Thursday, October 20, 2022
Wednesday, October 19, 2022
Dalton’s Atomic Theory
• It fails to explain why atoms of different kinds should differ in mass and valency etc.
• The discovery of isotopes and isobars showed that atoms of same elements may have different atomic masses (isotopes) and atoms of different kinds may have same atomic masses (isobars).
• The discovery of various sub-aomic particles like X-rays, electrons, protons etc. during late 19th century lead to the idea that the atom was no longer an indivisible and smallest particle of the matter.
Thursday, October 13, 2022
Cbse Chemistry Class XI sample paper term -I 2022-23
Central board secondary examination 2022
TERM I 2022-23
CHEMISTRY M.Marks:70
General Instructions:
Read the following instructions carefully
a) There are 35 questions in this question paper with internal choice.
b) SECTION A consists of 18 multiple-choice questions carrying 1 mark each.
c) SECTION B consists of 7 very short answer questions carrying 2 marks each.
d) SECTION C consists of 5 short answer questions carrying 3 marks each.
e) SECTION D consists of 2 case- based questions carrying 4 marks each.
f) SECTION E consists of 3 long answer questions carrying 5 marks each.
g) All questions are compulsory. h) Use of log tables and calculators is not allowed
SECTION A
1. The number of significant figures in 2.004 is.
i 3 ii. 4 iii. 5 iv. 1
2. Photoelectric effect is shown by
a. Alkali metals b. Halogens c. Noble gases d. None
3. The number of sigma bonds in ethylene molecule is
a.4 b.6 c.5 d. 2
4. Isoelectronic species have same number of
a. electrons b. Protons c. Neutrons d. None
5. A light radiation of wavelength 400nm has a frequency of
a. 8x1014 b. 7.5x 1014 c. 6 d. none
6. Molarity depends on
a. Temperature
b. Pressure
c. both
d, none
7.An orbital can accommodate a maximum of 2 electrons with opposite spins. This is
a. Hund’s rule
b. Pauli’s exclusion Principle
‘c. Heisenberg’s Uncertainity Principle
d. none
8.The energy of bonding molecular orbital is ..........than antibonding molecular orbital.
a. greater
b. smaller
c. equal
d. none of these
9. In PCl5 molecule the bond angles are ..............
a. 1200 and 900
b. 900 and 1040
c. 1070
d. none
10. The shape of ammonia molecule is
a. Tetrahedral
b. Octahedral
c. Pyramidal
d. None
11. The properties which are independent of the amount of the substance contained in the system but depend upon the nature opf the substance only are called
a. Extensive properties
b. Intensive properties
c. State properties
d. none
12. In Open system there is an exchange of ................with the surroundings.
a.Matter and energy
b. Energy
c. Matter
d. None
13.For a spontaneous process the value of Gibb’s free energy is
a. Negative
b. Positive
c. Zero
d. None
14. The enthalpies of all elements in their standard states are
a. unity
b. zero
c. <0
d. different for each element
15. In a process, 701 J of heat is absorbed by a system and 394 J of work is done by the system. What is the change in internal energy for the process?
a. 305 J
b. 307 J
c. 404J
d. none
16. Sigma bond is .........than pie bond
a. stronger
b. weaker
c. equal in strength
d. none
17.Aufbau’s Principle is based on
a. n+l rule
b.2l+1 rule
c. n-l rule
d.none
18.Which of the following species will have the largest and the smallest size
Mg,Mg+2,Al,Al3+
a. Mg and Al3+
b. Al and Mg2+
c. Al and Mg
d. Mg2+ and Al3+
SECTION B
19State Heisenberg’s Uncertainity Principle . Write its mathematical expression.
20. Define Coordinate bond and give one example.
21. A golf ball has a mass of 40g and a speed of 45 m/s. If the speed can be measured within accuracy of 2%, calculate the uncertainity in the position.
22. Define Photoelectric Effect and write its expression.
23. Half filled and fully filled orbitals are more stable than any other electronic configuration. Why?
24. The ionization enthalpy of Nitrogen is more than that of Oxygen. Why?
25. Draw the Lewis dot structure of CO32- ion
SECTION C
26. What are the differences between sigma and pie bond?
27. Explain the geometry of NH3 molecule on the basis of VSEPR theory.;
28. Derive the relation between Cp and Cv
29. Calculate the energy needed to raise the temperature of 10 g of iron from 250C to 5000C if specific heat capacity of iron is 0.45 J/K/g.
30. Define
a. Ionization enthalpy
b. Electron gain enthalpy
c. Atomic radius
SECTION D
31. Molecular orbitals are formed by the linear combination of wave functions of atoms and Antibonding molecular orbitals are formed by the subtraction of wave functions of atoms. The no of molecular orbitals and antibonding molecular orbitals formed are equal to the number of atomic orbitals participated. The energy of antibonding molecular orbital is more than molecular bonding orbital.
a. What is the formula for bond order according to molecular orbital theory ?
b. How is bond order related to stability?
c. He2 molecule does not exist Why?
d. Hoe is bbopnd order rel;ated to bond length?
32. The first law of Thermodynamics is also called law of conservation of energy . During a physical or chemical change the energy may be converted from one form into another but the total energy remains constant.
a. Write the expression for I law of Thermodynamics.
b. What are the sign conventions for heat energy?
c. What is the sign conventions for work done?
d. What is the value of q for an adiabatic change ?
SECTION –E
33. (a) Arrange the following species in the increasing order of stability on the basis of Molecular Orbital Theory
O2, O2+,O-2
(b) Draw the molecular orbital diagram for O2-
34. What is Hybridisation ? Explain it in Ethylene molecule.
35. What are the frequency and wavelength of a photon emitted during a transition from n1=5 to n2 = 2 state in the hydrogen atom ?
ANSWER KEY
1. ii
2. a
3. c
4. a
5. b
6. a
7. b
8. b
9. a
10. c
11. b
12. a
13. a
14. b
15. b
16. a
17. a
18. a
19. Principle, Mathematical Expression
20. Correct Definition
21. 6.5x10-5m
22. Definition and expression
23. Symmetry , Exchange of energy
24. Due to half filled (stable) electronic configuration of Nitrogen
25. Correct structure
26. Any 3 differences
27. Pyramidal
28. Cp – Cv =R
29. 2.137 KJ
30. Definitions
31. (a) B.O=1/2 (Nb-Na)
(b)Bond order is directly proportional to stability
(c) B.O=0
(d) Bond order is inversely proportional to bond length
32.(a) ΔU=q+W
(b) Δq=+ve (heat gained by the system)
Δq = -ve (heat lost by the system)(
c) W= = +ve ( work done on the system )
W
= - ve ( work done by the system)
(d) q =0
33.(a) O2- <O2
<O2+
(b) Correct
Diagram
34. Definition and Explanation
35. ΔE =4.58x10-19 J, υ =6.91x1014 s-1 ,λ=434nm
Monday, October 3, 2022
Sunday, May 8, 2022
CBSE CLASS XI CHEMISTRY ANNUAL SYLLABUS 2022-23
CLASS–XI (THEORY) (2022-23)
Time:3Hours Total Marks70
|
S.NO |
UNIT |
PERIODS |
MARKS |
|
1 |
Some Basic Concepts of Chemistry |
18 |
7 |
|
2 |
Structure of Atom |
20 |
9 |
|
3 |
Classification of Elements and Periodicity in Properties |
12 |
6 |
|
4 |
Chemical Bonding and Molecular Structure |
20 |
7 |
|
5 |
Chemical Thermodynamics |
23 |
9 |
|
6 |
Equilibrium |
20 |
7 |
|
7 |
Redox
Reactions |
9 |
4 |
|
8 |
Organic Chemistry: Some basic Principles and Techniques |
20 |
11 |
|
9 |
Hydrocarbons |
18 |
10 |
|
|
TOTAL |
160 |
70 |
Unit I: Some Basic Concepts
of Chemistry 18 Periods
General
Introduction: Importance and scope of Chemistry. Nature of matter, laws of chemical combination, Dalton's atomic
theory: concept of elements, atoms and molecules.
Atomic and molecular masses, mole concept and molar mass, percentage composition, empirical and molecular
formula, chemical reactions, stoichiometry and
calculations based on stoichiometry.
Unit II:
Structure of Atom 20 Periods
Discovery of Electron, Proton and Neutron,
atomic number, isotopes
and isobars. Thomson's model and its limitations.
Rutherford's model and its limitations, Bohr's model and its limitations, concept of shells and subshells, dual
nature of matter and light, de Broglie's
relationship, Heisenberg uncertainty principle, concept of orbitals, quantum numbers, shapes of s, p and d orbitals,
rules for filling electrons in orbitals - Aufbau principle, Pauli's exclusion principle and Hund's rule,
electronic configuration of atoms, stability
of half-filled and completely filled orbitals.
Unit III: Classification of Elements and Periodicity in Properties 12 Periods
Significance of classification, brief history of the development of periodic table, modern periodic law and the present form of periodic table, periodic trends in properties of elements -atomic radii, ionic radii, inert gas radii, Ionization enthalpy, electron gain enthalpy, electronegativity, valency. Nomenclature of elements with atomic number greater than 100.
Unit IV: Chemical Bonding and Molecular Structure 20 Periods
Valence electrons, ionic bond, covalent bond, bond parameters, Lewis’s structure, polar character of covalent bond, covalent character of ionic bond, valence bond theory, resonance, geometry of covalent molecules, VSEPR theory, concept of hybridization,involving s, p and d orbitals and shapes of some simple molecules, molecular orbital theory of homonuclear diatomic molecules (qualitative idea only), Hydrogen bond.
Unit VI: Chemical Thermodynamics 23 Periods
Concepts of System and types of systems,
surroundings, work, heat, energy, extensive and intensive
properties, state functions. First law of thermodynamics -internal energy and enthalpy, heat capacity and specific
heat, measurement of ΔU and ΔH, Hess's law of
constant heat summation, enthalpy of bond dissociation, combustion, formation, atomization, sublimation, phase transition, ionization, solution and dilution.
Second law of Thermodynamics
(brief introduction) Introduction of entropy as a state function, Gibb's energy change for spontaneous and non- spontaneous processes, criteria for equilibrium. Third
law of thermodynamics (brief
introduction).
Unit VII: Equilibrium 20 Periods
Equilibrium
in physical and chemical processes, dynamic nature of equilibrium, law of mass action, equilibrium constant, factors
affecting equilibrium - Le Chatelier's principle, ionic equilibrium- ionization of acids and bases, strong and
weak electrolytes, degree of ionization,
ionization of poly basic acids, acid strength, concept of pH, hydrolysis of
salts (elementary idea), buffer
solution, Henderson Equation, solubility product, common ion effect (with
illustrative examples).
Unit VIII: Redox Reactions 09 Periods
Concept of oxidation and reduction, redox reactions, oxidation number, balancing redox reactions, in terms of loss and gain of electrons and change in oxidation number, applications of redox reactions.
Unit XII: Organic Chemistry -Some Basic Principles and Techniques 20 Periods General introduction, methods of purification, qualitative and quantitative analysis, classification and IUPAC nomenclature of organic compounds. Electronic displacements in a covalent bond: inductive effect, electromeric effect, resonance and hyper conjugation. Homolytic and heterolytic fission of a covalent bond: free radicals, carbocations, carbanions, electrophiles and nucleophiles, types of organic reactions.
Unit XIII: Hydrocarbons 18 Periods
Classification of Hydrocarbons Aliphatic Hydrocarbons:
Alkanes
- Nomenclature, isomerism, conformation (ethane only), physical properties, chemical reactions including free radical
mechanism of halogenation, combustion and pyrolysis.
Alkenes - Nomenclature, the structure of double bond (ethene),
geometrical isomerism, physical
properties, methods of preparation, chemical reactions: addition of hydrogen, halogen,
water, hydrogen halides
(Markovnikov's addition and peroxide effect),
ozonolysis, oxidation, mechanism of electrophilic addition.
Alkynes - Nomenclature, the structure of triple bond (ethyne), physical properties, methods of preparation, chemical reactions: acidic character of alkynes, addition reaction of - hydrogen, halogens, hydrogen halides and water.
Aromatic Hydrocarbons:
Introduction, IUPAC nomenclature, benzene: resonance, aromaticity, chemical properties: mechanism of electrophilic substitution. Nitration,
sulphonation, halogenation, Friedel Craft's alkylation and acylation, directive
influence of the functional group in monosubstituted benzene. Carcinogenicity and toxicity.
PRACTICALS
3 HOURS/ 30 Marks
|
Evaluation Scheme for Examination |
Marks |
|
Volumetric Analysis |
08 |
|
Salt Analysis |
08 |
|
Content Based
Experiment |
06 |
|
Project Work |
04 |
|
Class record
and viva |
04 |
|
Total |
30 |
PRACTICAL SYLLABUS Total Periods: 60
Micro-chemical methods are available for several of the practical experiments, wherever possible such techniques should be used.
A. Basic Laboratory Techniques
1.
Cutting glass tube and glass rod
2.
Bending a glass tube
3.
Drawing out a glass
jet
4. Boring a cork
B. Characterization and Purification of Chemical Substances
1.
Determination of melting point of an organic compound.
2. Determination of boiling point of an organic compound.
3. Crystallization of impure sample of any one of the following: Alum, Copper Sulphate, Benzoic Acid.
C. Experiments based on pH
1. Any one of the following experiments:
• Determination of pH of some solutions obtained from fruit juices, solution of known and varied concentrations of acids, bases and salts using pH paper or universal indicator.
• comparing the pH of solutions of strong and weak acids of same concentration. Study the pH change in the titration of a strong base using universal indicator.
2. Study the pH change by common-ion in case of weak acids and weak bases.
D. Chemical Equilibrium
One of the following experiments:
1. Study the shift in equilibrium between ferric ions and thiocyanate ions by increasing/decreasing the concentration of either of the ions.
2. Study the shift in equilibrium between [Co(H2O)6]2+ and chloride ions by changing the concentration of either of the ions.
E. Quantitative Estimation
1. Using a mechanical balance/electronic balance.
2. Preparation of standard solution of Oxalic acid.
3. Determination of strength of a given solution of Sodium hydroxide by titrating it against standard solution of Oxalic acid.
4. Preparation of standard solution of Sodium carbonate.
5. Determination of strength of a given solution of hydrochloric acid by titrating it against standard Sodium Carbonate solution.
F. Qualitative Analysis
1. Determination of one anion and one cation in a given salt
Cation:
Pb2+, Cu2+ As3+, Aℓ3+, Fe3+, Mn2+, Zn2+, Ni2+, Ca2+, Sr2+, Ba2+, Mg2+, NH +
Anions:
|
G. PROJECTS
Scientific investigations involving laboratory testing and collecting information from other sources.
A few suggested Projects
· Checking the bacterial contamination in drinking water by testing sulphide ion
· Study of the methods of purification of water
· Testing the hardness, presence of Iron, Fluoride, Chloride, etc., depending upon the regional variation in drinking water and study of causes of presence of these ions above permissible limit (if any).
· Investigation of the foaming capacity of different washing soaps and the effect of addition of Sodium carbonate on it
· Study the acidity of different samples of tea leaves.
· Determination of the rate of evaporation of different liquids.
· Study the effect of acids and bases on the tensile strength of fibers.
· Study of acidity of fruit and vegetable juices.
Note: Any other investigatory project, which involves about 10 periods of work, can be chosen with the approval of the teacher.
PRACTICAL EXAMINATION FOR VISUALLY IMPAIRED STUDENTS
Note: Same Evaluation scheme and general guidelines for visually impaired students as given for Class XII may be followed.
A. List of apparatus for identification for assessment in practical (All experiments)
Beaker, tripod stand, wire gauze, glass rod, funnel, filter paper, Bunsen burner, test-tube, test-tube stand, dropper, test tube holder, ignition tube, china dish, tongs, standard flask, pipette, burette, conical flask, clamp stand, dropper, wash bottle
• Odour detection in qualitative analysis
•
Procedure/Setup of the apparatus
B. List of Experiments A. Characterization and Purification of Chemical Substances
1. Crystallization of an impure sample of any one of the following: copper sulphate, benzoic acid
C. Experiments based on pH
1. Determination of pH of some solutions obtained from fruit juices, solutions of known and varied concentrations of acids, bases and salts using pH paper
2. Comparing the pH of solutions of strong and weak acids of same concentration.
D. Chemical Equilibrium
1. Study the shift in equilibrium between ferric ions and thiocyanate ions by increasing/decreasing the concentration of either ions.
2. Study the shift in equilibrium between [Co(H2O)6]2+ and chloride ions by changing the concentration of either of the ions.
E. Quantitative estimation
1. Preparation of standard solution of oxalic acid.
2. Determination of molarity of a given solution of sodium hydroxide by titrating it against standard solution of oxalic acid.
F. Qualitative Analysis1. Determination of one anion and one cation in a given salt
2. Cations - NH4 +
Anions – (CO3)2-, S2-, (SO3)2-, Cl-, CH3COO-
(Note: insoluble salts excluded)
3. Detection of Nitrogen in the given organic compound.
4. Detection of Halogen in the given organic compound.
Note: The above practical may be carried out in an experiential manner rather than recording observations.
PRESCRIBED BOOKS:
1. Chemistry Part – I, Class-XI, Published by NCERT.
2. Chemistry Part – II, Class-XI, Published by NCERT.
3. Laboratory Manual of Chemistry, Class XI Published by NCERT
4. Other related books and manuals of NCERT including multimedia and online sources
Note:The content indicated in NCERT textbooks as excluded for the year 2022-23 is not to be tested by schools.