Chemistry 5 Years Paper Objective
Five Year Papers
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
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
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)
(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
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
Energetic 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
(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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