reactions in which heat is evolved are called __________ reactions.

5. Such reactions in which heat is absorbed are called __________ reactions.

6. In exothermic reactions, heat evolved is given by __________ sign of DH.

7. In endothermic reactions heat absorbed is given by __________ sign of DH.

8. The total heat change in a reaction is the same whether it takes place in one or several steps.

9. The first law of thermodynamics is also known as __________.

10. The part of universe under observation is called __________.

11. The system plus its surrounding is called __________.

12. Such properties, which give description of a system at a particular moment, is called __________.

13. The term E + PV is called __________.

14. DH represents change in __________.

15. The temperature of water is raised up when sulphuric acid is added to it. This is an __________ reaction.

16. The characteristic properties of a system which is independent of amount of material concerned is called __________ properties.

17. The characteristic properties of a system which depend on amount of substance present in it is called __________ properties.

18. Density, pressure and temperature are the examples of __________ properties.

19. Mole numbers and enthalpy are the examples of __________ properties.

20. A system, which exchange both energy and matter with its surrounding, is called __________ system.

21. A system, which only exchange energy with the surrounding but not matter is, called __________ system.

22. A system which neither exchange energy nor matter with its surrounding is called __________ system.

23. A system is __________ if it contains only one phase.

24. A system is __________ if it contains more than one phase.

25. 1 kilojoule is equal to __________ joules.

26. 1 Calorie is equal to __________ joules.

27. 1 kilo calorie is equal to __________ joules.

28. The work done (w) is mathematically denoted by __________.

29. The change in enthalpy is denoted by __________.

30. __________ law is used in calculating heat of reaction.

31. __________ is defined as the change in enthalpy when one gram mole of a compound is produced from its element.

32. Heat of formation is denoted by __________.

33. When the work is done on the system by the surrounding the sign of work done (w) is __________.

34. When the work is done by the system on surrounding the sign of work done is __________.

35. First law of Thermodynamics is mathematically represented as __________.

36. Standard enthalpies are measured at __________.

37. Hess’s Law is employed to calculate __________ of a chemical reaction.

38. Heat absorbed by the system at constant volume is completely utilize to increase the __________ of the system.

39. Heat change at constant pressure from initial to final state of the system is simply equal to the __________.

40. SI unit of measurement of heat change is __________.

Chapter 6

Chemical Equilibrium

1. The reactions, which proceed in both the directions, are called __________ reactions.

2. The reactions, which proceed to one direction only, are called __________ reactions.

3. Reversible reactions are __________ completed.

4. Irreversible reactions are __________ after some time.

5. A reversible reaction is said to be in __________ when the rate of forward reaction becomes equal to the rate of backward reaction.

6. The concentrations of reactants and products are __________ at equilibrium point.

7. The value of Kc depends upon the __________ of the reactants.

8. A increase of the value of Kc tends to move the reaction to the __________ direction.

9. A decrease of the value of Kc tends to move the reaction to the __________ direction.

10. An increase in the concentration of the reactants will move the reaction to the __________ direction.

11. A decrease in the concentration of the reactants will move the reaction to the __________ direction.

12. Equilibrium constant is denoted by __________.

13. When the equilibrium constant value is very __________, we can conclude that the forward reaction is almost completed.

14. When equilibrium constant value is very __________ we can conclude that forward reaction will occur to very little extent.

15. According to __________ principle, if system in equilibrium is subjected to a stress, the equilibrium shifts in a direction to minimize or undo the effect of the stress.

16. In exothermic reaction, the __________ of temperature favour the forward rate of reaction.

17. In endothermic reactions, the __________ of temperature favour the forward rate of reaction.

18. A __________ is a substance which effects the rate of reaction but remains unaltered at the end of the reaction.

19. A catalyst increases the velocity of the reaction by decreasing the __________.

20. The suppression of degree of ionization of a sparingly soluble weak electrolyte by the addition of a strong electrolyte containing an ion in common is called __________.

21. __________ is purified in industries by Common Ion Effect.

22. A reaction moves to the left when the concentrations of the products are __________.

23. A reaction moves to the right when the concentrations of the products are __________.

24. Increase in pressure will move the reaction in the direction of __________ volume.

25. Decrease in pressure will move the reaction in the direction of __________ volume.

26. An increase of temperature favours the formation of products in case of __________ reaction.

27. A decrease of temperature fovours the formation of products in case of __________ reaction.

28. Heating moves an endothermic reaction to the __________.

29. Cooling move an exothermic reaction to the __________.

30. The product of ionic concentration in a saturated solution is called __________ constant.

31. When HCl is added to NaCl, the concentration of __________ ion is increased.

32. Chemical reaction involving the substances in more than one phases are called __________.

33. The formation of NH3 is exothermic process hence __________ temperature will favour the formation of NH3.

34. The formation of NO from N2 and O2 is endothermic process hence __________ temperature will favour the formation of NO.

35. Chemical Equilibrium is __________ equilibrium.

36. Molar concentration is also called __________.

37. The rate at which a substance takes part in a chemical reaction depends upon its __________.

38. __________ principle is applied to all reversible reaction.

39. A common ion __________ the solubility of the salt.

40. Number of moles present per dm3 of a substance is called __________.

Chapter 7

Solutions and Electrolytes

1. A mixture of two or more substances, which are homogeneously mixed, is called a __________.

2. __________ is defined as the amount of solute dissolved in a given amount of solvent.

3. A solution is composed of two components __________ and __________.

4. A solution containing one mole of solute per dm3 of solution is called one __________ solution.

5. Molarity is denoted by __________.

6. 1M solution of NaOH contains __________ gms of it dissolved per dm3 of solution.

7. A solution containing one mole of solute dissolved by per kg of solvent is called __________ solution.

8. Molality is denoted by __________.

9. 1M solution of H2SO4 contains __________ gms of it per kg of solvent.

10. The process in which ions are surrounded by water molecules is called __________.

11. The water molecules attached with the hydrated substance are called __________.

12. Hydrated copper sulphate evolves __________ water molecules on heating.

13. The interaction between salt and water to produce acids and bases is called __________.

14. The products of ionic concentration in a saturated solution at a certain temperature are called the __________.

15. Solubility product constant expressed as __________.

16. The suppression of ionization by adding a common ion is called __________.

17. The process of dissociation of an electrolyte into ions is known as __________.

18. The chemical decomposition of a compound in a solution or in fused state brought about by a flow of electric current is known as __________.

19. Electrolysis is performed in an electrolytic cell, which is known as __________.

20. The positive electrode of a voltmeter is called __________ and negative as __________.

21. A solution, which tends to resist changes in pH is called a __________ solution.

22. A mixture of acetic acid and sodium acetate acts as a __________.

23. According to Sorenson __________ is defined as negative logarithm of the hydrogen ion concentration.

24. pH is mathematically expressed as __________.

25. The pH of a neutral solution is __________.

26. __________ substances have pH values lower than 7.

27. __________ solutions have pH values more than 7.

28. Oxidation is __________ of electron.

29. Reduction is the __________ of electron.

30. Such chemical reactions in which the oxidation number of atoms or ions is changed are called __________ reactions.

31. Oxidation number of a free element is __________.

32. Oxidation number of Oxygen in a compound is __________.

33. The sum of oxidation number of any formula of a compound is __________.

34. The oxidation number of any ion is equal to the __________ on the ion.

35. __________ is the reaction in which an acid reacts with a base to form salt and water.

36. __________ are organic compounds which change colour in accordance with the pH of the medium.

37. An indicator that changes from colourless to pink in the presence of an alkaline solution is called __________.

38. An indicator that changes from red to yellow in the presence of an alkaline solution is called __________.

39. Dissociation constant is denoted by __________.

40. According to Bronsted-Lowry Concept, __________ is the donor of proton and __________ is the acceptor of proton.

41. According to Arrhenius, acid is substance that produces __________ ions when dissolved in water.

42. According to Arrhenius, base is a substance that produces __________ ions when dissolved in water.

43. When ionic product is less than ksp, the solution will __________.

44. When ionic product is greater than ksp, the solution will __________.

45. The electrode at which oxidation takes place is called __________.

46. The electrode at which reduction takes place is called __________.

47. H3O+ ion is called __________ ion.

48. The logarithm of reciprocal of hydroxyl ion (OH)- is called __________.

49. Aqueous solution of NH4Cl is __________ while that of NaHCO3 is __________.

50. The ionic product of [H+] and [OH-] of pure water is __________.

51. An increase in the oxidation number of an element or ion during a chemical change is called __________.

52. A decrease in the oxidation number of an element or ion during a chemical change is called __________.

53. The degree of dissociation __________ with the increase in temperature.

54. The degree of dissociation __________ with the dilution of electrolytic solution.

55. A __________ consists of an electrode immersed in solution of its ion.

56. The potential difference between the electrode and the solution of its salt at equilibrium position is called __________ potential.

57. If the pH of a solution is 14, the solution is __________.

58. If the pH of a solution is 4, the solution is __________.

59. The oxidation number of Mn in KMnO4 is __________.

60. The oxidation number of Fe in FeCl3 is __________.

Chapter 8

Introduction to Chemical Kinetics

1. The branch of chemistry, which deals with the study of rates and mechanisms of chemical reactions, is known as __________.

2. Such reactions, which proceeds with very high velocities and are completed very quickly are called __________ reactions.

3. Such reactions, which take place very slowly, are called __________ reactions.

4. Reactions between silver nitrate and sodium chloride to form white precipitates of silver chloride are an example of __________ reaction.

5. Reactions of Organic compounds are slow and are called __________ reactions.

6. There are some reactions, which proceed slowly with a __________ speed.

7. The rate of __________ reaction can only be determined.

8. The amount of chemical change taking place in concentration of the per unit time is called __________ of reaction.

9. Rate of reaction is expressed in __________.

10. The rate of reaction between two specific interval of time is called __________.

11. The addition energy required to bring about a chemical reaction is called __________.

12. According to __________ theory for a chemical reaction to take place, the reacting molecules must come closed together.

13. The addition of __________ helps the reaction by lowering the energy of activation.

14. The rate of reaction __________ with the increase in concentration of the reacting molecules.

15. When the concentration of both the reacting molecules is double, the probability of collisions between them will be __________ times.

16. By __________ the surface area of the reactants, the rate of reaction is increased.

17. Rate of reaction generally __________ with the rise of temperature.

18. A __________ is a substance, which either accelerates or retards the rate of reaction without taking part in the reaction.

19. In the preparation of Oxygen from Potassium Chlorate, __________ is used as catalyst.

20. In the oxidation of SO2 to SO3 by the contact process for the manufacture of H2SO4 __________ is used as catalyst.

21. An unstable intermediate compound formed during a chemical reaction is called __________.

22. When a catalyst and the reactants are in the same phases, it is known as __________ catalyst.

23. When a catalyst and the reactants are in different phases, it is called __________.

24. When a catalyst increases the rate of reaction, it is called __________ catalyst.

25. When a catalyst retards the rate of reaction, it is called __________ catalyst.

26. A negative catalyst __________ the energy of activation, hence the rate of reaction is decreased.

27. The ratio between the rate of reaction and concentration of reactants is known as __________.

28. Velocity constant is independent of concentration but depends on __________.

29. Ionic reactions are __________ than molecular reactions.

30. The value of specific rates constant for a reaction __________ with time.

31. The sum of all exponents of concentration terms in the equation is called __________.

32. The sum of moles taking part in a chemical reaction is called __________ of the reactions.



Volumetric Analysis (Titrations)



Qs.1 What is Titration?

Ans. The process of adding one solution from the burette into another in the conical flask in order to determine its volume after the completion of the chemical reaction is known as Titration.

Qs.2 What is Neutralization?

Ans. Neutralization is simply a reaction between Hydrogen ion (H+) of an acid and Hydroxide ion (OH-) of the base to form water. In this reaction another class of compound known as “Salt is also produced which remains in the solution as ions, when water is boiled off these ions re-unite to form salt.

Acid + Base → Salt + Water

Qs.3 What is an acid?

Ans. An acid is a substance which gives off proton (H+) in solution or in other words it is a donor of proton e.g. HCl, H2SO4 or HNO3. When dissolve in water ionizes to give Hydrogen ion (H+)

HCL ↔ H+ + Cl-

H2SO4 ↔ 2H+ + SO4--

HNO3 ↔ H+ + NO3-

Qs.4 What do you mean by basicity of an acid?

Ans. It is the number of ionizable Hydrogen presents in the molecule of an acid.

e.g. In the above examples basicity of HCl and HNO3 is one, while that of H2*SO4 is 2.

Qs.5 What is a base?

Ans. A base is now regarded as a molecule or an ion which furnishes OH- ion or accepts a proton given up by an acid. Thus it is proton acceptor.

e.g. NaOH, KOH, Ca (OH)2

NaOH ↔ Na+ + OH-

KOH ↔ K+ + OH-

Ca(OH)2 ↔ CA++ + 2(OH)-

Qs.6 What is meant by acidity of a base?

Ans. It is the number of hydroxyl (OH-) groups present in the molecule of a base.

e.g. In the above examples the acidity of NaOH and KOH is one while that of

Ca (OH)2 is 2.

Qs.7 Why Phenolphthalein is added into the solution of titration flask?

Ans. Phenolphthalein serves as an indicator for determining the end point. It gives pink colour in presence of an alkali and becomes colourless with slight excess of an acid.

Qs.8 While using a burette what precautions are necessary to be observed?

Ans. Burette must be washed first with ordinary water and then rinsed with the solution which is to be taken in it. It must be held vertically and air bubbles must be removed.

While taking a reading the eyes should be in level with the surface of the liquid.

Qs.9 What is a pipette?

Ans. It is an instrument which delivers a definite volume of a liquid. It consists of a glass tube with a cylindrical bulb in the middle and the lower end is drawn into a jet. There is a circular mark on the upper tube.

Qs.10 What is a standard solution?

Ans. A standard solution is a solution of known strength.

Qs.11 What do you mean by strength of a solution?

Ans. Strength of a solution is the quantity of a substance present in any known volume of the solution.

Qs.12 Define a “Normal solution”?

Ans. A standard solution which contains 1 gram equivalent of a substance per dm3 is known as a Normal solution and is denoted by 1 N.

Qs.13 What is a decinormal solution?

Ans. A decinormal solution contains 1/10 fraction of gram equivalent weight of a substance dissolved per dm3 and is denoted by 0.1 N or N/10.

Qs.14 What is the relationship between Normality and Strength per dm 3 of solution?

Ans. Normality =

Qs.15 What is “Acidimetry”?

Ans. It is an operation by which the strength of an alkali is determined by neutralizing it with an acid of known strength in presence of an indicator.

Qs.16 What is “Alkalimetry”?

Ans. It is an operation by which the strength of an alkali is determined by neutralizing it with an alkali of known strength in presence of an indicator.

Qs.17 Define “Equivalent Weight”?

Ans. It is the number of parts by weight that will combine with or displace from 1 part by weight of H2, or 8 parts by weight of O2 or 35.5 by weight of Cl2.

Qs.18 What is “Gram Equivalent Weight”?

Ans. It is the equivalent weight of a substance expressed in gram.

Qs.19 How equivalent weight of an acid is determined?

Ans. Equivalent weight of an acid =

Where basicity is the total number of replaceable Hydrogen.

Qs.20 How equivalent weight of a base is determined?

Ans. Equivalent weight of a base =

Where acidity is the total number of hydroxyl (OH-) groups.

Qs.21 What do you mean by “Standardization of a solution”?

Ans. To standardize means to determine its strength by titration against some standard solution.

Qs.22 What is the equivalent weight of NaOH?

Ans. Eq. Wt. Of a base = = = 40

Therefore, equivalent weight of NaOH is 40.

Qs.23 How a decinormal solution of NaOH is prepared?

Ans. A decinormal solution (0.1 N) of NaOH can be prepared by dissolving 1/10 fraction of its equivalent wt. i.e. 40/10 = 4 gms, in one dm3 of distilled water.

Qs.24 What do you mean by “End-point”?

Ans. It is the exact stage at which the chemical reaction of the titrating solutions is just completed.

Qs.25 Why alkali is taken in burette when phenolphthalein is used as an indicator?

Ans. The appearance of a pink colour at the end point is easily detectable. So it is a better criterion than the disappearance of colour when acid is used in the burette.

Qs.26 How Equivalent weight of Na2CO3 is calculated?

Ans. Na2CO3 is a basic salt. Its equivalent weight can be calculated from the equation of its reaction with an acid e.g. HCl.

Hence the equivalent weight of Sodium Carbonate is determined as follows:

Na2CO3 + 2HCl 2 NaCl + H2O +CO2

2x23+12 2(1+35.5)

+3x16 = 73

= 106

2 HCl = Na2CO3

1 HCl = ½ Na2CO3

Eq. Wt. of Na2CO3= = 53.

Qs.27 What is Methyl Orange?

Ans. It is Sodium salt of an azo dye. It is very good indicator for titrating strong acid against strong base or strong acid against weak base.

Qs.28 What is meant by Anhydrous salt?

Ans. A Salt without water molecule is called Anhydrous.

Qs.29 What indicator is suitable for Sodium Carbonate titration against strong acids and why?

Ans. The pH range of Methyl Orange is pH 3.0 to 4.4.

Hence, it is very suitable indicator when a weak alkali like Sodium Varbonatye is neutralized with a strong acid. In such vases the end point would be at a pH some what below 7.0.

Qs.30 Give the structural formula of Phenolphthalein?

Ans. It is as follows:

Qs.31 Which is the suitable indicator for the titration of?

(i) Weak acid against strong alkali.

(ii) Strong acid against weak alkali.

(iii) Strong acid against strong alkali.

Ans. (i) Phenolphthalein for the titration of weak acid with strong alkali.

(ii) Methyl Orange for the titration of strong acid with a weak alkali.

(iii) Phenolphthalein or Methyl orange for the titration of strong acid with strong alkali (preferably phenolphthalein)

Qs.32 Define Oxidation?

Ans. The loss of electron from an atom, ion or molecule is called Oxidation.

Fe++ ® Fe+++ + e-

Qs.33 Define Reduction?

Ans. Gain of electron or the loss of positive valence is called reduction.

Fe+++ + e- ® Fe++

Qs.34 What are “Oxidation-Reduction Titration”?

Ans. Titrations based upon the reaction between an oxidation agent and reducing agent are known as “Oxidation-Reduction Titrations”.

Qs.35 What are Oxidizing and Reducing agents?

Ans. An oxidizing agent is that substance which oxidizes the other substance e.g., KMnO4 in this titration.

A reducing agent is that which reduces the other substance e.g., Oxalic acid in this titration.

Qs.36 Why do we add equal volume of dilute Sulphuric acid in KMnO4 titration?

Ans. In presence of acid it acts as a strong oxidizing agent and liberates atomic Oxygen from KMnO4.

2 KMnO4 + 3H2SO4 ® K2SO4 + 2 MnSO4 + 3 H2O + 5 [O]

Qs.37 Why do we heat Oxalic acid solution to 60-70º C?

Ans. Oxalic acid reacts with Potassium Permanganate very slowly at room temperature. In order to facilitate the reaction, it is heated to 60º to 70ºC.

Qs.38 What indicator is used in KMnO4 titrations?

Ans. KMnO4 itself acts as an indicator. So no external indicator is required. Near end point it produces a permanent pinkish tinge.

Qs.39 Explain how the change of colour takes place near end point while titrating Oxalic acid with KMnO4 in presence of sulphuric acid?

Ans. In the presence of Sulphuric acid, Potassium permanganate reacts with reducing agent as follows.

2 KMnO4 + 3H2SO4 ® K2SO4 + 2 MnSO4 + 3 H2O + 5 [O]

As the titration proceeds, Potassium Sulphate and Manganese Sulphate are formed, both give colourless solution. As soon as KMnO4 is in excess, the solution becomes pink and so it acts as its own indicator.

Qs.40 Why upper meniscus is noted while using KMnO4 solution in the burette?

Ans. Potassium Permanganate solution is highly coloured and lower meniscus is not distinctly visible. That is why reading of upper meniscus of the liquid is noted.

Qs.41 What is the nature of FeSO4 · 7H2O in Redox titration?

Ans. Ferrous Sulphate acts as reducing agent.

Qs.42 How equivalent weight of FeSO4 · 7H2O is calculated?

Ans. According to the equation:

10 FeSO4 · 7H2O + 5H2SO4 + 5[O] → 5Fe2 (SO4)3 + 12H2O

5[O] combines with 10 FeSO4 · 7H2O

5 x 16 parts by wt. of [O] combines with 10 x 278 parts by

wt. of FeSO4 · 7H2O

8 parts by wt. of [O] combines with = 278

Hence equivalent weight of FeSO4 · 7H2O = 278

Qs.43 Why no external Indicator is required for this titration?

Ans. During titration the dark purple colour of permanganate solution disappears entirely. As soon as the reaction is completed, a single drop of permanganate produces a permanent pinkish tinge to the solution. Thus KMnO4 acts as internal indicator and no external indicator is required.

Qs.44 How equivalent weight of hydrated and Anhydrous Oxalic acid are calculated?

Ans. Mol. Wt. of hydrated Oxalic acid H2C2O4 • 2H2O

2+24+64+2x18=126

Basicity of acid =2

Eq. wt. of hydrated Oxalic acid = = = 63

Mol. wt. of Anhydrous Oxalic acid H2C2O4 = 2+24+64 = 90

Eq. wt. of Anhydrous Oxalic acid = = = 45.

Qs.45 What do you mean by Mohr’s Salt?

Ans. Ferrous Ammonium Sulphate FeSO4 (NH4)2 SO4.6H2O is commonly known as “Mohr’s salt”.

Qs.46 What is the nature of Ferrous Ammonium Sulphate in redox titration?

Ans. Ferrous Ammonium Sulphate acts as a reducing agent.

Qs.47 How can you calculate the equivalent weight of KMnO4?

Ans. KMnO4 *is an Oxidizing agent. The equivalent weight of an Oxidant is the number of parts by weight which gives 8 parts by weight of Oxygen for oxidation.

2 KMnO4 *+ 3H2SO4 → K2SO4 + 2MnSO4 + 3H2O + 5 [O]

5x16 parts by wt. of O2 comes from 2x158 parts of KMnO4

׃٠ 8 ״ ״ ״ ״ ״ ״ = 31.6

Therefore, the equivalent weight of KMnO4 is 31.6

Qs.48 Why Mohr’s salt solution is prepared in water containing dilute Sulphuric acid?

Ans. Solution of Ferrous Ammonium Sulphate (Mohr’s salt) is always prepared by dissolving it in water containing some dilute Sulphuric acid which prevents hydrolysis.

Qs.49 What happens when KMnO4 solution is added in acidified Ferrous Ammonium Sulphate solution?

Ans. When KMnO4 solution is added to Ferrous Ammonium Sulphate is presence of dulute H2SO4, the Ferrous salt is Oxidized to ferric state.

Qs.50 What is the equivalent weight of Ferrous Ammonium Sulphate ?

Ans. Equivalent weight of Ferrous Ammonium Sulphate.

FeSO4. (NH4)2SO4.6H2O = 152+36+32+64+108 = 392



(B) Melting Point & Boiling Point


Qs.1 What is meant by melting or fusion of a substance?

Ans. It is the change of a substance from solid to the liquid state.

Qs.2 Define melting point?

Ans. The temperature at which the solid substance fuses (i.e. changes into liquid) and continue to take place until the whole of the solid is converted into liquid is known as the “Melting point” of the substance.

Qs.3 Why a thin walled capillary tube and not a thick walled tube is selected to determine the Melting point?

Ans. The wall of the capillary tube should be thin otherwise the temperature of the bath may not be equal to the temperature of the substance inside the capillary tube.

Qs.4 Why it is necessary to heat the bath slowly with constant stirring?

Ans. It is necessary to heat the bath slowly with constant stirring to ensure uniformity of temperature; otherwise the rate of rise of temperature is so rapid that it will be difficult to observe the temperature at which the substance just melts.

Qs.5 Why water as a bath cannot be used for the substance having Melting point above 100˚C?

Ans. Boiling point of water is 100˚C. So it cannot be used as a bath for determination of Melting points of such substances. Instead of water, Sulphuric acid or Glycerine can be used.

Qs.6 What is “Boiling”?

Ans. It is a rapid change from the liquid to the gaseous state.

Qs.7 What is meant by the “Boiling point” of a liquid?

Ans. The Boiling point of a liquid is the temperature at which the vapour pressure of the liquid is equal to the atmospheric pressure (i.e. 760 mm.)

Qs.8 Why temperature remains constant at boiling point of a liquid?

Ans. Because the heat is utilized for converting liquid into vapours, so temperature remains constant.

Qs.9 What is the Boiling point of pure water? Is the same at all places?

Ans. The Boiling point of pure water is 100˚C at a pressure of 760 torr. The pressure of air at all places is not the same. So B.P. differs at various places.

Qs.10 What is the difference between Boiling and Evaporation?

Ans. Boiling is a rapid change and takes place through out the mass of the liquid at a definite temperature (i.e., boiling point). While Evaporation is a slow change and takes place at the surface of the liquid at all temperature.

Qs.11 What is the effect of pressure on Boiling point?

Ans. It is raised by the increase of pressure and is lowered by the decrease of pressure. (An increase or decrease of 26.7 m.m. of pressure increase or decrease the Boiling point by 1˚ C).

Qs.12 What is the effect of height on the Boiling point of a liquid?

Ans. With heights, the boiling points are reduced as the pressure decreases.

Qs.13 Why should we stir the liquid in the beaker?

Ans. We should stir it constantly because, this helps the liquid to maintain uniform temperature.




Introduction to Fundamental Concepts of Chemistry


Atom

It is the smallest particle of an element which can exist with all the properties of its own element but it cannot exist in atmosphere alone.


Molecule

When two or more than two atoms are combined with each other a molecule is formed. It can exist freely in nature.


Formula Weight

It is the sum of the weights of the atoms present in the formula of a substance.


Molecular Weight

It is the sum of the atomic masses of all the atoms present in a molecule.


Chemistry

It is a branch of science which deals with the properties, composition and the structure of matter.


Empirical Formula

Definition
It is the simplest formula of a chemical compound which represents the element present of the compound and also represent the simplest ratio between the elements of the compound.

Examples
The empirical formula of benzene is "CH". It indicates that the benzene molecule is composed of two elements carbon and hydrogen and the ratio between these two elements is 1:1.
The empirical formula of glucose is "CH2O". This formula represents that glucose molecule is composed of three elements carbon, hydrogen and oxygen. The ratio between carbon and oxygen is equal but hydrogen is double.


Determination of Empirical Formula

To determine the empirical formula of a compound following steps are required.
1. To detect the elements present in the compound.
2. To determine the masses of each element.
3. To calculate the percentage of each element.
4. Determination of mole composition of each element.
5. Determination of simplest ratio between the element of the compound.


Illustrated Example of Empirical Formula

Consider an unknown compound whose empirical formula is to be determined is given to us. Now we will use the above five steps in order to calculate the empirical formula.

Step I - Determination of the Elements
By performing test it is found that the compound contains magnesium and oxygen elements.

Step II - Determination of the Masses
Masses of the elements are experimentally determined which are given below.
Mass of Mg = 2.4 gm
Mass of Oxygen = 1.6 gm

Step III - Estimation of the Percentage
The percentage of an element may be determined by using the formula.
% of element = Mass of element / Mass of compound x 100
In the given compound two elements are present which are magnesium and oxygen, therefore mass of compound is equal to the sum of the mass of magnesium and mass of oxygen.
Mass of compound = 2.4 + 1.6 = 4.0 gm
% Mg = Mass of Mg / Mass of Compound x 100
= 2.4 / 4.0 x 100
= 60%
% O = Mass of Oxygen / Mass of Compound x 100
= 1.6 / 4.0 x 100
= 40%

Step IV - Determination of Mole Composition
Mole composition of the elements is obtained by dividing percentage of each element with its atomic mass.
Mole ratio of Mg = Percentage of Mg / Atomic Mass of Mg
= 60 / 24
= 2.5
Mole ratio of Mg = Percentage of Oxygen / Atomic Mass of Oxygen
= 40 / 16
= 2.5

Step V - Determination of Simplest Ratio
To obtain the simplest ratio of the atoms the quotients obtained in the step IV are divided by the smallest quotients.
Mg = 2.5 / 2.5 = 1
O = 2.5 / 2.5 = 1
Thus the empirical formula of the compound is MgO

Note
If the number obtained in the simplest ratio is not a whole number then multiply this number with a smallest number such that it becomes a whole number maintain their proportion.


Molecular Formula

Definition
The formula which shows the actual number of atoms of each element present in a molecule is called molecular formula.
OR
It is a formula which represents the element ratio between the elements and actual number of atoms of each type of elements present per molecule of the compound.

Examples
The molecular formula of benzene is "C6H6". It indicates that
1. Benzene molecule is composed of two elements carbon and hydrogen.
2. The ratio between carbon and hydrogen is 1:1.
3. The number of atoms present per molecule of benzene are 6 carbon and 6 hydrogen atoms.
The molecular formula of glucose is "C6H12O6". The formula represents that
1. Glucose molecule is composed of three elements carbon, hydrogen and oxygen.
2. The ratio between the atoms of carbon, hydrogen and oxygen is 1:2:1.
3. The number of atoms present per molecule of glucose are 6 carbon atoms. 12 hydrogen atoms and 6 oxygen atoms.
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Determination of Molecular Formula

The molecular formula of a compound is an integral multiple of its empirical formula.
Molecular formula = (Empirical formula)n
Where n is a digit = 1, 2, 3 etc.
Hence the first step in the determination of molecular formula is to calculate its empirical formula by using the procedure as explained in empirical formula. After that the next step is to calculate the value of n
n = Molecular Mass / Empirical Formula Mass

Example
The empirical formula of a compound is CH2O and its molecular mass is 180.
To calculate the molecular formula of the compound first of all we will calculate its empirical formula mass
Empirical formula mass of CH2O = 12 + 1 x 2 + 16
= 30
n = Molecular Mass / Empirical Formula Mass
= 180 / 30
= 6
Molecular formula = (Empirical formula)n
= (CH2O)6
= C6H12O6


Molecular Mass

Definition
The sum of masses of the atoms present in a molecule is called as molecular mass.
OR
It is the comparison that how mach a molecule of a substance is heavier than 1/12th weight or mass of carbon atom.

Example
The molecular mass of CO2 may be calculated as
Molecular mass of CO2 = Mass of Carbon + 2 (Mass of Oxygen)
= 12 + 2 x 16
= 44 a.m.u
Molecular mass of H2O = (Mass of Hydrogen) x 2 + Mass of Oxygen
= 1 x 2 + 16
= 18 a.m.u
Molecular mass of HCl = Mass of Hydrogen + Mass of Chlorine
= 1 + 35.5
= 36.5 a.m.u


Gram Molecular Mass

Definition
The molecular mass of a compound expressed in gram is called gram molecular mass or mole.

Examples
1. The molecular mass of H2O is 18. If we take 18 gm H2O then it is called 1 gm molecular mass of H2O or 1 mole of water.
2. The molecular mass of HCl is 36.5. If we take 36.5 gm of HCl then it is called as 1 gm molecular mass of HCl or 1 mole of HCl.


Mole

Definition
It is defined as atomic mass of an element, molecular mass of a compound or formula mass of a substance expressed in grams is called as mole.
OR
The amount of a substance that contains as many number of particles (atoms, molecules or ions) as there are atoms contained in 12 gm of pure carbon.

Examples
1. The atomic mass of hydrogen is one. If we take 1 gm of hydrogen, it is equal to one mole of hydrogen.
2. The atomic mass of Na is 23 if we take 23 gm of Na then it is equal to one mole of Na.
3. The atomic mass of sulphur is 32. When we take 32 gm of sulphur then it is called one mole of sulphur.
From these examples we can say that atomic mass of an element expressed in grams is called mole.
Similarly molecular masses expressed in grams is also known as mole e.g.
The molecular mass of CO2 is 44. If we take 44 gm of CO2 it is called one mole of CO2 or the molecular mass of H2O is 18. If we take 18 gm of H2O it is called one mole of H2O.
When atomic mass of an element expressed in grams it is called gram atom
While
The molecular mass of a compound expressed in grams is called gram molecule.
According to the definition of mole.
One gram atom contain 6.02 x 10(23) atoms
While
One gram molecule contain 6.02 x 10(23) molecules.


Avagadro's Number

An Italian scientist, Avagadro's calculated that the number of particles (atoms, molecules) in one mole of a substance are always equal to 6.02 x 10(23). This number is known as Avogadro's number and represented as N(A).

Example
1 gm mole of Na contain 6.02 x 10(23) atoms of Na.
1 gm mole of Sulphur = 6.02 x 10(23) atoms of Sulphur.
1 gm mole of H2SO4 = 6.02 x 10(23) molecules H2SO4
1 gm mole of H2O = 6.02 x 10(23) molecules of H2O
On the basis of Avogadro's Number "mole" is also defined as
Mass of 6.02 x 10(23) molecules, atoms or ions in gram is called mole.

Determination Of The Number Of Atoms Or Molecules In The Given Mass Of A Substance

Example 1
Calculate the number of atoms in 9.2 gm of Na.

Solution
Atomic mass of Na = 23 a.m.u
If we take 23 gm of Na, it is equal to 1 mole.
23 gm of Na contain 6.02 x 10(23) atoms
1 gm of Na contain 6.02 x 10(23) / 23 atoms
9.2 gm of Na contain 9.2 x 6.02 x 10(23) /23
= 2.408 x 10(23) atoms of Na

Determination Of The Mass Of Given Number Of Atoms Or Molecules Of A Substance

Example 2
Calculate the mass in grams of 3.01 x 10(23) molecules of glucose.

Solution
Molecular mass of glucose = 180 a.m.u
So when we take 180 gm of glucose it is equal to one mole So,
6.02 x 10(23) molecules of glucose = 180 gm
1 molecule of glucose = 180 / 6.02 x 10(23) gm
3.01 x 10(23) molecules of glucose = 3.01 x 10(23) x 180 / 6.02 x 10(23)
= 90 gm


Stoichiometry

(Calculation Based On Chemical Equations)

Definition
The study of relationship between the amount of reactant and the products in chemical reactions as given by chemical equations is called stoichiometry.
In this study we always use a balanced chemical equation because a balanced chemical equation tells us the exact mass ratio of the reactants and products in the chemical reaction.
There are three relationships involved for the stoichiometric calculations from the balanced chemical equations which are
1. Mass - Mass Relationship
2. Mass - Volume Relationship
3. Volume - Volume Relationship

Mass - Mass Relationship
In this relationship we can determine the unknown mass of a reactant or product from a given mass of teh substance involved in the chemical reaction by using a balanced chemical equation.

Example
Calculate the mass of CO2 that can be obtained by heating 50 gm of limestone.

Solution
Step I - Write a Balanced Equation
CaCO3 ----> CaO + CO2

Step II - Write Down The Molecular Masses And Moles Of Reactant & Product
CaCO3 ----> CaO + CO2

Method I - MOLE METHOD
Number of moles of 50 gm of CaCO3 = 50 / 100 = 0.5 mole
According to equation
1 mole of CaCO3 gives 1 mole of CO2
0.5 mole of CaCO3 will give 0.5 mole of CO2
Mass of CO2 = Moles x Molecular Mass
= 0.5 x 44
= 22 gm

Method II - FACTOR METHOD
From equation we may write as
100 gm of CaCO3 gives 44 gm of CO2
1 gm of CaCO3 will give 44/100 gm of CO2
50 gm of CaCO3 will give 50 x 44 / 100 gm of CO2
= 22 gm of CO2

Mass - Volume Relationship
The major quantities of gases can be expressed in terms of volume as well as masses. According to Avogardro One gm mole of any gas always occupies 22.4 dm3 volume at S.T.P. So this law is applied in mass-volume relationship.
This relationship is useful in determining the unknown mass or volume of reactant or product by using a given mass or volume of some substance in a chemical reaction.

Example
Calculate the volume of CO2 gas produced at S.T.P by combustion of 20 gm of CH4.

Solution
Step I - Write a Balanced Equation
CH4 + 2 O2 ----> CO2 + 2 H2O

Step II - Write Down The Molecular Masses And Moles Of Reactant & Product
CH4 + 2 O2 ----> CO2 + 2 H2O

Method I - MOLE METHOD
Convert the given mass of CH4 in moles
Number of moles of CH4 = Given Mass of CH4 / Molar Mass of CH4
From Equation
1 mole of CH4 gives 1 moles of CO2
1.25 mole of CH4 will give 1.25 mole of CO2
No. of moles of CO2 obtained = 1.25
But 1 mole of CO2 at S.T.P occupies 22.4 dm3
1.25 mole of CO2 at S.T.P occupies 22.4 x 1.25
= 28 dm3

Method II - FACTOR METHOD
Molecular mass of CH4 = 16
Molecular mass of CO2 = 44
According to the equation
16 gm of CH4 gives 44 gm of CO2
1 gm of CH4 will give 44/16 gm of CO2
20 gm of CH4 will give 20 x 44/16 gm of CO2
= 55 gm of CO2
44 gm of CO2 at S.T.P occupy a volume 22.4 dm3
1 gm of CO2 at S.T.P occupy a volume 22.4/44 dm3
55 gm of CO2 at S.T.P occupy a volume 55 x 22.4/44
= 28 dm3

Volume - Volume Relationship
This relationship determine the unknown volumes of reactants or products from a known volume of other gas.
This relationship is based on Gay-Lussac's law of combining volume which states that gases react in the ratio of small whole number by volume under similar conditions of temperature & pressure.
Consider this equation
CH4 + 2 O2 ----> CO2 + 2 H2O
In this reaction one volume of CH4 gas reacts with two volumes of oxygen gas to give one volume of CO2 and two volumes of H2O

Examples
What volume of O2 at S.T.P is required to burn 500 litres (dm3) of C2H4 (ethylene)?

Solution
Step I - Write a Balanced Equation
C2H4 + 3 O2 ----> 2 CO2 + 2 H2O

Step II - Write Down The Moles And Volume Of Reactant & Product
C2H4 + 3 O2 ----> 2 CO2 + 2 H2O

According to Equation
1 dm3 of C2H4 requires 3 dm3 of O2
500 dm3 of C2H4 requires 3 x 500 dm3 of O2
= 1500 dm3 of O2
Limiting Reactant

In stoichiometry when more than one reactant is involved in a chemical reaction, it is not so simple to get actual result of the stoichiometric problem by making relationship between any one of the reactant and product, which are involved in the chemical reaction. As we know that when any one of the reactant is completely used or consumed the reaction is stopped no matter the other reactants are present in very large quantity. This reactant which is totally consumed during the chemical reaction due to which the reaction is stopped is called limiting reactant.
Limiting reactant help us in calculating the actual amount of product formed during the chemical reaction. To understand the concept the limiting reactant consider the following calculation.

Problem
We are provided 50 gm of H2 and 50 gm of N2. Calculate how many gm of NH3 will be formed when the reaction is irreversible.
The equation for the reaction is as follows.
N2 + 3 H2 ----> 2 NH3

Solution
In this problem moles of N2 and H2 are as follows
Moles of N2 = Mass of N2 / Mol. Mass of N2
= 50 / 28
= 1.79
Moles of H2 = Mass of H2 / Mol. Mass of H2
= 50 / 2
= 25
So, the provided moles for the reaction are
nitrogen = 1.79 moles and hydrogen = 25 moles
But in the equation of the process 1 mole of nitrogen require 3 mole of hydrogen. Therefore the provided moles of nitrogen i.e. 1.79 require 1.79 x 3 moles of hydrogen i.e. 5.37 moles although 25 moles of H2 are provided but when nitrogen is consumed the reaction will be stopped and the remaining hydrogen is useless for the reaction so in this problem N2 is a limiting reactant by which we can calculate the actual amount of product formed during the reaction.
N2 + 3 H2 ----> 2 NH3

1 mole of N2 gives 2 moles of NH3
1.79 mole of N2 gives 2 x 1.79 moles of NH3
= 3.58 moles of NH3

Mass of NH3 = Moles of NH3 x Mol. Mass
= 3.58 x 17
= 60.86 gm of NH3

33 comments:

  1. for god sake why dont you people give answers of these mcqs?

    ReplyDelete
  2. thanks alot for mcqs...............

    ReplyDelete
  3. there is a mistake in 19th Question,....

    ReplyDelete
  4. there is a mistake in 19th Question of 5th chapter.

    ReplyDelete
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  6. mcqs are good but there is no answerkeys

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  8. WHeree ISs LOnGG Questionn ISs NOtt Available....Sooo WHyy..!!!!!!

    ReplyDelete
  9. very nice descriptive questions.

    ReplyDelete
  10. hello dear
    just share me answer key of all mcqs

    ReplyDelete
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  11. This comment has been removed by the author.

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  12. where ix the key of these mcqz????????

    ReplyDelete
  13. There is very good place for any kind of study notes..please check here
    www.webcrase.com

    ReplyDelete
  14. tooooooooooooo goooooooooooood mcqs but answers are not available

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  15. why ans r nt available

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  16. haleema iftikhar22 July 2014 12:38

    toooo much helpful but answers are not available

    ReplyDelete
  17. haleema iftikhar22 July 2014 12:43

    there is mistake in 19 mcqs of 5th chapter plz correct this one

    ReplyDelete
  18. haleema iftikhar22 July 2014 12:46

    there is mistake in19 mcqs

    ReplyDelete
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  19. Want to test / improve your knowledge before attempting pre-entry test for university or college?
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