BASIC
CONCEPTS
Subjective
Q.1 Define an atom. Name some subatomic particles
of atom. Which of them are regarded as fundamental
particles? 11201001
Q.2 What is the direct evidence
of existence of atoms? OR 11201002
Can we actually see the
atoms?
Q.3 (a) Define. 11201003
(i) Molecule
(ii) Macromolecules
(b) What is atomicity? What are
monoatomic, diatomic and triatomic molecules?
Q.4 (a) Define: 11201004
(i) Substance (ii)
Element (iii) Compound
Q. 5 Define
molecular ion. What do you mean by cationic molecular ion and anionic molecular
ion? Give example. 11201005
Q.6 Differentiate between. 11201006
(i)
Atom and Ion (ii)
Atom and molecule
Q.7 (a) Define:
(i) Relative atomic mass (ii) amu 11201007
(b) Write a note on isotopes.
(c) What is isotopy?
Q.8 (a) Explain the term relative abundance of isotopes. 11201008
(b) What are mono-isotopic elements?
Q.9 (a) What is the Mass Spectrometer? How does it work? 11201009
Name the first mass
spectrometer.
(b) Discuss the construction and working of
Dempster’s Mass Spectrometer.
OR
How relative atomic masses of isotopes are determined by mass
spectrometer?
(c) What is the mass spectrum?
(d) Name different techniques used for the separation of Isotopes.
Q.10 (a) Define fractional atomic mass. 11201010
(b) How do you deduce the fractional atomic masses of elements from the
relative isotopic abundance?
Q.11 (a) Define. 11201011
(i) Qualitative
analysis (ii) Quantitative analysis (iii)
Percentage of the elements
(b) How is the percentage
composition of the compounds determined?
Q.12 (a) Define organic
compounds. 11201012
(b) How are organic
compounds analyzed by combustion analysis and how is the percentage
of C, H, O determined? (Board 2005, 2006)
Q.13 (a) Define the following. 11201013
(i) Empirical formula. (ii) Molecular formula.
(iii) Formula mass. (iv) Molecular mass.
(b) How are the followings
determined?
(i) Empirical
formula of a compound. OR (Write down the steps to calculate
empirical formula.) (Board
2014)
(ii) Molecular
formula of a compound
Q.14 (a) Define the following terms and give
three examples of each. 11201014
(i)
Gram atom (ii)
Gram Molecular mass
(iii)
Gram formula (iv)
Gram ion
(b) Define Mole.
Q.15 (a) Explain
the term Avogadro's number. Give its significance. 11201015
(b) Define
Molar volume.
(c) Why
22.414 dm3 of each gas has a different mass but the same
number of molecules?
Q. 16 (a) Define (Board 2004) 11201016
(i) Stoichiometry (ii)
Stoichiometric amount
(b) State two laws on which Stoichiometric calculations
are based.
(c) What are assumptions that must be considered
while doing Stoichiometric calculations?
Q.17 (a) Define
and explain the term Limiting Reactant. Give an example. 11201017
(b) How
Limiting Reactant is determined in a reaction?
Q.18 (a) Define: (Board 2005, 06, 10) 11201018
(i) Yield (ii) Actual Yield (iii) Theoretical Yield (iv) Percentage
Yield
(b) Why
is actual yield of a chemical reaction usually less than its theoretical yield?
SOLVED EXAMPLES
Example: 1
A sample of neon is found to consist of 
and 
in the percentages of 90.92%, 0.26%, 8.82% respectively.
Calculate the fractional atomic mass of neon. 11201019
Given data:
Percentage of 
= 90.92%
Percentage of 
= 0.26%
Percentage of = 8.82%
Fractional atomic mass of Ne = ?
Example: 2
8.657g of a compound were decomposed into
its elements and gave 5.217g of carbon, 0.962g of hydrogen, 2.478g of oxygen.
Calculate the percentage composition of the compound under study. 11201020
Given data:
Mass of compound = 8.657g
Mass of carbon = 5.217g
Mass of hydrogen = 0.962g
Mass of oxygen = 2.478g
% age composition = ?
Example: 3
Ascorbic acid (vitamin C) contains 40.92%
carbon, 4.58% hydrogen and 54.5% of oxygen by mass. What is the empirical formula
of the ascorbic acid? 11201021
Given data:
Mass of carbon = 40.92%
Mass of hydrogen = 4.58%
Mass of oxygen = 54.5%
Empirical formula = ?
Example: 4
A sample of liquid
consisting of carbon, hydrogen and oxygen was subjected to combustion analysis.
0.5439g of the compound gave 1.039g of 
of
. Determine the empirical
formula of the compound. 11201022
Given data:
Mass of organic compound = 0.5439g
Mass of H2O = 0.6369g
Mass of CO2 = 1.039g
Example: 5
The
combustion analysis of an organic compound shows it to contain 65.44% carbon,
5.50% hydrogen and 29.06% oxygen. What is the empirical formula of the
compound? If the molecular mass of this compound is 110.15gmol–1. Calculate
the molecular formula of the compound. 11201023
Given data:
%
age of carbon = 65.44
%
age of hydrogen = 5.50
%
age of oxygen = 29.06
Molar
mass = 110.15 gmol–1
Example: 6
Calculate
the gram atoms (moles) in 11201024
a) 0.1g of sodium
b) 0.1kg of silicon
a) Given data:
Mass of
sodium = 0.1g
Atomic mass
of sodium = 23g/mol
Example: 7
Calculate
the mass of 10–3 moles of MgSO4. 11201025
Given data:
Formula
mass of MgSO4 = 24
+ 96 = 120g mol–1
No. of
moles of MgSO4 = 10–3
Mass of 10–3
moles of MgSO4 = ?
Example: 8
How
many molecules of water are there in 10.0g of ice? Also calculate the number of
atoms of hydrogen and oxygen separately, the total number of atoms and the covalent
bonds present in the sample. 11201026
Given data:
Mass
of ice = 10.0g
Molar
mass of ice 
= 18gmol–1
No.
of atoms of hydrogen = ?
No.
of atoms of oxygen = ?
Total
No. of atoms in the sample = ?
Total
No. of covalent bonds = ?
Example: 9
10.0g
of 
has been
dissolved in excess of water to dissociate it completely into ions.
Calculate: 11201027
a) Number
of molecules in 10.0g of
b) Number of positive and negative
ions in case of complete dissociation in water.
c) Masses of individual ions.
d) Number of positive and negative
charges dispersed in the solution.
Example: 10
A well-known
ideal gas is enclosed in a container having volume 500cm3 at STP.
Its mass comes out to be 0.72g. What is the molar mass of this gas? 11201028
Given data:
Volume of
the gas = 500cm3
Mass of the
gas = 0.72g
Molar mass
of gas = ?
Example: 11
Calculate
the number of grams of 
and water
produced when 14g of KOH are reacted with excess of 
. Also
calculate the number of molecules of water produced. 11201029
Given data:
Mass
of KOH = 14g
Molar
mass of KOH = 39 + 16 + 1 = 56g mol–1
No.
of grams of = ?
No.
of grams of = ?
No.
of molecules of produced = ?
Example: 12
Mg metal
reacts with HCl to give hydrogen gas. What is the minimum volume of HCl
solution (27% by weight) required to produce 12.1g of H2. The
density of HCl solution is 1.14g/cm3. 11201030
Mg(s)
+ 2HCl(aq) → MgCl2(aq) + H2(g)
Given data:
Mass of H2
produced =
12.1g
Density of
HCl solution = 1.14 g/cm3
% of HCl
solution by mass = 27
Example: 13
NH3
gas can be prepared by heating together two solids NH4Cl and Ca(OH)2.
If a mixture containing 100g of each solid is heated then: 11201031
a)
Calculate
the no. of grams of NH3 produced.
b)
Calculate
the excess amount of reagent left unreacted.
2NH4Cl(s) + Ca(OH)2(s)
→ CaCl2(s) + 2NH3(g) + 2H2O(l)
Given data:
Mass of NH4Cl
=
100g
Mass of Ca(OH)2 = 100g
Molar mass
of NH4Cl = 53.5 gmol–1
Molar mass
of Ca(OH)2 = 74 gmol–1
Example: 14
When
limestone (CaCO3) is roasted, quicklime (CaO) is produced according
to the following equation. The actual yield of CaO is 2.5kg when 4.5kg of
limestone is roasted. What is the percentage yield of this reaction? 11201032
CaCO3(g) → CaO(g)
+ CO2(g)
Given data:
Mass of
CaCO3 = 4.5kg = 4.5 ´ 1000 = 4500g
Actual
yield = 2.5kg = 2.5 ´ 1000 = 2500g
Molar mass
of CaCO3 = 100 gmol-1
Molar mass
of CaO = 56 gmol-1
EXERCISE
Q.1 Select
the most suitable answer from the given ones in each question: 11201033
Ans. See in Objective.
Q.2 Fill
in the blanks. 11201034
(i) The unit of relative atomic mass is
__________.
(ii) The exact masses of isotopes can be determined
by __________ spectrograph.
(iii) The phenomenon of isotopy was first
discovered by ___________.
(iv) Empirical formula can be determined by
combustion analysis for those compounds which have __________ and __________in them.
(v) A limiting reagent is that which controls the
quantities of __________.
(vi) 1 mole of glucose has __________ atoms of
carbon, __________ atoms of oxygen and __________
atoms of hydrogen.
(vii) 4 g of CH4 at 0°C and 1 atmospheric
pressure has __________ molecules of CH4.
(viii) Stoichiometric calculations can be
performed only when __________ is obeyed.
Q.3 Indicate
true or false as the case may be. 11201035
(i) Neon has three isotopes and the fourth one
with atomic mass 20.18 amu.
(ii) Empirical formula gives the information about
the total number of atoms present in the molecule.
(iii) During combustion analysis Mg (ClO4)2
is employed to absorb water vapours.
(iv) Molecular formula is the integral multiple of
empirical formula and the integral multiple can never be unity.
(v) The number of atoms in 1.79 g of gold and
0.023 g of sodium are equal.
(vi) The number of electrons in the molecules of CO
and N2 are 14 each, so 1g of each gas will have same number of electrons.
(vii) Avogadro's
hypothesis is applicable to all types of gases i.e. ideal and non-ideal.
(viii)
Actual yield of a chemical reaction may
be greater than the theoretical yield.
Q.4 What
are ions? Under what conditions are they produced? 11201036
Q.5 (a) What are Isotopes? How do you deduce
the fractional atomic masses of elements from the relative isotopic abundance?
Give two examples in support of your answer. 11201037
(b) How does
a Mass Spectrograph shows the relative abundance of isotopes of an element? 11201038
(c) What is
the justification of two strong peaks in the mass spectrum for bromine; while
for iodine only one peak at 127 amu is indicated? 11201039
Q.6 Silver
has atomic number 47 and has 16 known isotopes but two occur naturally i.e.
Ag-107 and Ag-109. Given the following mass spectrometric data, calculate the
average atomic mass of silver. 11201040
Isotopes
Mass (amu) Percentage
abundance
107Ag 106.90509 51.84
109Ag 108.90476 48.16
Q.7 Boron
with atomic number 5 has two naturally occurring isotopes. Calculate the percentage abundance of 10B
and 11B from the following information. 11201041
Q. 8 Define
the following terms and give three examples of each. 11201042
(i) Gram
atom. (ii) Gram
molecular mass
(iii) Gram formula (iv) Gram
ion
(v) Molar volume (vi)
Avogadro's number
(vii)Stoichiometry
(viii)
Percentage yield
Q. 9 Justify
the following statements. 11201043
(a) 23 g of sodium and 238 g of uranium have equal
number of atoms in them.
(b) Mg atom is twice heavier than that of Carbon atom.
(c) 180 g of glucose and 342 g of sucrose have the
same number of molecules but different number
of atoms present in them.
(d) 4.9 g of H2SO4 when
completely ionised in water have equal number of positive and negative ions but the number of positively charged
ions are twice than the number of
negatively charged ions.
(e) One mg of K2CrO4 has
thrice the number of Ions than the number of formula units when ionized in water.
(f) Two grams of H2, 16 g of CH4
and 44 g of CO2 occupy separately the volumes of 22.414 dm3,
although the sizes and masses of
molecules of three gases are very different from each other.
Q.10 Calculate
each of the following quantities: 11201044
(a) Mass
in grams of 2.74 moles of KMnO4.
(b) Moles of O atoms in 9.00g of Mg (NO3)2.
(c) Number of O atoms in 10.037 g of CuSO4.5H2O.
(d) Mass
in kilograms of 2.6 ´ 1020
molecules of SO2.
(e) Moles of Cl atoms in 0.822g C2H4Cl2.
(f) Mass in grams of 5.136 moles of silver
carbonate.
(g) Mass in grams of 2.78 ´ 1021
molecules of CrO2Cl2.
(h) Number of moles and formula units in 100 g of
KClO3
(i) Number of K+ ions, ClO
ions, Cl atoms and O
atoms in (h).
Q.11 Aspartame
the artificial sweetner, has a molecular formula of C14H18N2O5 . 11201045
(a) What
is the mass of one mole of aspartame?
(b) How
many moles are present in 52 g of aspartame?
(c) What
is the mass in grams of 10.122 moles of aspartame?
(d) How
many hydrogen atoms are present in 2.43 g of aspartame?
Q.12 A sample of 0.600 moles of a metal M reacts
completely with excess of fluorine to form 46.8 g MF2. 11201046
(a) How
many moles of F are present in the sample of MF2 that forms?
(b) Which
element is represented by the symbol M?
Q.13 In
each pair, choose the larger of the indicated quantity, or state if the samples
are equal. 11201047
(a) Individual particles: 0.4 moles of
oxygen molecules or 0.4 moles of oxygen atoms.
(b) Mass: 0.4 moles of ozone molecules or
0.4 moles of Oxygen atoms.
(c) Mass: 0.6 moles of C2H4
or 0.6 moles of l2
(d) Individual particles: 4.0g N2O4
or 3.3 g of SO2
(e) Total ions: 2.3 moles of NaClO3
or 2.0 moles of MgCl2.
(f) Molecules: 11.0 g H2O or
11.0 g H2O2.
(g) Na+ ion: 0.500 moles of
NaBr or 0.0145 kg NaCl.
(h) Mass: 6.02 ´ 1023
atoms of 235 U or 6.02 ´ 1023 atoms of 238
U.
Q.14 (a)
Calculate the percentage of nitrogen in the four important fertilizers i.e. NH3,
NH2CONH2, (NH4)2 SO4, and
NH4NO3. 11201048
(b) Calculate
the percentage of nitrogen and phosphorus in each of the following: 11201049
(a) NH4H2PO4 (b) (NH4)2HPO4 (c) (NH4)3PO4
Q.15 Glucose C6H12O6
is the most important nutrient in the cell for generating chemical potential
energy. Calculate the mass %age of each element in glucose and determine the
number of C,H and O atoms in 10.5 g of the sample. 11201050
Given: Formula of glucose = C6H12O6
Problems:
(i) %age of C =
? (ii) %age
of H = ?
(iii) %age of O =
? (iv) Atoms
of H = ?
(v) Atoms of O
=? (vi) Atoms of C =?
Q.16 Ethylene glycol is used as automobile
antifreeze. It has 38.7% Carbon, 9.7% Hydrogen and 51.6% Oxygen. Its molar mass
is 62.1 grams mol-1. Determine its empirical formula. (Board 2005) 11201051
Given: (i) Compound = ethylene glycol (ii) %age of C = 38.7%
(iii)
%age of H = 9.7 % (iv) %age of O = 51.6%
Problem: Empirical
formula = ?
Q.17 Serotenin (molecular mass = 176 g mole–1)
is a compound that conducts nerve impulses in brain and muscles. It contains
68.2% C, 6.86% H, 15.09% N and 9.08% O. What is its molecular formula? 11201052
Given: %age
of C = 68.2 %age of H = 6.86
%age of N = 15.09 %age of O = 9.08
Problem: Molecular formula = ?
Q.18 An unknown metal M reacts with S to form a
compound with a formula M2S3. If 3.12 g of M reacts with
exactly 2.88 g of sulphur, what are the names of metal M and the compound M2S3
? 11201053
Ans: Given: Mass of M = 3.12 g
Problem: (i) Name of metal M = ?
(ii) Name of the compound M2S3
= ?
Q.19 The
octane present in gasoline burns according to the following equation. 11201054
2C8H18
(l) + 25O2 (g) ¾¾¾¾® 16 CO2(g)
+ 18H2O(l)
(a) How
many moles of O2 are needed to react fully with 4 moles of octane? 11201055
(b) How many
moles of CO2 can be produced from one mole of octane? 11201056
(c) How
many moles of water are produced by the combustion of 6 moles of octane? 11201057
(d) If this
reaction is to be used to synthesize 8 moles of CO2, how many grams
of Oxygen are needed? How many grams of
octane will be used? 11201058
Q.20 Calculate the number of grams of Al2S3
which can be prepared by the reaction of
20 g of Al and 30 g of Sulphur. How much the non-limiting reactant is in excess?
11201059
(Board 2005)
Q.21. A mixture of two liquids, hydrazine N2H4
and N2O4 are used as a fuel in rockets. They produce N2
and water vapours. How many grams of N2 gas will be formed by
reacting 100 g of N2H4 and 200 g of N2O4
? (Board 2005) 11201060
2N2H4 + N2O4
¾¾¾¾® 3N2
+ 4H2O
Q.22 Silicon Carbide (SiC) is an important ceramic
material. It is produced by allowing sand (SiO2) to react with
carbon at high temperature. (Board 2004, 2006) 11201061
SiO2 + 3C ¾¾¾¾¾® SiC + 2CO
When 100 kg sand is reacted with excess of
carbon, 51.4 kg of SiC is produced. What is the percentage yield of SiC?
Q.23 (a) What is Stoichiometry? Give its
assumptions. Mention two important laws which help to perform the
Stoichiometric calculations? 11201062
(b) What is limiting reactant? How does it control
the quantity of the product formed? Explain three examples.
Q.24 (a) Define yield. How do we calculate the percentage yield of a
chemical reaction?11201063
(b) What
are the factors which are mostly responsible for the low yield of the products in
chemical reactions? 11201064
Q.25 Explain
the following with reasons. 11201065
(i) Law of conservation of
mass has to be obeyed during stoichiometric calculations.
(ii) Many chemical reactions taking place in our surrounding
involve the limiting reactants.
(ii)
No individual neon atom in the sample of the
element has a mass of 20.18 amu.
(iii)
One mole of H2SO4 should
completely react with two moles of NaOH. How does Avogadro’s number help to
explain it.
(iv)
One mole of H2O has two moles of
bonds, three moles of atoms, ten moles of electrons and twenty-eight moles of
the total fundamental particles present in it.
(v)
N2 and CO have the same number of
electrons, protons and neutrons.
EXPERIMENTAL
TECHNIQUES IN CHEMISTRY
Subjective
Q. 1 (a) Define.
11202001
(i) Analysis
(ii) Analytical chemistry
(b) What is qualitative and
quantitative analysis?
(c) What are the steps
involved in quantitative analysis?
(d) Name the important
techniques of separation.
Q.2 (a) Define
the terms: 11202002
(i) Filtration (ii) Filter (iii) Filtrate (iv) Residue (v) Filter medium
(b) Explain different methods of filtration.
(c) What important points must be kept in mind for
smooth and fast filtration process?
(d) Discuss the important steps during folding to
form the regular filter paper and fluted filter paper. How rate of filtration can be increased?
(e) Draw a diagram of complete filtration assembly
and mention various apparatus/chemicals used.
(f) Write note on filter crucibles. What are their
advantages?
Q.3 (a) Define
crystallization. What is basic principle of crystallization? 11202003
(b) Explain the steps involved in the process
of crystallization.
Q.4 (a) Define
sublimation. Name some substances, which could be purified by sublimation. 11202004 (b) Explain
the process of sublimation.
Q.5 What is solvent extraction? Explain with a
common laboratory example. 11202005
Q.6 Explain Distribution Law or Partition Law. 11202006
Q.7 (a) Define: 11202007
(i) Chromatography (ii) Stationary
phase (iii) Mobile phase
(b) On which
principle chromatography is based?
(c) What is
distribution co-efficient K? How its value helps to explain the distribution of
the components of a mixture between the stationary and mobile phases?
(d) What is adsorption
chromatography and partition chromatography?
Q. 8 (a) What
is paper chromatography? Give its types. 11202008
(b) Explain the procedure involved in the
separation of the components of a mixture by Ascending Paper Chromatography?
(c) What is meant by Retardation factor (Rf)?
(d) Write down the uses of chromatography.
EXERCISE
Q.1 Multiple choice questions: 11202009
Q.2 Fill in the blanks: 11202010
1. A complete chemical characterization of a
compound must include __________.
2. During filtration the tip of the stem of
the funnel should touch the side of the beaker to avoid __________.
3. A fluted filter paper is used to
__________.
4. A solvent used for crystallization is
required to dissolve __________ of the substance at its boiling point and
__________ at room temperature.
5. Repeated solvent extractions using small
portions of solvent are __________ than using a single extraction with larger
volume of the solvent.
Q.3 Tick
the correct sentences. If the sentence is incorrect, write the correct
statement.11202011
(i) A
qualitative analysis involves the identification of elements present in a
compound.
(ii) If the process of filtration is to run
smoothly, the stem of the funnel should remain empty.
Correct
Statement
(ii) If the process of filtration is to run smoothly,
the stem of the funnel should remain continuously full of liquid until liquid
remains in the conical portion.
(iii) If none of the solvents is found suitable for
crystallization a combination of two or more immiscible solvents may be used.
(iv) Paper chromatography is a technique of
partition chromatography.
(v) A solute distributes itself between two
immiscible liquids in a constant ratio of concentrations depending upon the
amount of solvent added.
Correct
Statement
(v) A solute distributes itself between two
immiscible liquids in a constant ratio of concentrations irrespective of the
amount of solute added.
Q.4 Why is there a need to crystallize the
crude product? 11202012
Q.5 A water insoluble organic compound aspirin
is prepared by the reaction of salicylic acid with a mixture of acetic acid and
acetic anhydride. How will you separate the product from the reaction mixture? 11202013
Q.6 A solid organic compound is soluble in
water as well as in chloroform. During its preparation, it remains in aqueous
layer. Describe a method to obtain it from this layer. 11202014
Q.7 The following figure shows a developed
chromatogram on paper with five spots: 11202015
(i) Unknown mixture X (ii) Sample A
(iii) Sample
B (iv) Sample
C
(v) Sample
D
Find out (i) the composition of
unknown mixture X (ii) which sample is impure and what is its composition?
Q.8 In solvent extraction
technique, why repeated extraction using small portions of solvent are more efficient
than using a single extraction but larger volume of solvent? 11202016
Q.9 Write down the main characteristics of a
solvent selected for crystallization of a compound. 11202017
Q.10 You have been
provided with a mixture containing three inks with different colours. Write down the procedure to separate the mixture with the help of paper
chromatography. 11202018
|
Subjective
Q. 1 (a) Define matter. Name the different states of matter. 11203001
(b) Define gas. Why liquids are less common than solids, gases, and
plasma?
(c) Discuss general properties of gases.
Q. 2 Discuss the general properties of liquids and
solids. 11203002
Q. 3 (a) What is gaseous pressure? What are its different units? 11203003
(b) Define: (i) Pascal
(ii) One Atmospheric Pressure
(c) What
are gas laws? Explain.
Q. 4 (a) State and explain Boyle’s Law. 11203004
OR State and explain the Law which explains the
effect of pressure on the volume of gas.
(Board 2005)
(b) How Boyle’s Law is verified experimentally? OR
How
will you verify the law experimentally which relates the volume with pressure? OR
Describe an experiment to demonstrate that
the product of pressure and volume of a gas in an experiment remains constant.
Q. 5 (a) What are isotherms? What happens to position of isotherms when they
are plotted at high temperature for a particular gas? 11203005
(b) Why do we get a straight line when pressure
exerted on a gas is plotted against inverse of volume? This straight line
changes its position in the graph by varying the temperature. Justify it.
Q. 6 (a) State
and explain Charle’s Law. OR 11203006
State and explain the law which shows
temperature volume relationship.
(b) How is Charle’s Law verified experimentally?
(c) How absolute zero of temperature can be
derived from Charle’s Law? OR
Show that Charle’s Law is not obeyed
when temperature is measured on the Celsius scale.
(d) Plot a graph for one mole of an ideal gas to
prove that a gas becomes liquid earlier than
–273.16oC.
Q. 7 (a) What
is general gas equation? Derive it in various forms. OR 11203007
Derive an
expression for the general gas equation or an ideal gas equation. On what
factors the value of Universal gas constant “R” depends?
(b) How
Boyle’s Law, Charle’s law and Avogadro’s law could be derived from the general
gas equation?
(c) What
are the applications of general gas equation?
OR With the help of general gas equation, how will you calculate (i)
molecular mass (ii) density of a gas?
Q. 8 (a) Define universal or general gas
constant. 11203008
(b) Derive
the units for gas constant R in general gas equation.
(i) When
the pressure is in atmosphere and volume in dm3.
(ii) When the pressure is in Nm-2 and volume in
dm3.
(iii) When energy is expressed in ergs.
Q. 9 What is Avogadro’s law of gases? Explain with
examples. 11203009
Q.10 (a) Define
partial pressure. 11203010
(b) State and explain
(c) How general gas equation explains relationship
between the number of moles of a gas and its partial pressure? (Board 2005, 2006)
Q. 11(a) Derive an equation to
find out the partial pressure of a gas knowing the individual moles of
component gases and the total pressure of the mixture. 11203011
(b)
Give important applications
of Dalton’s law of partial pressure.
(c)
Define Aqueous Tension.
Q.12 (a) Define and explain the terms Diffusion and Effusion. 11203012
(b) Explain
Graham’s Law of diffusion. Derive its mathematical expression. (Board
2005)
(c) Describe an experiment to demonstrate the validity of Graham’s
law of diffusion.
Q. 13.What is kinetic molecular theory of gases? Give its postulates. 11203013
Q. 14 (a) What is kinetic gas equation? Explain. 11203014
(b) Define:
(i) Mean square velocity
(ii) Root mean square velocity
(c) Derive an expression to show that
Cr.m.s = 
(d) What is relationship between root mean square velocity and
absolute temperature?
Q. 15. How does the Kinetic molecular theory of gases explains the following
gas Laws?
(Board
2004, 2005 & 2006) 11203015
(i)
Boyle’s Law (ii) Charle’s Law
(iii) Avogadro’s
Law (iv) Graham’s
Law of diffusion
Q.16 (a) How kinetic theory helps to explain the term temperature? OR 11203016
Show that the Kelvin temperature of gas is directly proportional to the
average translational kinetic energy of its molecules.
(b) Is it true that for solids temperature is a measure of
vibrational kinetic energy only?
(c) Define temperature and absolute zero in the light of kinetic theory
of gases.
Q. 17 (a) What is the general principle
of liquefaction of gases? 11203017
(b) Define the terms, critical temperature,
critical pressure and critical volume. Explain
the significance of critical temperature
and pressure by giving a suitable example.
(c) How nature of the gas affect its
critical temperature?
Q. 18 (a) State
and explain Joule Thomson effect. How does it help to explain liquefaction of gases? 11203018
(b) Name different methods used for the
liquefaction of gases. Explain Lind’s method of liquefaction
of gases.
Q. 19.(a) Define an ideal gas and a non-ideal gas. 11203019
(b)Gases show non-ideal behaviour at low temperature and high pressure.
Explain with the help of a graph.
Q. 20. (a) Under
what conditions real gases show deviation from ideal gas laws? State and explain
the causes of the deviation. 11203020
(b) Do
you think that some of the postulates of kinetic molecular theory of gases are
faulty? Point out these postulates.
(c) Why do the gases deviate from
the ideal behaviour at low temperature and high pressure?
Q. 21. (a) What
is Van der Waal’s equation? Derive Van der Waal’s equation for real gases.
11203021
(b) What is physical significance
of Van der Waal’s constants, ‘a’ and ‘b’? Give their
units.
(c) Explain that the excluded volume is
four times the actual volume of the molecules.
Q. 22. (a) What
is plasma? How is it formed? Discuss its characteristics and applications.
(b) Discuss different scales of
thermometry. 11203022
SOLVED EXAMPLES
Example: 1
A gas having a volume of
10dm3 is enclosed in a vessel at 0oC and the pressure is
2.5 atm. This gas is allowed to expand until the new pressure is 2 atm. What will
be the new volume of this gas, if the temperature is maintained at 273 K? 11203023
Given data:
Initial volume of gas (V1) = 10 dm3
Initial temperature (T1)
= 0oC + 273K = 273K
Initial pressure (P1)
=
2.5 atm
Final pressure (P2)
= 2
atm
Final temperature (T2)
= 273K
Final volume (V2)
= ?
Example: 2
250cm3 of
hydrogen is cooled from 127oC to -27oC by maintaining the
pressure constant. Calculate the new volume of the gas at low temperature. 11203024
Given data:
Initial volume (V1) = 250cm3 = 0.25dm3
Initial temperature (T1)
= 127oC
+ 273K = 400K
Final temperature (T2)
= -27oC
+ 273K = 246K
Final volume (V2)
= ?
Example: 3
A sample of nitrogen gas is
enclosed in a vessel of volume 380cm3 at 120oC and
pressure of 101325 Nm–2. This gas is transferred to a 10 dm3
flask and cooled to 27oC. Calculate the pressure in Nm-2
exerted by the gas at 27oC. 11203025
Given data:
Initial volume of the gas (V1) =
380cm3 = 0.38dm3
Initial temperature (T1) = 120oC
+ 273K = 393K
Initial pressure (P1) = 101325Nm–2
Final temperature (T2) = 27oC
+ 273K = 300K
Final volume (V2) = 10dm3
Final pressure (P2) = ?
Example: 4
Calculate the density of CH4(g)
at 0oC and 1 atm pressure. What will happen to the density if (a)
temperature is increased to 27oC (b) the pressure is increased to 2
atmospheres at 0oC. 11203026
Given data:
Temperature of the gas = 0oC +
273K = 273K
Pressure of the gas = 1 atm
Molecular mass of the gas = 16 gmol-1
Gas constant (R) = 0.0821dm3 atm
K-1mol-1
Example: 5
Calculate the mass of 1 dm3 of
NH3 gas at 30oC and 1000 mm Hg pressure, considering that
NH3 is behaving ideally. 11203027
Given data:
Pressure of the gas = 1000 mm Hg =
Volume of the gas = 1 dm3
Temperature of the gas = 30oC + 273K = 303K
Molecular mass of the gas = 17 g mol-1
Example: 6
There is a mixture of hydrogen, helium and
methane occupying a vessel of volume 13 dm3 at 37oC and
pressure of 1 atmosphere. The masses of H2 and He are 0.8 g and 0.12
g respectively. Calculate the partial pressures in torr of each gas in the
mixture. 11203028
Given data:
Volume of the mixture of gases = 13 dm3
Temperature of the mixture = 37 + 273 = 310K
Pressure
of the mixture = 1 atm
Example: 7
250cm3 of the sample of
hydrogen effuses four times as rapidly as 250cm3 of an unknown gas.
Calculate the molar mass of unknown gas. 11203029
Given data:
Suppose unknown gas = x
Rate of effusion of unknown gas (rx)
= 1
Rate of effusion of hydrogen gas 
= 4
Molar mass of H2 gas 
= 2 g mol-1
Molar mass of unknown gas (Mx)
= ?
Example: 8
One mole of methane gas is maintained at
300K. Its volume is 250cm3. Calculate the pressure exerted by the
gas under the following conditions: 11203030
(i)
When the gas is ideal
(ii)
When the gas is non-ideal
a
= 2.253 atm dm6 mol-2 b
= 0.0428 dm3 mol-1
(a) Given data:
V = 250cm3 = 0.25dm3
n = 1 mole
T = 300K
R
= 0.0821 dm3 atm K-1 mol-1
(b) Given data:
V = 0.25 dm3
n = 1 mole
R = 0.0821 dm3 atm K-1
mol-1
T
= 300K
a = 2.253 dm6 atm mol-2
b = 0.0428 dm3 mol-1
EXERCISE
Q.1. Select the correct answer out of the
following alternative suggestions: 11203031
Ans. See in Objective
Q.2. Fill in the blanks. 11203032
(i) The product PV has the S.I. unit of _____________.
(ii) Eight grams each of O2 and H2
at 27oC will have total K.E. in the ratio of _____________.
(iii) Smell of the cooking gas during leakage
from a gas cylinder is due to the property of _____________ of gases.
(iv) Equal _____________ of ideal gases at the
same temperature and pressure contain _____________ number of molecules.
(v) The temperature above which a substance exists only as a gas
is called _____________.
Q.3. Label the following sentences as ‘True or
False’. 11203033
(i) Kinetic energy of molecules of a gas is zero at 0oC.
(ii) A gas in a closed container will exert
much higher pressure at the bottom due to gravity than at the top.
(iii) Real gases show ideal gas behavior at low
pressure and high temperature.
(iv) Liquefaction of gases involves decrease
in intermolecular spaces.
(v) An ideal gas on expansion will show Joule-Thomson effect.
Q. 4. (a) What
is Boyle’s law? Give its experimental verification. 11203034
(b) What are isotherms? What
happens to the positions of isotherms when they are plotted at high temperature for a particular gas?
(c) Why do we get a straight line
when pressures exerted on a gas are plotted against inverse of volumes? This straight line changes
position in the graph by varying the temperature. Justify it.
(d) How will you explain that the
value of the constant k in the equation PV = k depends upon:
(i) The
temperature of gas (ii) The quantity of a gas?
Q.5. (a) What
is Charles’s Law? Which scale of temperature is used to verify that
V/T = k (pressure and number of moles are constant)? 11203035
(b) A sample
of carbon monoxide gas occupies 150ml at 25oC. It is then cooled at
constant pressure until it occupies 100ml. What is the new temperature?
(c) Do you
think that the volume of any quantity of a gas becomes zero at – 273oC.
Is it not against the law of conservation of mass? How do you deduce the idea of
absolute zero from this
information?
Q. 6.(a) What is Kelvin scale of temperature? Plot a
graph of one mole of an ideal gas to prove that a gas becomes liquid, earlier
than –273.16 oC. 11203036
(b) Throw some light on a factor 1/273 in Charle’s Law.
Q.7 (a) What is the general gas equation? Derive it in various forms. 11203037
(b) Can we determine the molecular mass of an unknown gas if we know
the pressure, temperature and volume alongwith the
mass of the gas?
(c) How
do you justify from general gas equation that the increase in temperature or
decrease of pressure decreases the density of the gas?
(d) Why do we feel comfortable in expressing the densities of gases
in the units of g dm-3 rather
than g cm-3 a unit which
is used to express the densities of liquids and solids?
Q. 8 Derive the units of gas
constant R in general gas equation. 11203038
(a) When the pressure is in the atmosphere and volume in dm3.
(b) When the pressure is in Nm-2 and volume in m3.
(c) When energy is expressed in ergs.
Q. 9 (a) What is Avogadro’s law of gases? 11203039
(b) Do you think that 1 mole of H2, and 1 mole of NH3
at 0oC and 1 atm-pressure will have Avogadro’s number of particles. If not, why?
(c) Justify 1 cm3 of H2 and 1 cm3
of CH4, at STP will have same number of molecules, when
one molecule of CH4 is 8 times heavier than that of hydrogen.
Q. 10. (a). Dalton’s law of partial pressure is only
obeyed by those gases which don’t have attractive forces among their molecules.
Explain it. 11203040
(b). Derive an equation to find out the partial pressure
of a gas knowing the individual moles of component gases and the total pressure
of the mixture.
(c) Explain that the process of respiration
obeys the Dalton’s law of partial pressures.
(d) How do you
differentiate between diffusion and effusion? Explain Graham’s law of diffusion.
Q. 11. (a) What is critical
temperature of a gas? What is its importance for liquefaction of gases? Discuss Lind’s method of liquefaction
of gases. 11203041
(b) What
is the Joule Thomson Effect? Explain its importance in Lind’s method of liquefaction of gases.
Q. 12 (a) What is Kinetic
molecular theory of gases? Give its postulates. 11203042
(b) How does Kinetic molecular theory of gases? Explain the following
gas laws.
(i) Boyle’s Law. (ii) Charle’s law
(iii) Avogadro’s law (iv) Graham’s law of diffusion
Q. 13 (a) Gases show non ideal
behavior at low temperature and high pressure. Explain this with
the help of a graph. 11203043
(b) Do you think that some of the postulates of
kinetic molecular theory of gases are faulty? Point out these postulates.
(c) Hydrogen and Helium are ideal at room temperature, but SO2
and Cl2 are
non-ideal.
How will you explain this?
Q. 14 (a) Derive Van der Waal’s
equation for real gases. 11203044
(b) What
is the physical significance of Van der Waal’s constants, ‘a’ and ‘b’? Give
their units.
Q. 15 Explain
the following facts. 11203045
(i) The plot of PV versus P is a straight line at
constant temperature and with a fixed number of moles of an ideal gas.
(ii) The straight line in (a) is parallel to x-axis
and goes way from the pressure axis at higher pressure for many gases.
(iii) Pressure of NH3 gas at given
conditions (say 1 atm pressure and room temperature) is less as calculated by
Van der Waal’s equation than that calculated by general gas equation.
(iv) Water vapours do not behave ideally at 273 K.
(v) SO2 is comparatively non-ideal at
273oC but behaves ideally at 327oC.
Q.16. Helium
gas in a 100cm3 container at a pressure of 500 torr is transferred
to a container with a volume of
250cm3. What will be the new pressure (a) if no change in
temperature occurs (b) if its
temperature changes from 20° C to 15°C? 11203046
Q.17 (a) What
are the densities in kg/m3 of the following gases at STP (P=101325
Nm-2,
T=273K,
molecular mass is in kg mole–1)? 11203047
(i) methane (ii) oxygen (iii) hydrogen
(b) Compare the values of densities in proportion
to their molar masses.
(c) How
do you justify that increase of volume upto 100 dm3 at 27°C .of 2
moles of NH3 will allow the gas behave ideally as compared to STP
conditions?
Q 18. A sample of krypton with a volume of 6.25
dm3, a pressure of 765 torr and a temperature of 20°C is expanded to a volume of 9.55 dm3
and a pressure of 375 torr. What will be its final
temperature in °C? 11203048
Q 19. Working at a vacuum line, a chemist
isolated a gas in a weighing bulb with a volume of 255cm3,
at a temperature of 25°C and under a pressure in the bulb of 10.0 torr. The gas
weighed 12.1 mg. What is the molecular
mass of this gas? 11203049
Q.20 What
pressure is exerted by a mixture of 2.00 g of H2 and 8.00 g of N2
at 273 K in a 10dm3 vessel? 11203050
Q 21.(a) The relative densities
of two gases A and B are 1:1.5. Find out the volume of B which will
diffuse in the same time in which 150 dm3 of A will diffuse? 11203051
(b) Hydrogen (H2) diffuses through a porous plate at a
rate of 500 cm3 per minute at 0°C. What is the rate of diffusion of
oxygen through the same porous plate at 0°C?
(c) The rate of effusion of an unknown gas A through a pinhole is
found to be 0.279 times the rate of effusion of H2 gas through the
same pinhole. Calculate the molecular mass of the unknown gas at STP.
Q. 22 Calculate
the number of molecules and the number of atoms in the given amounts of each
gas. 11203052
(a) 20 cm3 of CH4 at 0°C and pressure of 700 mm of mercury.
(b) 1 cm3 of NH3 at 100°C and pressure of 1.5 atm.
Q.23 Calculate
the masses of 1020 molecules of each of H2, O2
and CO2 at STP. What will happen
to the masses of these gases, when the temperature of these gases is increased by
100oC and the pressure is
decreased by 100 mm of Hg? 11203053
Q. 24. (a) Two moles of NH3 are enclosed in 5 dm3
flask at 27°C. Calculate the pressure exerted by the gas
assuming that 11203054
(i) Gas
behaves like an ideal gas (ii) Gas behaves like real gas
a = 1.17
atm dm6 mol–2
b = 0.0371
dm3 mol–1
(b) Also
calculate the amount of pressure lessened due to forces of attractions at these
conditions
of volume and temperature.
(c) Do you expect the same decrease in the pressure of
2 moles of NH3 having a volume of 40 dm3 and at
temperature of 27°C?
|
Subjective
Q1. Discuss
different types of forces among the molecules in detail. 11204001
Q2. Explain
the factors on which London forces depend. 11204002
Factors affecting the London forces:
Q3. What is Hydrogen bonding? Discuss it with examples. 11204003
Hydrogen Bonding
Q4. Describe
the importance of hydrogen bonding in different compounds. 11204004
Properties of Compounds containing hydrogen
bonding:
Q5. (a) Define evaporation. 11204005
(b) How evaporation causes cooling?
(c) What factors affect evaporation of a liquid?
Q6. (a) What is meant by vapour pressure of a
liquid? 11204006
(b) On what factors, vapour pressure of a liquid
depends?
(c) How vapour pressure of a liquid can be
measured by Manometric method?
Q7. (a) What is boiling point? 11204007
Q8. (a) What is the
effect of external pressure on boiling point of a liquid? 11204008
(b) Discuss vacuum
distillation and its advantages.
Q9. Discuss
energetics of phase changes in detail. 11204009
Energetics of phase changes: (Board 2003)
Q10. What
are liquid crystals? Give their uses in daily life. (Board 2014) 11204010
|
Section-II
Q1. (a) Describe some characteristics of
solid substances. 11204011
(b) Discuss types of solids on the basis of
arrangement of particles.
Q2.(a) Discuss properties of crystalline solids in detail. 11204012
(b) What is transition temperature? Give some
examples.
Q3. What
is crystal lattice and a unit cell? Describe different crystal system in which
crystals have been grouped. 11204013
Q4. In how
many types crystalline solids have been classified on the basis of type of
bonds? Discuss each type in detail. 11204014
Q5. (a) What
are ionic solids? Give examples. 11204015
(b) Discuss
properties of Ionic solids.
Q6. What
are molecular solids? Write their properties. (Board 2014) 11204016
Q7. (a) What are metallic solids? 11204017
(b) How the structure of metals explained
by different theories?
(c) Discuss properties of metallic solids.
Q8. How
study of crystalline solids help to calculate avogadro’s number? 11204018
Determination of Avogadro’s Number (NA):
EXERCISE (LIQUIDS)
Q.1. MCQs.
Choose the best answers: 11204019
Q.2. Fill
in the blanks with suitable words. 11204020
(i) The
polarizability of noble gases __________ down the group and results in the
increase in their boiling points.
(ii) __________
is developed in acetone and chloroform when they are mixed together.
(iii) Exceptionally
weak __________ of HF is due to strong hydrogen bonding present in it.
(iv) The
rate of increase of vapour pressure of water __________ at high temperature.
(iii)
During the formation of ice from liquid water
there is a _________ % increase in volume.
(iv)
A layer of ice on the surface of water _______
the water underneath for further heat loss.
(vii) Evaporation is a __________ process.
(viii) Liquid Crystals are used in the display of __________
devices.
(ix) The
concept of dynamic equilibrium is the ultimate __________ of all reversible
systems.
(x) DHv
of C6H14 should be ___________ than that of C2H6.
Q.3. Indicate the true or false as the case may
be: 11204021
(i) Dipole-Dipole forces are weaker than
dipole-induced dipole forces.
(ii) The ion-dipole interactions are responsible
for the dissolution of an ionic substance in water.
(iii) The high polarizability of iodine is
responsible for its existence in solid form, different from other halogens.
(iv) The strong hydrogen bonding in H2S
makes it different from water.
(v) Hydrocarbons are soluble in water because they
are polar compounds.
(vi) The viscosity’s of liquids partially depend
upon the extent of hydrogen bonding.
(vii) The state of equilibrium between liquid state
and vapours is dynamic in nature.
(viii) Heat of vaporization of liquids depend upon
the intermolecular forces of attraction present between their molecules.
(ix) Ice does not show any vapour-pressure on its
surface at –1oC
(x) Boiling point of a liquid is independent of
external pressure.
Q.4. (a) What
type of intermolecular forces will dominate in the following liquids? 11204022
(i) Ammonia, NH3 (ii) Octane C8H18
(iii) Argon, Ar (iv) Propanone
CH3COCH3
(v) Methanol (CH3OH)
Q.5. Explain
the following with reasons. 11204023
(i)
In the hydrogen bonded structure of HF, which is
the stronger bond: the shorter covalent bond or the longer bond between different
molecules.
(ii) In a
very cold winter the fish in garden ponds owe their lives to hydrogen bonding.
(iii)
Water and
ethanol can mix easily in all proportions.
(iv) The origin of the intermolecular forces in
water.
Q.6a) Briefly
consider some of the effects on our lives if water has only a very weak
hydrogen bonding present among its molecules. 11204024
b) All gases have a characteristics critical
temperature. Above the critical temperature, it is possible to liquefy a gas.
The critical temperatures of CO2 and CH4 are 31.14oC
and
–81.9oC, respectively. Which gas has the stronger intermolecular
forces? Briefly explain your choice.
Q.7 Three liquids have the properties, mentioned
against their names? 11204025
(i) Molecular
formula Water Propanone Pentane
H2O C3H6O C5H12
(ii) Relative
molecular mass 18 58 72
(iii) Enthalpy
change of vaporization 41.1 31.9 27.7
(iv) Boiling
point 100 56 36
(a) What
type of intermolecular forces predominates in each liquid?
(i) in
water (ii) in propanone (iii)
in pentane
Q.8 Describe the various forces responsible for
keeping the particles together in the following elements and compounds and
their effects on physical properties making use of the data below: 11204026
|
Substance |
Formula |
Molar Mass |
M.PoC |
|
Neon |
Ne |
20 |
-248 |
|
Argon |
Ar |
40 |
-189 |
|
Water |
H2O |
18 |
0 |
|
Sodium fluoride |
NaF |
42 |
993 |
|
Diamond |
C |
12 |
3350 |
Q.9 The boiling points and molar masses of
hydrides of some first row elements are tabulated below: 11204027
|
Substance |
Boiling point(K) |
Molar Mass (g/mol) |
|
CH4 |
109 |
16 |
|
NH3 |
240 |
17 |
|
H2O |
373 |
18 |
Suggest reasons for the
difference in their boiling points in terms of their type of molecules involved
and the nature of the forces present between them.
Q.10 Explain
the term saturated vapour-pressure. Arrange in order of increasing vapour pressure
1 dm3 water, 1dm3 ethanol, 50 cm3 of water, 50
cm3 ethanol and 50 cm3 of ether.
11204028
Q.11 While
a volatile liquid standing in a beaker evaporates, the temperature of the
liquid remains the same as that of its surrounding if the same liquid is
allowed to vaporize into atmosphere in an insulated flame, its temperature
falls below that of its surrounding. Explain the difference in behaviour. 11204029
Q.12 How
does the hydrogen bonding explain the indicated properties of the following
substances? 11204030
(a) Structure of DNA (b) Formation of ice and its lesser
density than liquid water.
Q.13 What are liquid Crystals? Give their uses in
daily life? 11204031
Q.14 Explain the following with reasons: 11204032
EXERCISE (SOLIDS)
Q.1. Multiple
choice questions. 11204033
Ans. See in objective.
Q.2. Fill
in the blanks. 11204034
(i) In a
crystal lattice, the number of nearest neighbours to each atom is called the
__________.
(ii) There are
__________ Bravis lattices.
(iii) A
pseudo solid is regarded as __________ liquid.
(iv) Glass may
begin to crystallize by a process called __________.
(v) Crystalline
solids which exhibit the same __________
in all directions are called __________.
(vi) The branch
of science which deals with the __________ of crystals is called
crystallography.
Q.3. Indicate
True / False as the case may be: 11204035
(i) There are
five parameters in unit cell dimensions of a crystal.
(ii) Ionic
crystals are very hard, have low volatility and very low melting and boiling
points.
(iii) The
value of lattice energy of the ionic substances depends upon the size of ions.
(iv) Molecular
orbital theory of solids is also called band theory.
(v) Ionic
solid is a good conductor of electricity in the molten state.
Q.4. What are solids? Give general properties of solids.
How do you differentiate between crystalline solids and amorphous solids? 11204036
Q.5. What
is the co-ordination number of an ion? What is the coordination number of the cation
in (a) NaCl structure and (b) CsCl structure? Explain the reason for this
difference.
Q.6. Explain
the following with reasons: 11204038
(i) Sodium
is softer than Copper but both are very good electrical conductors.
|
Subjective
Q1: How were cathode rays
discovered? Discuss their properties.
Discovery of Electron (Cathode rays): 11205001
Q2: Write a note on J.J
Thomson’s experiment for the measurement of charge on an electron.
Measurement of charge to Mass Ratio of Electron: 11205002
Q3: (a)Write a note on
Millikan’s oil drop experiment for the measurement of charge on electron. (Board 2004, 2005) 11205003
(b) How can we calculate the mass of an electron?
Q4. What are positive rays? How were they
discovered? Discuss their properties.
Discovery of Proton (Positive rays) (Board 2003, 2005) 11205004
Q5: Write a note on discovery of neutron. Describe properties and uses of
neutrons.
Discovery of Neutron: (Board 2004,
2005) 11205005
Q6: (a) Discuss Rutherford’s
atomic model in detail. 11205006
(b) Give postulates of Planck’s quantum theory.
Q7: What is a spectrum?
Discuss different types of spectrum. 11205007
SPECTRUM: (Board 2003,
2005)
Q8. Discuss Rutherford’s atomic model in detail. Derive an expression for
radius of an orbit and energy of a revolving electron by this model. (Board
2004, 2005) 11205008
Q9: What is Hydrogen spectrum?
Discuss it on the basis of Bohr’s model. 11205009
Q10: (a) Discuss defects of
Bohr’s atomic model. (Board 2005, 2006) 11205010
(b) What are X-rays? How are they
produced?
EXERCISE
Q1. Select the most suitable answer for the
given questions.
Q2. Fill
in the blanks with suitable words. 11205011
(i)
b- particles are nothing but _______ moving with a very
high speed.
(ii)
The
charge on one mole of electrons is _______ coulombs.
(iii)
The
mass of hydrogen atom is ______ grams.
(iv)
The
mass of one mole of electrons is _____.
(v)
Energy
is ______ when electron jumps from higher to a lower orbit.
(vi)
The
ionization energy of hydrogen atom can be calculated from _____ model of atom.
(vii)
For d sub-shell, the azimuthal quantum number
has value of ______.
(viii)
The number of electrons in a given subshell is
given by formula________.
(ix)
The
electronic configuration of H– is___________.
Q3. Indicate
true or false as the case may be. 11205012
(i)
A
neutron is slightly lighter particle than a proton.
(ii)
A
photon is the massless bundle of energy but has momentum.
(iii)
The
unit of Rydberg constant is the reciprocal of unit of length.
(iv)
The
actual isotopic mass is a whole number.
(v)
Heisenberg’s
uncertainty principle is applicable to macroscopic bodies.
(vi)
The
nodal plane in an orbital is the plane of zero electron density.
(vii)
The number of
orbital present in a sub-level is given by the formula (2l + 1).
(viii)
The magnetic
quantum number was introduced to explain Zeeman and Stark effects.
(ix)
Spin quantum number tells us the
direction of spin of electron around the nucleus.
EXERCISE NUMERICALS
Q.1 A photon of light with energy 10-19 J is emitted
by a source of light. 11205013
(a) Convert this energy into wavelength, frequency
and wave number of the photon in terms of meters, hertz and m-1,
respectively.
(b) Convert this energy of photon into ergs and
calculate wavelength in cm, frequency in Hz and wave number in cm-1.
Q. 2 The formula for calculating the energy of
electron in hydrogen atom given by Bohr’s model is En = 
. Calculate the energy of electron in first orbit of
hydrogen. 11205014
Q. 3 Bohr’s equation for the radius of nth orbit
of electron in hydrogen atom is rn = 
. 11205015
(a) When the electron moves from n = 1 to
n = 2 show how much does the radius of the orbit increases.
(b) What is the
distance travelled by the electron when it goes from n=2 to n =3 and n=9 to
n=10?
Q. 4 Calculate the value of principle quantum
number if an electron in hydrogen atom revolves in an orbit of energy
= 0.242 ´ 10-18J. 11205016
Q. 5 Calculate
the wave number of the photon when the electron jumps from. 11205017
(i) n = 5 to n = 2
(ii) n = 5 to n = 1
In which series of spectral
lines and spectral regions these photons appear?
Q. 6 A photon of a wave number 102.70 ´ 105 m-1 is emitted
when electron jumps from higher orbitals to n = 1. 11205018
(a) Determine
that orbit from where the electron falls.
(b) Indicate
the name of the series to which this photon belongs.
(c) If
the electron will fall from higher orbit to n = 2, then calculate the wave
number of the photon emitted. Why this energy difference is so small as
compared to that in part (a)?
Q. 7 (a) What is De-Broglie’s wavelength of an
electron in meters travelling at half a speed of light? 11205019
(b) Convert the mass of electron into grams and
velocity of light into cm s-1 and then
calculate the wavelength of an
electron in cm.
(c) Covert
the wavelength of electron from meters to (i) nm (ii) 
(iii) Pm
Q. 8 Bohr’s
formula for the energy levels of hydrogen atom for any system, say H, He+1, Li+2
etc. 11205020
En
= 
En
= - k 
For HZ = 1 and for He+2 Z
= 2
(a) Draw an energy level diagram for hydrogen and
He+1
The energy level diagrams of H and He+1 are similar in the
sense that, the differences go on decreasing from lower to the higher levels,
but gaps of energies in He+2 are more than those of H.
(b) Thinking
that k = 2.18 ´ 10-18J, Calculate the energy needed to remove the
electron from hydrogen atom
and He+1 to give H+1 and He+2
Q.9. Answer the following questions, by
performing the calculations. 11205021
(a) Calculate
the energy of first five orbits of hydrogen atom and determine the energy difference
between them.
(b)
Justify that energy
difference between second and third orbits is approximately five times
smaller than that between first and second orbits.
(c)
Calculate the energy of
electron in He+ in first five orbits and justify that the energy differences
are different from those of hydrogen atom.
(d)
Do you think that groups
of the spectral lines of He+ are at different places than those for hydrogen
atom? Give reasons.
CHEMICAL BONDING
Subjective
Q.1 What
is a Chemical Bond? Discuss the formation of Ionic and Covalent Bond. 11206001
Q2: Discuss the energetics of bond formation in
detail. 11206002
Q.3 What
is Coordinate Covalent Bond? (Board 2014) 11206003
Q. 4 Define
ionization energy and electron affinity. What factors affect these properties? How
these quantities change with increase in mass number? 11206004
Q. 5 Write
Lewis structures for the following compounds. 11206005
|
(i) |
HCN |
(ii) |
CCl4 |
(iii) |
CS2 |
(iv) |
H3NAlF3 |
Q. 6 Define
electronegativity. Give its trend in the periodic table. (Board 2005, 14) 11206006
Q. 7 Explain
the following: 11206007
(i) The
melting and boiling points of electrovalent compounds are very high as compared
with those of covalent compounds.
(ii) Why
does solid sodium chloride not conduct electricity? What will happen if
electric current is passed through
molten sodium chloride or its aqueous solution?
(iii) In
many cases, the distinction between a co-ordinate covalent bond and a covalent
bond vanishes after a bond formation.
Explain with the help of an example.
Q.8 Explain
qualitatively the valence bond theory. How does it differ from molecular
orbital theory? 11206008
Q. 9 Give the main postulates of VSEPR theory. Apply
this theory to derive the shapes of following
molecules: 11206009
|
(i) |
BeCl2 |
(ii) |
BCl3 |
(iii) |
SO2 |
(iv) |
CH4 |
(v) |
H2O |
Q.10 Explain
atomic orbital hybridization with specific examples of sp3, sp2
and sp mode of hybridization. 11206010
Q. 11 (a)
Give the main points of the Molecular Orbital Theory? Explain the formation of Bonding
and anti-bonding Molecular orbitals with the help of different types of overlaps. 11206011
(b) Apply the M.O. treatment to the following
molecules:
(i) H2 (ii) He2 (iii) N2
Figure 6.12: formation of N2 Molecule – Molecular Orbital
Treatment
Q.12 How does
Molecular orbital theory explains Paramagnetic behaviour of oxygen? 11206012
Q.13 Define bond energy. Explain various
parameters which determine its strength.
(Board 2003, 2005, 2010) 11206013
Q.14 Define dipole moment. How does it help
to find the shapes of molecules? 11206014
Q. 15 The bond length of H – Br is 1.4 ´
10-10 m. Its observed dipole – moment is
0.79. Find the percentage ionic character of the bond. 11206015
Q. 16 PF3 is a polar molecule with dipole
moment 1.02 D so P – F bond is polar-Si, being in proximity of P in the periodic
table it is expected that Si – F bond would also be polar but Si-F4 has zero dipole moment. Why? 11206016
Q.
17 Explain hybridization schemes for
geometrical shapes of molecules. 11206017
Q.
18 The linear geometry of BeCl2
suggests that central Be atom is sp hybridized. What type of hybridization a
central atom undergoes when the atoms bonded to it are located at the corners
of (a) an equilateral triangle and (b) a regular tetrahedron? 11206018
Q. 19 What
do you understand by the term Electronegativity? Discuss its variations with
respect to periodic table. How does it affect the bond strength? 11206019
Q. 20 Why
chemical reactions of ionic compounds are faster than those of covalent
compounds? 11206020
Q. 21(a) Differentiate between VBT and MOT.
(b) How can the bonding in the following molecules
be explained with respect to valence bond theory? 11206021
Cl2,
O2, N2, HF, H2S
Q. 22(a) Write
short notes on: 11206022
(i) Ionic Radii (ii) Covalent
Radii (iii) Atomic Radii
(a) What is the effect of bonding on the
properties of compounds?
Q. 23 The molecules NF3, BF3
and CIF3 all have molecular formula of the type KF3. But
they have different structural formulas. Keeping in view VSEPR theory sketch
the shape of each molecule and explain the origin of differing in shapes. 11206023
Q. 24 These species NH
, NH3, NH
have bond angles of
105° and 109.5° respectively. Justify these values by drawing
their structures. 11206024
Q. 25 (a) Explain atomic orbital hybridization
with reference to sp3, sp2 and sp modes of hybridizations
for PH3, C2H4 and C2H2.
Discuss geometries of CCl4, PCl3 and H2S by
hybridization of central atoms. 11206025
Q. 26 (a) Give the basis of the molecular
orbital configurations of the following: 11206026
(i) He2 (ii) N2 (iii) O2
(iv) O
(v) O
(b) How does
molecular orbital theory explain the paramagnetic character of O2, O
and O
species?
Q. 27 (a) Sketch the molecular orbital pictures
of 11206027
(i) p (2px) and p* (2px)
(ii) O2, 
, 
(iii) He2 and Ne2
(b) Sketch the hybrid orbitals of the
species, PCl3 SF6, SiCl4 and NH4+.
Q. 28 (a)
Define bond energy. Explain the various parameters which determine its
strength.
11206028
Q. 29 Which
of the following molecules will be polar or non-polar? Sketch the structures
and justify your answer. 11206029
(i) CCl4 (ii) SO3 (iii) SF4 (iv) NH3
(v) PF5 (vi) SO2 (vii) SF6 (viii) IF7
(SOLVED
EXAMPLE)
Example 1:
The
observed dipole moment of HF is 1.90D. Find the percentage ionic character in 
bond. The distance
between the charges is 
.
(Unit
positive charge = 
(EXERCISE)
Q.1. Select
the correct statement. 11206030
Ans. See in Objective.
Q2. Fill in the blanks. 11206031
(i)
The
tendency of atoms to attain maximum of ______ electrons in the valence shell is
called completion of ______.
(ii)
The
geometrical shape of SiCl4 and PCl3 can be explained on
the basis of ________ and hybridizations.
(iii)
The
VSEPR theory stands for ________.
(iv)
For N2
molecule, the energy of s 2px orbital is ________ than p 2py
orbital.
(v)
The
paramagnetic property of O2 is well explained on the basis of MO theory
in terms of the presence of ________ electrons in two MO orbitals.
(vi)
The bond
order of N2 is ________ while that of Ne2 is ________.
(vii)
The
values of dipole moment for CS2 is ________ while for SO2
is ________.
Q3. Classify the statements as true or false. Explain with reasons. 11206032
(i)
The core
of an atom is the atom minus its valence shell.
(ii)
The
molecules of nitrogen N º N and acetylene HC º HC are not
isoelectronic.
(iii) There are four coordinate covalent bonds in 
ion.
(iv) A s – bond is stronger than a p – bond and
electrons of s –bond are more
diffused than
p – bond.
(v)
The bond
energy of heteroatomic diatomic molecules increases with the decrease in the electronegativities of the bonded atoms.
(vi) With increase in bond order, bond length
decreases and bond strength increases.
(vii) The first ionization energies of the elements
rise steadily with increasing atomic number from top to bottom in a group.
(viii) A double bond is stronger than single bond
and a triple bond is weaker than a double bond.
(ix) The bonds formed between the elements having
electronegativity difference more than 1.7 are said
to be covalent in nature.
(x)
The
repulsive force between the two bonding pairs is less than that between the two
lone pairs.
(xi) The number of covalent bonds an atom can form
is related to the number of unpaired electrons it
has.
(xii) The rules which govern the filling of
electrons into the atomic orbitals also govern filling of electrons into the molecular
orbitals.
Q.4. Write
the Lewis structures for the following compounds: 11206033
(i)
HCN (ii) CCl4 (iii) CS2 (iv) H3N
A1F3
(v) H2SO4 (vi) N2O5 (viii) K2Cr2O7
(ix)
NH4OH (x) Ag(NH3)2NO3
Q. 5. Calculate the bond energy of H-Br. The bond energy of H-H is 436
kJ mol–1 and that of Br-Br is 193 kJ mol–1. 11206034
THERMOCHEMISTRY
Subjective
Q. 1 (a) What
is meant by Thermochemistry? Explain thermo-chemical reactions.
(Board 2004) 11207001
(b) What are the units used in thermochemical
measurements?
Q. 2 Explain
the terms: 11207002
(i) Spontaneous Reactions (ii)
Non-spontaneous Reactions
Q. 3 Explain
the following terms: (Board 2004) 11207003
(i) System
(ii) Surroundings (iii)
State (iv) State Function (v) Internal Energy
Q. 4 (a) Define and explain first law of
thermodynamics and derive its mathematical expression. 11207004
(b) Calculate an expression for pressure volume
work of a system.
Example 1: When
2.00 moles of H2 and 1.00 mole
of O2 at 100oC and 1 torr
pressure react to produce 2.00 moles of gaseous water, 484.5 kJ of energy are
evolved. What are the values of (a) DH and (b) DE for
the production of one mole of H2O(g)? 11207005
Q.5 Define
Enthalpy of a system. Derive the following expression by applying first law of
Thermodynamics. 11207006
(i) Expression for change in internal energy at
constant volume.
(ii) Expression for change in enthalpy at constant
pressure.
Q. 6 (a) Explain the following terms: 11207007
(i) Enthalpy
of a reaction. (Board 2013)
(ii) Enthalpy
of Formation.
(iii) Enthalpy
of Atomization.
(iv) Enthalpy
of Neutralization. (Board 2013)
(v) Enthalpy
of combustion.
(vi) Enthalpy
of Solution.
(b) How
is enthalpy of a reaction measured?
Example-2: Neutralization of 100 cm3 of 0.5 M NaOH at 25°C with 100 cm3 of 0.5 M HCl
at 25°C raised the temperature
of the reaction mixture to 28.5°C.
Find the enthalpy of neutralization. Specific heat of water = 4.2 J g-1 K-1. 11207008
Q.7 If
10.16g of graphite is burnt in a bomb calorimeter and the temperature rise
recorded is 3.87 K. Calculate enthalpy of combustion of graphite, if the heat
capacity of the calorimeter (bomb, water etc.) is 86.02 kJ K-1. 11207009
Q. 8 State and explain Hess’s
Law. (Board 2005,
2006) (Board 2014) 11207011
Q. 9 What is lattice energy? How does Born-Haber
cycle help to calculate the lattice energy of NaCl? 11207013
EXERCISE
Q. 1 Multiple Choice Questions. 11207014
Ans. See objective.
Q. 2 Fill in the blanks with suitable words: 11207015
(i) The
substance undergoing a physical or a chemical change forms a chemical
_________.
(ii) The change
in internal energy __________ be measured.
(iii) Solids
which have more than one crystalline forms possess __________ values of heats
of formation.
(iv) A process
is called ___________ if it takes place on its own without any external
assistance.
(v) A
___________ is a macroscopic property of a system which is ___________ of the
path adopted to bring that change.
Q. 3 Indicate true or false as the case may be. 11207016
(i) Total heat
content of a system is called enthalpy of the system.
(ii) Enthalpy
is a state function but internal energy is not.
(iii) The work
done by the system is given by the +ve sign.
(iv) Amount of
heat absorbed at constant volume is internal energy change.
(v) It is
necessary that a spontaneous reaction should be exothermic.
Q. 4 Define the following terms and give three examples
of each. 11207017
(i) system
(ii) surroundings
(iii) state function (Board 2014)
(iv) units of energy
(v) Exothermic reaction
(vi) Endothermic reaction
(vii) Internal energy of the system
(viii) Enthalpy of the system
Q. 5 Differentiate between the following: 11207018
(i) Internal energy and enthalpy
(ii) Internal energy change and enthalpy change
(iii) Exothermic and endothermic reaction
(b) Define the following enthalpies and give two
examples of each: 11207019
(i) standard
enthalpy of reaction (ii) standard enthalpy of combustion
(iii) standard
enthalpy of atomization (iv) standard enthalpy of solution
Q. 6 (a) What
are spontaneous and non spontaneous processes? Give examples.
(b) Explain
that burning of a candle is spontaneous process. 11207020
(c) Is
it true that a non-spontaneous process never happens in the universe? Explain
it:
Q. 7 (a) What is the first law of thermodynamics? How does it
explain that: 11207021
(i) qv = DE (ii) qp = DH
(b) How
will you differentiate between DE and
DH? Is it true that DH and DE
have the same values for the reactions taking place in the solution state?
Q. 8 (a) What
is the difference between heat and temperature? Write a mathematical
relationship between these two parameters. 11207022
(b) How
do you measure the heat of combustion of a substance by bomb calorimeter?
(b) See in chapter.
Q.9: Define heat of neutralization. When a dilute
solution of a strong acid is neutralized by a dilute solution of a strong base,
the heat of neutralization is found to be nearly the same in all the cases. How
do you account for this? 11207023
Q.10 (a) State
the laws of thermochemistry and show how are they based on the first law of
thermodynamics. 11207024
(b) What
is a thermochemical equation? Give three examples. What information do they
convey?
(c) Why
is it necessary to mention the physical states of reactants and products in a
thermochemical reaction? Apply Hess’s law to justify your answer.
Q.11 (a) Define
and explain Hess’s law of constant heat summation. Explain it with examples and
give its applications. 11207025
(b) Hess’s
law helps us to calculate the heats of those reactions, which cannot be normally
carried out in a laboratory. Explain it.
Q.12 (a) What
is lattice energy? How does Born-Haber cycle help to calculate the lattice energy
of NaCl? 11207026
(b) Justify that heat of formation of compound is
the sum of all the other enthalpies.
Q. 13 50cm3 of 1.0M HCl is mixed with 50cm3
of 1.0M NaOH in a glass calorimeter. The temperature of the resultant mixture
increases from 21.0oC to 27.5oC. Assume that calorimeter
losses of heat are negligible. Calculate the enthalpy change mole-1
for the reactions. The density of solution to be considered is 1gcm-3
and specific heat is 4.185Jg-1 K-1. 11207027
Q. 14 Hydrazine N2H4 is a
rocket fuel. It burns in O2 to give N2 and H2O.
11207028
N2 H4(l)
+ O2(g) 
N2(g) + 2H2O(g)
1.0g of N2H4 is
burned in a bomb calorimeter. An increase of temperature 3.51oC is
recorded. The heat capacity of calorimeter is 5.5kJ kg-1. Calculate
the quantity of heat evolved. Also calculate the heat of combustion of 1 mole
of N2H4.
Q. 15 Octane C8 H18 is a motor
fuel. 1.80g of a sample of octane is burned in a bomb calorimeter having heat capacity
11.66 kJK-1. The temperature of the calorimeter increases from 21.36oC
to 28.78oC. Calculate the heat of combustion for 1g of octane. Also
calculate the heat for 1 mole of octane. 11207029
Q.16 By applying, Hess’s law calculate the enthalpy
change for the formation of an aqueous solution of NH4Cl from NH3
gas and HCl gas. The results for the various reactions and pressures are as
follows: 11207030
(i) NH3(g)
+ aq → NH3(aq) ∆H
= -35.16 kJ / mol
(ii) HCl(g)
+ aq → HCl(aq) ∆H
= -72.41 kJ / mol
(iii) NH3(aq)
+ HCl(aq) → NH4Cl(aq) ∆H = -51.48 kJ / mol
Q.
17 Calculate the heat of formation of
ethyl alcohol from the following information.
11207031
(i) Heat of formation of CO2
is –393.7 kJ/mole
(ii) Heat of formation of H2O
is –285.8 kJ/mole
(iii) Heat of combustion of ethyl alcohol is –1367
kJ/mole
Q. 18 If the heats of combustion of C2H4, H2 and C2H6 are –337.2, –68.3
and – 372.8 k calories respectively, then calculate the heat of the following
reaction. 11207032
C2H4(g) + H2(g) ¾®
C2H6(g)
Q. 19 Graphite and diamond are two forms of carbon.
The enthalpy of combustion of graphite at 25°C
is – 393.51 kJ/mol-1 and that of diamond is – 395.41 kJ
mol-1.
What is the enthalpy change of the
process? Graphite ®
Diamond (at the same temperature)? 11207033
Q. 20 What is the meaning of the
term enthalpy of ionization? If the heat of neutralization of HCl and NaOH is
–57.3 kJ/mol-1 and heat of neutralization of CH3OOH is –55.2 kJ/mol-1. Calculate the enthalpy of ionization
of CH3COOH. 07(013)
Q. 21 (a) Explain
what is meant by the following terms. 11207034
(i) Atomization
energy. (ii) Lattice energy.
(b) Draw
a complete, fully labeled Born Haber cycle for the formation of potassium
bromide.
(c) Using
the information given in the table below, calculate the lattice energy of
potassium bromide.
Reactions: DH
kJ mol-1
K(s) + ½ Br2(l)
®K+
Br(s) -392
K(s)
® K(g) +
90
K(g) ®
K+ (g) + e– +
420
½ Br2(l)
® Br-1 (g) + 112
Br
(g) + e- ® Br-1 (g) -342
CHEMICAL EQUILIBRIUM
Subjective
Q1. Differentiate between irreversible and reversible reactions. 11208001
Q2. What is
meant by state of chemical equilibrium for a reversible reaction? 11208002
Q3. (a)
State law of mass action. 11208003
(b)
Define “active mass”.
(c)
Derive an expression for equilibrium constant using law of mass action.
Q4: Discuss the units of equilibrium constant in
different conditions. 11208004
Kc
and its units:
Q5. Discuss applications of Kc. / How the
value of Kc helps in detecting the direction and extent of a
chemical reaction? 11208005
Q6. State Le-Chatelier’s principle. Discuss its
applications. (Board 2005) 11208006
Q:7 Write a note on synthesis of Ammonia by Haber’s
process. 11208007
Q:8 How is SO3
produced industrially by applying Le-chatelier’s principle? 11208008
Q:9 What is
meant by Ionic product of water? Give its mathematical representation. 11208009
Q:10
(a)
What
is common ion effect? How does it affect the solubility of a substance. (Board 2014)
(b)
Give
applications of common ion effect. 11208010
Q:11(a) Define buffer solution. Discuss its
composition. 11208011
(b) Why
do we need a buffer solution?
(c) What is the mechanism of
buffer action? / How do the buffers act?
Q:12 Write a detailed note on Buffer Capacity. 11208012
Q:13 (a)
Define solubility. Discuss effect of common ion on solubility. 11208013
(b) What is meant by solubility
product. Give examples.
(c) Give applications of solubility
product.
EXERCISE
Q. 1 Select the most suitable answer. 11208014
Ans. See in objective.
Q. 2 Fill in the blanks. 11208015
(i) Law
of Mass action states that the _________ at which a reaction proceeds is directly
proportional to the product of the active masses of the _________.
(ii) In
an exothermic reversible reaction _________ temperature will shift the
equilibrium towards the forward direction.
(iii) In
a gas phase reaction, if the number of moles of reactants are equal to the
number of moles of the products, Kc of the reaction is _________ to the Kp.
(iv) The
equilibrium constant for the reaction 2O3 ® 3O2 is 1055 at 25°C, it
tells that ozone is _________ at room temperature.
(v) Buffer
solution is prepared by mixing together a weak base and its salt with _________
or a weak acid and its salt with __________.
Q. 3 Indicate true or false as the case may be: 11208016
(i) When
a reversible reaction attains equilibrium both reactants and products are present
in reaction mixture.
(ii) The
Kc of the reaction A + B 
C + D is given by Kc = 
therefore it is
assumed that [A] = [B] = [C] = [D].
(iii) Ionic
product Kw of water at 25°C is 10-14 mole2 dm-6, and is represented by an expression Kw = [H+] [OH-]
= 10-14 mole2 dm-6.
(iv) A
catalyst is a compound which increases the speed of the reaction and
consequently increases the yield of the product.
(v) AgCl
is a sparingly soluble ionic solid in water. Its solution produces excess of Ag+ and Cl- ions.
Q. 4 (a) Explain
the terms “reversible reaction” and “state of equilibrium”. 11208017
(b) Define
and explain the Law of Mass action and derive the expression for the
equilibrium constant (Kc).
(c) Write Kc for the following reactions.
(i) 
+ 
+ 
(ii) 
+ 
+ Ag(s)
(iii) 
+ 
2NO(g)
(iv) 
+ 
4NO(g) + 6H2O(g)
(v) 
+ 
Q. 5 (a) Reversible
reactions attain the position of equilibrium which is dynamic in nature and not
static. Explain it. 11208018
(b) Why do the rates of forward reactions
slow down when a reversible reaction approaches the equilibrium stage?
Q. 6 When a graph is plotted between time on x-axis
and the concentrations of reactants and products on y-axis for a reversible
reaction, the curves become parallel to time axis at a certain stage. 11208019
(a) At what stage the curves become parallel?
(b) Before the curves become parallel, the
steepness of curves falls? Give reasons.
(c) The rate of decrease of concentration of any
of the reactants and rate of increase of concentrations
of any of products may or may not be equal, for various types of reactions, before the equilibrium time. Explain it.
Q. 7 (a) Write
down the relationship of different types of equilibrium constants i.e. Kc, and Kp, for the following
general reaction aA + bB cC + dD. 11208020
(b) Decide
the comparative magnitudes of Kc and Kp, for the
following reversible reactions (i) Ammonia synthesis (ii) Dissociation of PCl5.
Q. 8 (a) Write
down Kc for the
following reversible reactions. Suppose that the volume of reaction mixture in
all cases is V dm3 at equilibrium stage. 11208021
(i) CH3COOH+CH3CH2OH 
CH3COOC2H5 + H2O
(ii) H2 + l2 
2Hl
(iii) 2Hl 
H2 + l2
(iv) PCl5 PCl3 + Cl2
(v) N2 + 3H2 2NH3
(b) How
do you explain that some of the reactions mentioned above are affected by
change in volume at equilibrium stage?
Q.9 Explain the following two applications of
equilibrium constant. Give examples.
(i) Direction of reaction (ii) Extent of reaction. 11208022
Q.10 . Explain
the following with reasons:
(a) The change of volume disturbs the
equilibrium position for some of the gaseous phase reactions but not the
equilibrium constant. 11208023
(b) The change of temperature disturbs both the
equilibrium position and the equilibrium constant of a reaction.
(c) The solubility of glucose in water is
increased by increasing the temperature.
Q.11 (a) What
is ionic product of water? How does this value vary with the change in
temperature? Is it true that its value increases 75 times when the temperature
of water is increased from 0°C to 100°C? 11208024
(b) What is the justification for the
increase of ionic product with temperature?
(c) How would you prove that at 25°C, 1dm3 of
water contains 10–7 moles of H3O+ and 10-7 moles of OH-?
Q. 12 (a) Define
pH and pOH. How are they
related with pKw? (Board 2014) 11208025
(b) What happens to the acidic and basic
properties of aqueous solutions when pH varies from zero to 14.
(c) Is
it true that the sum of pKa and pKb is always equal to 14
at all temperatures for any acid? If not why?
Q.13 (a) What
is Lowry Bronsted idea of acids and bases? Explain conjugate acid and bases.
11208026
(b) Acetic acid dissolves in water and
gives proton to water, but when dissolved in H2SO4, it accepts proton.
Discuss the role of acetic acid in both cases.
Q.14 In the equilibrium 
+ 
DH = 90kJ/mole. 11208027
What is the effect on?
(a) The position of equilibrium
(b) Equilibrium constant if:
(i) temperature is increased (ii) volume
of the container is decreased
(iii) catalyst is added (iv) chlorine
is added
Explain your
answer.
Q. 15 Synthesis of ammonia by Haber’s process is an
exothermic reaction. 11208028
N2(g)
+ 3H2(g) 
2NH3(g)
(a) What should be the possible effect of
change of temperature at equilibrium stage?
(b) How does the change of pressure or volume
shifts the equilibrium position of this reaction?
(c) What is the role of catalyst in this
reaction?
(d) What happens to equilibrium position
of this reaction if NH3 is removed from the reaction vessel time to
time?
Q. 16 Sulphuric acid is a king of chemicals. It is
produced by the burning of SO2 to SO3 through an exothermic reversible process. 11208029
(a) Write the balanced reversible
reaction.
(b) What is the effect of pressure change
on this reaction?
(c) Reaction is exothermic but still
temperature of 400 - 500°C is required to increase the yield of SO3. Give reasons.
Q. 17 (a) What are buffer solutions? Why do we need them in daily life?
11208030
(b) How does the mixture of sodium
acetate and acetic acid give us acidic buffer?
(c) Explain that a mixture of NH4OH and NH4Cl gives us the basic buffer?
(d) How
do you justify that the greater quantity of CH3COONa in acetic acid
decreases the dissociating power of acetic acid and so the pH increases.
(e) Explain the term buffer capacity.
Q.18 (a) What
is solubility product? Derive the solubility product expression for sparingly
soluble compounds, AgCl, Ag2CrO4, and PbCl2. 11208031
(b) How do you determine the solubility
product of a substance when its solubility is provided in grams/100 grams of
water?
(c) How do you calculate the solubility
of a substance from the value of solubility product?
Q.19 Kc for the reaction 2Hl 
H2 + l2 is 0.016 at 520°C. The equilibrium mixture
contains [Hl] = 0.08 M, [H2] = 0.01 M and
[l2] = 0.01 M. To
this mixture more Hl is added so that its new concentration is 0.096 M. What
will be the concentration of [Hl], [H2] and [l2], when equilibrium is re-established? 11208032
Q. 20 The equilibrium constant for the reaction
between acetic acid and ethyl alcohol is 4.0. A mixture of 3 moles of acetic
acid and 1 mole of ethyl alcohol is allowed to come to equilibrium stage.
Calculate the amount of ethyl acetate at equilibrium in number of moles and
grams. Also calculate the masses of reactants left behind. 11208033
Q.21 Study the equilibrium H2O(g) + CO(g) H2(g) + CO2(g). 11208034
(a) Write an expression of Kp.
(b) When 1.00 mole of steam and 1.00 mole
of carbon monoxide are allowed to reach equilibrium, 33.3% of the equilibrium
mixture is hydrogen. Calculate the value of Kp. State the units of Kp.
Q.22 Calculate the pH of 11208035
(a) 10-4 mole dm-3 of HCl
(b) 10-4 mole dm-3 of Ba(OH)2
(c) 1.0 mole dm-3 of H2
X which is only 50% dissociated
(d) 1.0 mole dm-3 of NH4OH which is 1% dissociated
Q.23 (a) Benzoic
acid, C6H5COOH, is a weak mono-basic acid (Ka = 6.4 ´ 10-5 mole dm-3). What is the
pH of a solution containing 7.2g of sodium benzoate in one dm3 of 0.02 mole dm-3 benzoic acid? (Board 2015) 11208036
(b) A buffer solution has been prepared
by mixing 0.2 M CH3COONa and 0.5
M CH3COOH in 1 dm3
of solution. Calculate the pH of solution. pKa of acid = 4.74 at 25oC.
How the values of pH will change by adding 0.1 mole of NaOH and 0.1 mole of HCl
separately.
Q. 24 The solubility of CaF2 in water at 25°C is found to be 2.05 ´ 10-4 mole dm-3. What is the
value of Ksp at this
temperature? 11208037
Q. 25 The solubility product of Ag2CrO4 is 2.6 ´ 10-2 at 25°C. Calculate the solubility of the compound.
SOLUTIONS
Subjective
Q. 1 (a) Define
the following terms. 11209001
(i) Solution (ii) Solute (iii) Solvent
(iv)
Dilute Solution (v) Concentrated
solution (vi) Binary Solution
(b) What
are the concentration units of solutions? Explain in detail.
Q.2 (a) Calculate the molarity of glucose solution when
9g of it are dissolved in 250 cm3 of solution. 11209002
(b) Calculate
the mass of urea in 100g of H2O in 0.3 molal
solution.
(c) 250
cm3 of 0.2 molar K2SO4
solution is mixed with 250 cm3 of 0.2 molar
KCl solution. Calculate the molar concentration of K+1
ions in the solution.
(d) Calculate
the concentration of a solution in terms of molality kg-1,
which is obtained by mixing 250g of 20% solution of NaCl with 200g of 40%
solution of NaCl.
Q.3 How various
concentration units of solutions are inter-converted? Explain it by giving
examples. 11209003
Q.4 You are
provided with 80% H2SO4 w/w having density 1.8g/cm3. How much
volume of this H2SO4 sample is required to obtain one dm3
of 20% H2SO4
w/w which has a density of 1.25g cm-3. 11209004
Q. 5 An aqueous solution of sucrose has been
labeled as 1 molal. Find the mole fraction of the solute and the solvent. 11209005
Q.6 5g of NaCl are dissolved in 1000g of water. The
density of resulting solution is 0.997 g/cm3.
Calculate molarity, molality and mole fraction of this solution. Assume that
the volume of the solution is equal to that of solvent. 11209006
Q. 7 (a) What
are the common types of solution? 11209007
(b) Discuss
in detail the following:
(i) Solution
of solids in liquids (ii) Solutions of liquids in liquids
Q. 8 (a) Compare
ideal and non-ideal solutions. 11209008
(b) What
is Raoult’s Law? Give its different statements.
(c) Explain
lowering of vapour pressure.
(d) The
vapour pressure of water at 30°C is
28.4 torr. Calculate the vapour pressure of a solution containing 70g of cane
sugar (C12H22O11) in 1000g of water at the same
temperature. Also calculate the lowering of vapour pressure.
Q. 9 Explain in detail Raoult’s Law when both
components are volatile. 11209009
Q10 (a) Define Binary mixtures and give its types. 11209010
(b) What is ideal solution? Explain
fractional distillation of ideal mixture of two liquids.
(c) What are non-ideal solutions? Explain clearly
positive and negative deviations.
Q.11 (a) What is solubility? How is it determined? 11209011
(b) What are solubility curves? Give their
different types.
(c) Write a note on Fractional
Crystallization.
Q.12 (a) What
are colligative properties of solutions? Give their different types. 11209012
(b) Why
some of the properties are called Colligative? How Kb
and Kf values are calculated?
Q.13 (a) What is lowering of vapour pressure?
Derive an equation for calculating molecular mass of solute from lowering of
vapour pressure. 11209013
(b) Pure benzene has a vapour pressure of
122.0 torr at 32 °C. When 20g of a
non-volatile solute were dissolved in 300g of benzene, a vapour pressure of
120 torr was observed. Calculate the molecular mass of the solute. The
molecular mass of benzene is 78.1.
Q. 14 (a) Define
elevation in boiling point and explain it.
(Board 2010,2015) 11209014
(b) How is boiling point elevation measured
by Landsberger’s method? (Board 2015)
(c) The
boiling point of water is 99.725°C. To
a sample of 600g of water are added 24.0g of a solute having molecular mass of
58g mole-1, to form a solution. Calculate the
boiling point of the solution.
Q.15 The boiling point of a solution containing
0.2g of a substance A in 20.0g of ether (molar mass = 74) is 0.17K higher than
that of pure ether. Calculate the molar mass of A. Molal boiling point constant
of ether is 2.16K. 11209015
Q. 16 3g of a non-volatile, non-electrolyte Solute
“X” are dissolved in 50g of ether (molar mass = 74.00) at 293K. The vapour
pressure of ether falls from 442 torr to 426 torr under these conditions.
Calculate the molar mass of solute “X”. 11209016
Q. 17. (a) Define
freezing point depression in solutions. Explain it in detail. 11209017
(b) How is depression in freezing point
measured? Give experimental detail.
(c) The freezing point of pure camphor is 178.4oC.
Find the freezing point of a solution containing 2.0 g of a non-volatile
compound, having molecular mass 140, in 40g of camphor. The molal freezing
point constant of camphor is 37.7oC kg mol-1. 11209018
Q.18 4.675g of a compound with empirical formula
C3H3O were dissolved in 212.5g of pure benzene. The
freezing point of solution was found 1.02°C
less than that of pure benzene. The molal freezing point constant of benzene is
5.1°C. Calculate (i) the
relative molar mass and (ii) the molecular formula of the compound. 11209019
Q.19 (a) What
is heat of solution? Explain energetics of solution. 11209020
(b) Define
and explain hydration process. Also discuss the extent of hydration.
(c) What
are hydrates and water molecules of crystallization?
Q. 20 (a) Explain
the steps involved in the process of hydration. Also give hydration energies of
common ions. 11209021
(b) What
is hydrolysis? Explain hydrolytic reactions with respect to hydrolysis.
EXERCISE
Q.1 Choose the correct answer for the given ones:
Ans. See
in objective.
Q.2 Fill in the blanks with suitable words. 11209022
(i) Number of molecules of sugar in 1 dm3
of 1 M sugar solution is __________.
(ii) 100 g of a 10% aqueous solution of NaOH
contains 10g of NaOH in _____ g of water.
(iii) When an azeotropic mixture is distilled, its
__________ remains constant.
(iv) The molal freezing point constant is also
known as _________ constant.
(v) The boiling point of an azeotropic solution
of two liquids is lower than either of them because the solution shows
__________ from Raoult’s law.
(vi) Among equimolal aqueous solutions of NaCl,
BaCl2
and FeCl3
the maximum depression in freezing point is shown by __________ solution.
(vii) A solution of ethanol in water shows
__________ deviations and gives azeotropic solution with __________ boiling
point than other components.
(viii) Colligative properties are used to calculate
__________ of a compound.
(ix) The hydration energy of Br-
ion is __________ than that of F- ion.
(x) The aqueous solution of NH4Cl is
__________ while that of Na2SO4 is __________.
Q.3 Indicate True or False from the given
statements. 11209023
(i) At a definite temperature the amount of a
solute in a given saturated solution is fixed.
(ii) Polar solvents readily dissolve non-polar
covalent compounds.
(iii) The solubility of a substance decreases with
increase in temperature, if the heat of a solution is negative.
(iv) The rate of evaporation of a liquid is
inversely proportional to the intermolecular forces of attraction.
(v) The molecular mass of an electrolyte
determined by lowering of vapour pressure is less than the theoretical
molecular mass.
(vi) Boiling point elevation is directly
proportional to the molality of the solution and inversely proportional to
boiling point of solvents.
(vii) All solutions containing 1g of non-volatile,
non-electrolyte solutes in some solvent will have the same freezing point.
(viii) The freezing point of a 0.05 molal solution of
a non-volatile, non-electrolyte in water is
– 0.93 °C
(ix) Hydration and hydrolysis are different
processes for Na2SO4.
(x) The hydration energy of an ion only depends
upon its charge.
Q. 4 Define and explain the following with one
example in each case. 11209024
(a) A homogeneous phase (b) A
concentrated solution
(c) A solution of solid in a solid (d) Consulate
temperature
(e) A non-ideal solution (f) Zeotropic solutions
(g) Heat of Hydration (h) Water
of Crystallization
(i) Azeotropic solution (j) Conjugate
solution
Q.5 (a) What
are the concentration units of solutions? Compare molar and molal solutions.
11209025
(b) One
has one molal solution of NaCl and one molal solution of glucose.
(i) Which solution has greater number of
particles of solute?
(ii) Which
solution has greater amount of the solvent?
(iii) How do we convert these concentrations
into weight by weight percentage?
Q.6 Explain the following with reasons. 11209026
(i) The concentration in terms of molality is
independent of temperature but molarity depends upon temperature.
Q.7 What are non-ideal solutions? Discuss their
types and give three examples of each.
11209027
Q.8 (a) Explain fractional distillation. Justify
the two curves when composition is plotted against boiling point of solutions.
(b) The
solutions showing positive and negative deviations cannot be fractionally
distilled at their specific compositions. Explain it. 11209028
Q.9 (a) What
are azeotropic mixtures? Explain them with the help of graphs. 11209029
(b) Explain the effect of temperature on
phenol-water system.
Q.10 (a) What
are Colligative properties? Why are they called so? 11209030
(b) What
is the physical significance of Kb and Kf values of solvents?
Q.11 How do you explain that the lowering of
vapour pressure is a Colligative property? How do we measure the molar mass of
a non-volatile, non-electrolyte solute in a volatile solvent?
11209031
Q.12 How do you justify the given statements. 11209032
Q.13 What is Raoult’s Law? Give
its three statements. How does this help in understanding the ideality of a
solution. 11209033
Q.14 Give Graphical explanation for elevation of
boiling point of a solution.
Describe
one method to determine the elevation in Boiling point of a solution. 11209034
Q.15 Freezing points of
solutions are depressed when non-volatile solutes are present in volatile solvent.
Justify it. Plot a graph to elaborate your answer. Also give one method to
record the depression of freezing point of a solution. 11209035
Q.16 Discuss the energetics of solution. Justify the heat of solutions as exothermic and endothermic properties.
ELECTROCHEMISTRY
Subjective
Q. 1 (a) What is Electrochemistry? Define
cell. Also write different types of cells. 11210001
(b) What
is conductor? Describe its types.
Q. 2 (a) What is ionization? Give examples. 11210002
(b) What
is electrolytic cell? Explain the working of electrolytic cell.
(c) What
is voltaic or galvanic cell? Give its construction and working with cell
reaction.
(d) Explain
the Reversibility in Voltaic cell.
Q. 3 Explain fully the electrolysis of following
substances with the help of diagrams and equations: 11210003
(a) Fused
sodium chloride.
(b) Concentrated
aqueous solution of sodium chloride.
(c) Aqueous
solution of salt
(d) Explain
processes of some industrial importance. OR
Outline the important industrial
applications of electrolysis. Write the electrochemical reactions involved
there in. (Board 2014)
Q. 4 (a) What
is electromotive force or EMF of a cell? Explain it. 11210004
(b) What
is electrode potential?
(c) Write
a note on standard hydrogen electrode (SHE).
Q. 5 How will you use the Standard Hydrogen
Electrode to measure the standard electrode potential of Zinc Electrode? (Board 2014) 11210005
Q. 6 (a) Define
and explain electro-chemical series or e.m.f. series. 11210006
(b) What
are the applications of electrochemical series?
Q. 7 (a) What are the two
basic types of modern cells? 11210007
(b) What
is lead accumulator? Explain it with reference to discharging and recharging.
(c) Write
a note on the following:
(1) Alkaline
battery.
(2) Silver
oxide battery.
(3) Nickel
cadmium cell. (Board 2014)
(4) Fuel
cells.
Q. 8 (a) What
is Oxidation Number? Describe the rules for allocating oxidation numbers to the
elements. Apply these rules to allocate oxidation numbers to the following
elements.
(i) Cr
in K2 Cr2
O7 and CrO4-2 (ii) Mn
in MnO4-1 11210008
(iii) S
in H2SO4
and SO4-2 (iv) N in HNO3
and NH4+1
(b) How oxidation number of an
element in a compound is found?
Q. 9 How are redox equations balanced by oxidation
number method? 11210009
Q.10 Balance the following equation by Oxidation number method. 11210010
(a) K2Cr2O7 + HCl ¾¾®
KCl + CrCl3 + Cl2 + H2O
(b) Zn
+ HNO3 ¾¾®
Zn(NO3)2 + NO + H2O
(c) NaOH
+ Br2 ¾¾®
NaBr + NaBrO3 + H2O
Q.11 (a) Describe
the general rules for balancing a redox equation by oxidation number method. 11210011
EXERCISE
Q. 1 Multiple choice questions. For each question
there are four possible answers a, b, c and d choose
the one you consider correct. 11210012
Ans. See in Objective.
Q. 2 Fill in the Blanks: 11210013
(i) The oxidation number of O-atom is
__________ in OF2 and is __________ in H2O2.
(ii) Conductivity of metallic conductors is due
to the flow of __________ while that of electrolytes is due to flow of
__________.
(iii) Reaction taking place at the __________ is
termed as oxidation and at the __________ is called as reduction.
(iv) __________ is set up when a metal is dipped in
its own ions.
(v) Cu metal __________ the Cu-cathode when
electrolysis is performed for CuSO4 solution with Cu-cathode.
(vi) The reduction potential of Zn is ______ volts
and its oxidation potential is __________ volts.
(vii) In a fuel cell __________ react together in
the presence of __________
Q. 3 Mark the following statements as true or
false. 11210014
(i) In electrolytic conduction, electrons flow
through the electrolyte.
(ii) In the process of electrolysis, the
electrons in the external circuit flow from cathode to anode.
(iii) Sugar is a non-electrolyte in solid form and
when dissolved in water will allow the passage of an electric current.
(iv) A metal will only allow the passage of an
electric current when it is in cold state.
(v) The electrolytic products of aqueous copper
(II) chloride solution are copper and chlorine.
(vi) Zinc can displace iron from its solution.
(vii) SHE acts as cathode when connected with
Cu-electrode.
(viii) A voltaic cell produces electrical energy at
the expense of chemical energy.
(ix) Lead storage battery is not a reversible
cell.
(x) Cr change its oxidation number when K2Cr2O7
is reacted with HCl.
Q. 3 True/False: 11210015
Q. 4 (a) Explain
the term oxidation number with examples. 11210016
(b) Describe
the rules used for the calculation of oxidation number of an element in
molecules and ions giving examples.
(c) Calculate
the oxidation number of chromium in the following compounds.
(i) CrCl3 (ii) Cr2(SO4)3 (iii) K2CrO4 (iv) K2Cr2O7
(v) CrO3 (vi) Cr2O3 (vii) Cr2O-27
Q. 4 A, b and c; Ans. is given. 11210017
(i) Ca (ClO3)2
(ii) Na2CO3 (iii) Na3 PO4
(iv) HNO3
(v) Cr2(SO4)3 (vi)
HPO3 (vii)
K2MnO4.
Q. 5 (a) Describe
the general rules for balancing a redox equation by oxidation number method. 11210018
(b) Balance
the following equations by oxidation number method.
(i) Cu + HNO3 ¾® Cu(NO3)2 + NO2 + H2O
(ii) Zn + HNO4 ¾® Zn(NO3)2 + NO + H2O
(iii) Br2 + NaOH ¾® NaBr + NaBrO3 + H2O
(iv) MnO2 + HCl ¾® MnCl2 + H2O + Cl2
(v) FeSO4+K2Cr2O7+H2SO4 ® Fe2(SO4)3+Cr2(SO4)3+H2O+K2SO4
(vi) HNO3 + HI ¾® NO + H2O + I2
(vii) Cu + H2SO4 ¾® CuSO4 + SO2 + H2O
(viii) HI + H2SO4 ¾® I2 + SO2 + H2O
(ix) NaCl + H2SO4 + MnO2 ¾® Na2SO4 + MnSO4 + Cl2 + H2O
Q. 6 (a) Describe the general rules for
balancing a redox equation by Ion electron method.
(b) Balance
the following equations by Ion electron method: 11210019
(i) Fe+3
+ Sn+2 ¾® Fe+2 + Sn+4
(ii) MnO4-1 + C2 O4-2 ¾® Mn+2
+ CO2
(iii) Cr2O7-2 + Cl-1 ¾® 2Cr+3 + 3Cl2
(iv) Cu + NO3-1 ¾® Cu+2
+ 2NO2
(v) 
+ Fe+2 ¾® Cr+3 + Fe+3 (Acidic media)
(vi) S2O3-2 + OCl-1 ¾® Cl-1 + S4O6-2 (Acidic
media)
(vii) IO3-1 + AsO3-3 ¾® I-1 + AsO4-3 (Acidic
media)
(viii) Cr+3 + BiO3-1 ¾® Cr2O7-2 + 3Bi+3 (Acidic
media)
(ix) AsO3-3 + Cr2O7-2 ¾® AsO4-3 + 2Cr+3 (Acidic
media)
(x) CN-1 + MnO4-1 ¾® CNO-1 + MnO2 (Basic media)
Q. 7 See in chapter. 11210020
Q. 8 Single Electrode Potential. 11210021
Q. 9 Application of Electrolysis. 11210022
Q. 10 Describe the construction and working of standard
Hydrogen Electrode (SHE).11210023
Q. 11 Fe+3
+ Ag ¾¾® Fe+2 + Ag + 11210024
Q. 12 Explain the difference between ionization and electrolysis. 11210025
Q.13 Describe the
galvanic cell explaining the functions of redox reaction generates electric
current. 11210026
Q. 14 Write
comprehensive notes on: 11210027
(a) Spontaneity of
oxidation reduction reactions.
(b) Alkali, Silver oxide
and nickel cadmium batteries, fuel cells.
(c) Lead Accumulator, its
desirable and undesirable features.
Q. 15 Will the reaction be spontaneous for the
following set of half reactions? What will be the value of 
? 11210028
(i) Cr+3 + 3e-
¾¾®
Cr E°red
= - 0.74 Volt
(ii) MnO2 + 4H+
+ 2e¢ ¾¾®
Mn+2 +2H2O E° Red = + 1.28 Volt.
Q. 16 Explain
the following with reasons. 11210029
(a) A porous plate or a salt bridge is
not required in lead storage cell.
CHEMICAL
KINETICS OR
REACTION KINETICS
Subjective
Q. 1 (a) Define order of reaction. Discuss
various types of reactions on the basis of their order. 11211001
(b) What is half-life period?
Q. 2 (a) What
are the physical methods of determination of rate of a reaction? 11211002
(b) Write
a note on chemical method for the determination of rate of a reaction.
Q. 3 Write a detailed note on activation energy.
Explain with the help of graphs. 11211003
Q. 4 How order of a reaction
can be determined by half-life method? 11211004
Q:5 What factors affect the rate of a chemical
reaction? Discuss. 11211005
Q. 6 How does Arrhenius equation help us to
calculate the energy of activation? 11211006
Q:7 Discuss
different types of catalysis in detail. / Differentiate between Homogeneous and
Heterogeneous catalysis. (Board 2014) 11211007
Q:8 What is enzyme catalysis? Give its mechanism
and characteristics. 11211008
Enzyme Catalysis: (Board
2015)
EXERCISE
Q. 1 Multiple choice questions: 11211009
Ans. See in objective.
Q. 2 Fill in the blanks with suitable words. 11211010
(i). The rate of endothermic reaction _________
with the increase in temperature.
(ii). All radioactive disintegration reactions are
of __________ order.
(iii). For a fast reaction the rate constant is
relatively ________ and half life is ________.
(iv). The second order reaction becomes _________
if one of the reactants is in large excess.
(v). Arrhenius equation can be used to find out
_________ of a reaction.
Q. 3 Indicate TRUE or FALSE as the case may be : 11211011
(i). The half life of a first order reaction
increases with temperature.
(ii). The reactions having zero activation
energies are instantaneous.
(iii). A catalyst makes a reaction more exothermic.
(iv). There is difference between rate law and the
law of mass action.
(v). The order of reaction is strictly determined
by the stoichiometry of that balanced equation.
Q. 4 What is chemical kinetics? How do you
compare chemical kinetics with chemical equilibrium and thermodynamics? 11211012
Q. 5 The rate of a chemical reaction with respect
to products is written with positive sign, but with respect to reactants is
written with a negative sign. Explain it with reference to the following
hypothetical reaction. 11211013
aA
+ bB ¾¾® cC +
dD
Q.6 What are instantaneous and average rates? Is
it true that the instantaneous rate of a reaction at the beginning of the
reaction is greater than average rate and becomes far less than the average
rate near the completion of reaction? 11211014
Q.7 Differentiate between: 11211015
(i) Rate and rate constant of a reaction.
(ii) Homogeneous and Heterogeneous catalysis. (Board 2014)
(iii) Fast step and the rate determining step.
(iv) Enthalpy change of reaction and energy of
activation of reaction.
Q.8 Justify the following statements: 11211016
(i) Rate of chemical reaction is an ever changing
parameter under the given conditions.
(ii) The
reaction rate decreases every moment but rate constant “k” of the reaction is a
constant quantity under the given conditions.
Q.9 Explain that half-life method for
measurement of the order of reaction can help us to measure the order of even
those reactions, which have a fractional order. 11211017
Q.10 A
curve is obtained when a graph is plotted between time on x-axis and
concentration on y-axis. The measurement of the slopes of various points give
us the instantaneous rates of reaction. Explain with suitable examples. 11211018
Q.11 The rate determining step of a reaction is
found out from the mechanism of that reaction. Explain it with few examples. 11211019
Q.12 Discuss
the factors which influence the rates of chemical reactions. 11211020
Q.13 Explain the following facts about the
reaction. 11211021
2NO(g) + 2H2(g) ¾¾® 2H2O(g) + N2(g)
(i) The changing concentrations of reactants
change the rates of this reaction.
(ii) Individual orders with respect to NO and H2 can be measured.
(iii) The overall order can be evaluated
by keeping the concentration of one of the substances constant.
Q.14 The
collision frequency and the orientation of molecules are necessary conditions
for determining the proper rate of reaction. Justify the statement. 11211022
Q.15 How
does Arrhenius equation help us to calculate the energy of activation of a
reaction?
11211023
Q.16 Define
the following terms and give examples. 11211024
(i) Homogeneous
catalysis.
(ii) Heterogeneous
catalysis.
(iii) Activation
of a catalyst.
(iv) Auto-catalysis
(v) Catalytic
poisoning
(vi) Enzyme
catalysis
Q. 17 Briefly
describe the following with examples: 11211025
(i) Change of
physical state of a catalyst at the end of reaction.
(ii) A very
small amount of a catalyst may prove sufficient to carry out a reaction.
(iii) A
finely divided catalyst may prove more effective.
(iv) Equilibrium
constant of a reversible reaction is not changed in the presence of a catalyst.
(v) A catalyst
is specific in its action.
Q. 18 What
are enzymes? Give examples in which they act as catalyst. Mention the
characteristics of enzyme catalysis. 11211026
Q. 19 In
the reaction of NO and H2 it was observed that mixture of gases at 340.5
mm pressure was half changed in 102 seconds. In another experiment with an
initial pressure of 288 mm of Hg the reaction was half completed in 140 seconds.
Calculate the order of reaction.
11211027
Q. 20 A
study of chemical kinetics of a reaction 11211028
A + B ¾® products
gave
the following data at 25 °C. Calculate the rate law.
Q. 21 Some
reactions taking place around room temperature have activation energies around
50 kJ mol-1. 11211029
(i) What is the value of the factor 
at 25°C
(ii) Calculate the factor at 35°C and 45°C and note the increase in this factor for every
10°C rise in temperature.
(iii) Prove that for every 10°C rise of temperature, the factor doubles and so
rate constant also doubles.
Q.22 H2
and I2 react to produce HI. Following data for rate constant at
various temperatures (K) have been collected. 11211030
|
Temp (K) |
Rate constant (cm3 mol-1 s-1)
(k) |
|
500 550 600 650 700 |
6.814
´ 10-4 2.64
´ 10-2 0.56
´ 10o 7.31
´ 10o 66.67
´ 10o |
(i)
Plot a graph between 
on x-axis and log k on
the y-axis.
(ii) Measure the slope of this
straight line and calculate the energy for activation of this reaction.
