iii - KopyKitab...and Hund s rule, electronic confi guration of atoms, stability of half fi lled and completely fi lled orbitals. Unit III : Classifi cation of Elements and Periodicity

iii

Unit I : Some Basic Concepts of Chemistry 12 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 14 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, cle Broglie’s relationship, Heisenberg uncertainty principle, concept of orbitals, quantum numbers, shapes of s, p and d orbitals, rules for fi lling electrons in orbitals : Aufbau principle, Pauli’s exclusion principle and Hund’s rule, electronic confi guration of atoms, stability of half fi lled and completely fi lled orbitals.

Unit III : Classifi cation of Elements and Periodicity in Properties

08 Periods

Signifi cance of classifi cation, brief history of the development of periodic table, modern periodic law and the present from of periodic trends in properties of elements : atomic radii, inert gas fadii, lonization enthalpy, electron gain enthalpy, electronegativity, valency. Nomenclature of elements with atomic number greater than 100.

Unit IV : Chemical Bonding and Molecular structure 14 Periods

Valence electrons, ionic bond, covalent bond, bond parameters, Lewis structure, polar character of covalent bond, covalent character of ionic bond, valence bond theory, resonance, geometry covalent molecules, VSEPR theory, concept of hybridization, involving s, p and d orbitals and shape of some simple molecules, molecular orbital theory of homonuclear diatomic molecules (qualitative idea only), hydrogen bond.

iv

Unit V : States of Matter: Gases and Liquids 12 Period

Three states of matter, intermolecular interactions, types of bonding, melting and boiling points, role of gas laws in elucidating the concept of the molecule, Boyle’s law, Charles; law, Gay Lussacs law, Avogadro’s law, ideal behaviour, empirical derivation of gas equation, Avogadro’s number, ideal gas equation. Deviation from ideal behaviour, liquefaction of gases, critical temperature, kinetic energy and molecular speeds (elementary idea), Liquid State- vapour pressure, viscosity and surface tension (qualitative idea only, no mathematical derivations)

Unit VI : Chemical Thermodynamics 14 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 specifi c heat, measurement of U and H, 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 14 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, Henderson Equation, hydrolysis of salts (elementary idea), buffer solution, solubility product, common ion effect (with illustrative examples).

Unit VIII: Redox Reactions 06 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.

v

Unit IX : Hydrogen 08 Periods

Position of hydrogen in periodic table, occurrence, isotopes, preparation, properties and uses of hydrogen, hydrides-ionic, covalent and interstitial; physical and chemical properties of water, heavy water, hydrogen peroxide-preparation, reactions and structure and use; hydrogen as a fuel.

Unit X : s-Block Elements (Alkali and Alkaline Earth Metals) 10 Periods

Group 1 and Group 2 Elements : General introduction, electronic confi guration, occurrence, anomalous properties of the fi rst element of each group, diagonal relationship, trends in the variation of properties (such as ionization enthalpy, atomic and ionic radii), trends in chemical reactivity with oxygen, water, hydrogen and halogens, uses.

Preparation and Properties of Some Important Compounds :

Sodium Carbonate, Sodium Chloride, Sodium Hydroxide and Sodium Hydrogencarbonate, Biological importance of Sodium and Potassium.

Calcium Oxide and Calcium Carbonate and their industrial uses, biological importance of Magnesium and Calcium.

Unit XI : Some p-Block Elements 14 Periods

General Introduction to p-Block Elements :

Group 13 Elements : General introduction, electronic confi guration, occurrence, variation of properties, oxidation states, trends in chemical reactivity, anomalous properties of fi rst element of the group, Boron-physical and chemical properties, some important compounds, Borax, Boric acid, Boron Hydrides, Aluminium: Reactions with acids and alkalies, uses.

Group 14 Elements : General introduction, electronic confi guration, occurrence, variation of properties, oxidation states, trends in chemical reactivity, anomalous behaviour of fi rst elements. Carbon-catenation, allotropic forms, physical and chemical properties; uses of some important compounds : oxides. Important compounds of Silicon and a few uses: Silicon Tetrachloride, Silicones, Silicates and Zeolites, their uses.

vi

Unit XII : Organic Chemistry -Some Basic Principles and Technique

14 Periods

General introduction, methods of purifi cation, 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, carbamons, electrophiles and nucleophiles, types of organic reactions.

Unit XIII: Hydrocarbons 12 Periods

Classifi cation of Hydrocarbons

Aliphatic Hydrocarbons : Alkanes : Nomenclature, isornerism, conformation (ethane only), physical properties, chemical reactions including free radical mechanism of halogenation, combustion and pyrolysis.

Alkenes : Nomenclature, structure of double bond (ethene), geometrical isornerism, physical properties, methods of preparation, chemical reactions: addition of hydrogen, halogen, water, hydrogen halides (Markownikov’s addition and peroxide effect), ozonolysis, oxidation, mechanism of electrophilic addition.

Alkynes : Nomenclature, 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 infl uence of functional group monosubstituted benzene. Carcinogenicity and toxicity.

Unit XIV : Environmental Chemistry 06Periods

Environmental pollution : Air, water and soil pollution, chemical in atmosphere, smog, major atmospheric, acid rain, ozone and its reactions, effects of depletion of ozone layer, greenhouse effect and global warning : pollution due to industrial wastes, green chemistry as an alternative tool for reducing pollution, strategies for control of environmental pollution.

vii

Chemistry (Code No. 043)Question Paper DesignClass : XI (2016-17)

Time : 3 Hours] [Max. Marks : 70

S. Typology V.S.A. S.A.-1 S.A.-II V.B.Q. L.A. Total % of Questions Marks Weightage

1. Remembering-(Knowledge based Since recall questions, to know specifi c facts, terms, concepts, principles, or theories, Identify, defi ne, or recite, information)

2. Understanding Comprehension-to be familir with meaning and to understand conceptually, interpret, compare, contrast, explain, paraphrase information.

3. Application (Use abstract information in concrete situation, to apply knowledge to new situations, Use given content to interpret a situation, provide an example, or solve a problem)

4. High order Thinking Skills (Analysis and Synthesis : Classify, compare, contrast, or differentiate between different pieces of iformation, organize and/or integrate uniqaue pieces of informtion from a variety of sources.)

5. Evalution A: (Appraise, Judge, and/or justify the value or worth of a decision or outcome, to to predict outcomes based on values)

Question Wise Break UpType of Ques. Mark per Ques. Total No. of Ques. Total Marks

VSA 1 5 05

SA-I 2 5 10

SA-II 3 12 36

VBQ 4 1 04

LA 5 3 15

Total 26 70

1. Internal Choice : There is no overall choice in the paper. However, there is an internal choice in one question of 2 marks weightage, one question of 3 marks weightage and all the three questions of 5 marks weightage.

2. The above template is only a sample. Suitable internal variations may be made for generating similar templates keeping the overall weightase to different form of questions and typology of questions same.

viii

Chemistry - XIIndex

S. No. Chapter Name Page No.

1. Some Basic Concepts 1

2. Atomic Structure 13

3. Classifi cation of Elements and Periodicity in Properties 31

4. Chemical Bonding and Molecular Structure 39

5. States of Matter 51

6. Thermodynomic 59

7. Equilibrium 69

8. Redox Reactions 83

9. Hydrogen 90

10. The s-Block Elements 98

11. The p-Block Elements 105

12. Organic Chemistry : Some basic Principles and Techniques 113

13. Hydro Carbons 133

14. Environmental Chemistry 146

1Some Basic Concepts Of Chemistry

Matter : Anything that has mass and occupy space. Precision : If refers to the closeness of various measurements for the same

quantity. Accuracy : It refers to the agreement of a particular value to the true value

of the result. Mass and weight : Mass of a substance is the amount of matter present

in it while weight is the force exerted by gravity on an object. The mass of a substance is constant whereas its weight may vary from one place to another due to change in gravity.

Volume : 1 L = 1 dm3 = 103cm3 = 10–3 m3

Temperature : K = °C + 273.15; =

Standard Temperature Pressure (STP) : 0°C (273.15 K) temperature and 1 atm pressure.

Normal Temperature Pressure (NTP) : 20°C (293.15 K) temperature and 1 atm pressure.

Standard Ambient Temperature Pressure (SATP) : 25°C (298.15 K) temperature and 1 atm pressure

Scientifi c Notation : Expressing a number in the form N × 10n, and N can vary b/w 1 to 10.

Signifi cant fi gures : These are meaningful digits which are known with certainty.

Laws of Chemical Combination : Law of Conservation of Mass (Antonie Lavoisier) : Mass can neither be created nor be destroyed. Law of Defi nite Proportions (Joseph Proust) : A given compound

Chemistry Class XI2

always contains the same elements in the same proportion by mass. Law of Multiple Proportions (John Dalton) : When two elements combine to form two or more compounds, then the different masses of one element, which combine with a fi xed mass of the other, bear a simple ratio to one another. Gay Lussac’s Law : When gases combine or are produced in a chemical reaction, they do so in a simple ratio provided all gases are in the same temperature and pressure.

e.g., 2H2 (g) + O2 (g) 2H2O (g) 2 vol 1 vol 2 Vol (at same T, P)

Atomic Mass : It is defi ned as the average relative mass of an atom of an element as compared to the mass of an atom of carbon – 12 taken as 12.

Atomic mass is represented by ‘u’ (unifi ed mass). 1u = 1.66056 × 10–24 g

Molecular mass : It is the sum of the atomic mass of the elements present in the molecule.For example : Molecular mass of CH4 = (1 × 12) + (4 × 1) = 16 u

Avogadro Number : It is the amount of atoms or molecules present in one mole of a substance.

Avogadro number (NA) = 6.022 × 1023

Molar Mass : The mass of one mole of a substance in grams is called its molar mass.

For example : Molar mass of CH4 = (1 × 12) + (4 × 1) = 16g mol–1

Mole (n) : It is amount of a substance that contains as many particles or entities as the number of atoms in exactly 12 grams of pure C-12.

1 mole of a substance = Molar mass of substance = Avogadro’s Number of chemical units = 22.4L volume at STP of gaseous substancee.g., 1 mole of CH4 = 16g of CH4 = 6.022 × 1023 molecules of CH4 = 22.4L at STP

n = = = =

Molar Volume (Vm) : It is volume occupied by one mole of any substance. Molar volume of a gas = 22.4L at STP (273 K, 1atm) or 22.7L at STP (273

3Some Basic Concepts Of Chemistry

K, 1 bar) Calculating Molar Volume: PV = nRT

V = Or

V =

Percentage Composition : Mass % of the element

=

Empirical Formula : It represents the simplest whole number ratio of various atoms present in a compound. e.g., CH is the empirical formula of benzene.

Molecular Formula : It shows the exact number of different of atoms present in a molecule of a compound. e.g., C6H6 is the molecular formula of benzene.

Relationship between empirical and molecular formulae : Molecular formula = n × Empirical formula

Where; n =

Information Conveyed by a chemical equation :

N2(g) + 3H2(g) 2NH3(g)

(i) 1 molecule of N2 + 3 molecules of H2 2 molegules of NH3

(ii) 1 mole of N2 + 3 mole of H2 2 mole of NH3

(iii) 1 × 28g of N2 + 3 × 2 g of H2 2 × 17 g of NH3

(iv) 1 × 22.4L of N2 + 3 × 22.4L of H2 2 × 22.4L of NH3

at STP at STP at STP

Limiting Reagent : It is the reactant which gets consumed fi rst or limits the amount of product formed.

Mass Percent : It is the mass of the solute in grams per 100 grams of the solution.

Chemistry Class XI4

Mass percent =

Parts per million (ppm) : It is part of solute per million part of solution by mass.

ppm =

Molarity (M) : It is number of moles of solute dissolved per litre (dm3) of the solution.

Molarity =

Molarity equation : M1V1 = M2V2

(Before dilution) (After Dilution) Malarity of a solution decreases on increasing temperature. Malarity of pure water is 55.56 mol L–1

Molality (m)—It is number of moles of solute dissolved per 1000g (1kg) of solvent.

Molality =

Molality is independent of temperature. Mole Fraction(x) is the ratio of number of moles of one component to

the total number of moles (solute and solvents) present in the solution.

x1 = and x2 =

The sum of all the mole fractions in a solution is equal to one. i.e., x1 + x2 = 1

Importance of Chemistry & Nature of Matter 1-Mark Questions

1. Name two chemical compounds used in treatment of cancer.

2. What is AZT ? Write its use.

3. Give an example each of homogeneous and heterogeneous mixture.

4. Differentiate solids, liquids & gases in terms of volume & shapes.

5. Classify following as pure substances and mixtures : Air, glucose, gold,

5Some Basic Concepts Of Chemistry

sodium and milk.

6. What is the difference between molecules and compounds? Give examples of each.

Properties of matter and their Measurement

7. What is the SI unit of density ?

8. What is the SI unit of molarity ?

9. Defi ne accuracy and precision.

10. What are the two different system of measurement ?

11. What is the difference between mass & weight ?

Uncertanity in Measurement

12. Defi ne signifi cant fi gures.

13. Defi ne accuracy and precision

14. Which measurement is more precise 4.0g or 4.00g ? [Ans. 4.00 g]

15. How many signifi cant fi gures are there in (i) 3.070 and (ii) 0.0025 ?

[Ans. (i) 4 (ii) 2]

16. Express the following in the scientifi c notation : (i) 0.0048 (ii) 234,000

Laws of Chemical Combinations & Dalton’s Atomic Theory

17. State Avogadro’s law.

18. State law of defi nite proportions.

19. State Gay Lussac’s Law of combining volumes of gases.

20. If ten volumes of dihydrogen gas react with fi ve volumes of dioxygen gas, how much volume of water vapour would be produced ? [Ans. 10 volumes]

Atomic and Molecular masses and Mole Concept

21. Defi ne unifi ed mass (u).

22. Calculate the number of atoms in 32.0 u of He. [Ans. 8]

23. Defi ne molar volume of a gas.

24. What is the volume of 17 g of NH3 gas at STP (298 K, 1 atm) ? [Ans. 22.4 L]

25. What is the value of one mole ?

Chemistry Class XI6

26. Calculate the number of molecules present in 22.0 g of CO2.

[Ans. 3.011 × 1023]

27. How many molecules of SO2 are present in 11.2 L at STP ?

[Ans. 3.011 × 1023]

28. Which has more number of atoms ? 1.0 g Na or 1.0 g Mg. [Ans. 1.0 g Na]

29. How many oxygen atoms are present in 16 g of ozone (O3) ?

[Ans. 2.007 × 1023]

30. At STP, what will be the volume of 6.022 × l023 molecules of H2 ?

[Ans. 22.4L]

31. 1L of a gas at STP weighs 1.97g. What is molecular mass ?

[Ans. 44.128 g mol–1]

Percentage Composition, Empirical and Molecular Formula

32. Write the relationship between empirical formula and molecular formula.

33. Which is more informative ? Empirical formula or Molecular formula.

34. A subtance has molecular formula C6H12O6. What is its empirical formula.

35. Empirical formula of a compound X(Molar mass = 78 mol–1) CH. Write its molecular formula.

Stochiometry and Stoichiometric Calculations

36. How are 0.5 mol Na2CO3 and 0.5 M Na2CO3 different from each other ?

37. Why molality is preferred over molarity of a solution ?

38. Defi ne molarity of a solution.

39. What is the effect of temperature on molarrity of solution ?

40. What is limiting reactant in a reaction ?

Importance of Chemistry & Nature of Matter 2 Mark Questions

1. How can we say that sugar is solid and water is liquid?

2. How is matter classifi ed at macroscopic level?

3. Classify following substances as element, compounds and mixtures : water, tea, silver, steel, carbon dioxide and platinum.

7Some Basic Concepts Of Chemistry

Properties of matter and their Measurement

4. The body temperature of a normal healthy person is 37°C. Calculate its value in°F.

5. At what temperature will both the Celsius and Fahrenheit scales read the same value?

6. Convert 5L into m3.

7. What does the following prefi xes stand for : (a) pico (b) nano (c) micro (d) deci

Uncertanity in Measurement

8. How many signifi cant fi gures are present in the answer of the following calculations :

(i) 0.0125 + 0.8250 + 0.025 (ii)

9. Convert ‘450 pm’ into SI unit and write the answer in scientifi c notation upto 2 signifi cant fi gures. [Ans. 4.5 × 10–10 m]

10. The density of vanadium is 5.96 g cm–3. Express this in SI unit.

[Ans.5960 kg m–3]

Laws of Chemical Combinations & Dalton’s Atomic Theory

11. 45.4 L of dinitrogen reacted with 22.7 L of dioxygen and 45.4 L of nitrous oxide was formed. The reaction is given below : 2 N2 (g)+ O2 (g) 2 N2O (g) Which law is being obeyed in this experiment? Write the statement of the law.

12. Write main points of Dalton’s Atomic Theory.

Atomic and Molecular masses and Mole Concept

13. Give one example each of a molecule in which empirical formula and molecular formula is (i) Same (ii) Different.

14. Calculate the number of moles in the following masses :

(i) 7.85g of Fe; (ii) 7.9mg of Ca

15. Calculate average atomic mass of chlorine using following data:

Chemistry Class XI8

Isotope % Natural abundance Molar mass35Cl 75.77 34.968937Cl 24.33 36.9659 [Ans. 35.5 u]

Percentage composition, empirical and molecular formula

16. Give one example of molecule in whch empirical formula and molecular formulaare (i) same (ii) different.

17. Calculate the present of carbon, hydrogen and oxygen in ethanol (C2H5OH)

[Ans. 52.14%, 13.13%, 34.73%]

18. How much copper can be obtained from 100 g of CuSO4 ? [Ans. 39.8g]

Stiochiometry and Stoichiometric Calculations

19. Calculate the amount of water (g) produced by the combustion of 16 g of methane. [Ans. 36g]

20. How many moles of methane are required to produce 22 g CO2 (g) after combustion? [Ans. 0.5 mol]

21. A solution is prepared by adding 2 g of a substance A to 18 g of water. Calculate the mass per cent of the solute. [Ans. 10%]

22. Calculate molarity of water if its density is 1.00 g mL-1. [Ans. 55.56 M]

23. Calculate the molarity of NaOH in the solution prepared by dissolving its 4 g in enough water to form 250 mL of the solution. [Ans. 0.4 M]

24. The density of 3 M solution of NaCl is 1.25 g mL-1. Calculate molality of the solution. [Ans. 2.8m]

25. Calculate the molarity of a solution of ethanol in water in which the mole fraction of ethanol is 0.040 (assume the density of water to be one).

[Ans. 2.31 M]

26. NH3 gas can be prepared by Haber’s process as, N2(g) + 3H2 (g) 2NH3(g). At a particular moment concentration of all the species is 2 moles; calculate the concentration of N2 and H2 taken initially.

[Ans. 3 mole, 5 moles]

27. A sample of drinking water was found to be severely contaminated with

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