Chemistry and Industry

1. Cement was invented in the year 1824 buy J.Aspin.
2. Cement is a mixture of calcium silicates and calcium aluminates.
3. Raw materials for cement are limestone and clay.
4. Cement is prepared by two methods. They are dry method and wet method.
5. At the time of preparing cement temperature maintain is 1700 - 1900°C
6. Hard greyish balls of calcium silicate and Aluminium silicates are called as ''Clinker Cement"
7. Glass is a mixture of sodium silicate, calcium silicate and sand.
8. Heating temperature of glass is 1000°C
9. Liquid glass is cooled by a special method known as annealing.
10. Ceramics are the products made from clay, felspar and sand.
11. Pottery or Terra-cotta are products of porous nature such as pots/jugs etc.
12. earthenware are glazed articles such as porcelain wall tiles crucibles and dishes.
13. Plastics are polymeric organic materials
14. Plastics are synthesized by addition or condensation polymerisation
15. Adhesives hold solid materials together by surface attachment
16. Gum-arabic, Shellac, animal glue etc are natural adhesives.
17. Urea formaldehyde, epoxy resins and silicons are some synthetic adhesives.
18. Plastics with certain properties are spun into fibres known as man-made fibres.
19. Melt spinning and dry spinning are two methods of spinning fibres from plastics.
20. Nylon, decoron, cellulose acetate and orion are some examples for man-made fibres.
21. Cosmetics are meant for improving attraction of a person.
22. Cold creams mainly contain almond-oil, bees-wax and rose water.
23. Dyes are useful for colouring textiles.
24. Chromophore and auxochrome are two groups that must present in a dye.
25. Dyes are classified into acidic, basic, direct, mordant, vat etc depending upon the method of application of dye to the fibre.
26. Depending upon chromophores in dyes they are classified as nitro dyes, nitrosodyes, Azo dyes and Quinon dyes etc.
27. Drugs are substances used in prevention, diagnosis treatment or cure of a disease.
28. Drugs are classified into six types based on therapeutic action. They are
      i) Drugs working on central nervous system
      ii) Drugs working on cardiovascualr system
      iii) Drugs working on foreign organisms
      iv) Chemo therapy drugs,
      v) Vitamins,
      vi) Hormones (Insulin, cartisone)
29. Pharmaceuticals are modified from of drugs such as tablets, capsules, lotions, syrups, tonics and injections.
30. Aspirin and paracetamol are the common drugs used in our daily life.
31. Petroleum products are converted into organic chemicals called pertochemicals, petrochemical in turn are converted into many useful products.
32. Petrol is a mixture of hydrocarbons with 5 to 9 carbon atoms.
33. Cooking gas in a mixture of butane, propane and propene compressed to liquid form called ''liquefied petroleum gas LPG".
33. Plants derive nutrients from soil, water, air
34. Required nutrients must be added to the soil in the form of manure or fertilizer.
35. Fertilizer provide N, P and K to the plants.
36. Urea, Ammonium Phosphates, Super Phosphates are common examples of fertilizers.
37. Plants require certain nutrients in micro quantities which are known as micronutrients.
Example: B, Cr, Mo, Mn, CO, Zn, Fe

Note: Maximum items are covered from this chapter in every public exam. One must be through with the chapter.

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Oils and Fats

1. Oils and Fats are tri esters of glycerol and fatty acids.
2. Lauric and stearic acids are the saturated fatty acids. Lauric acid → C11H23COOH.

     Stearic acid → C17H35COOH.

     Sources of Lauric acid → coconut oil and Butter
     Sources of stearic acid → Animals fat and Butter.
3. Oleic acid (C17H35COOH) and Lionoleic acid (C17H29COOH) are the unsaturated fatty acids.
              Sources of oleic acid → Butter, cotton seed oil, soya beans
              Sources of Linoleic acid → Lin seed.
4. Fats are prepared by the hydrogenation of oils in the presence of 'nickel' catalyst
5. Soaps are salts of fatty acids.
6. Chemically detergents are sodium salts of alkyl benzene sulphonates or fatty alcohol sulphates.

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Carbohydrates and Proteins

1. Polyhydroxy aldehydes or ketones called as carbohydrates (or) Compound of carbon and water are called carbohydrates.
2. General formula of carbohydrates Cx (H2O)y
3. Carbohydrates are classified based on taste as sugars and non-sugars.
4. Carbohydrates containing aldehyde functional groups are called Aldoses, if carbohydrates contain ketone functional group are called ketoses.
5. Depending on number of carbon atoms in carbohydrates they are classified as Trioses (Three Carbon atoms), tetroses (4 carbon atom) and Pentoses with 5 carbon atoms.
6. Based on the character of hydrolysis of carbohydrates they are classified as monosaccharirdes, oligosaccharrides and polysaccharrides.
              Example for
               Monosaccharirdes: a) Glucose, Fructose and Mannose
               Oligosaccharrides: (They form 2 to 9 monosacharrides on hydrolysis).Ex: Sucrose, Maltose.
                Polysaccharrides: (They give more monosaccharirdes on Hydrolysis. General formula is (C6H10O5)n ) Ex: Starch, Cellulose.
7. Carbohydrates (glucose) in prepared by plants by a process known as photosynthesis.
8. Glucose given energy to the living cells.
9. Carbohydrates provide food, clothes and shelter for us.
10.Glucose is tested by Tollen's and Benedict's reagent.
11.Sugar is manufactured from sugar-cane which involves broadly four steps.
          a. Extraction of juice from sugar-cane
          b. Purification of sugar-cane juice
          c. Concentration of juice and crystallisation.
          d. Separation of Crystals and drying.
12.The byproducts of sugar industry are bagasse, wax, press mud and molasses. Alcohol is prepared from molasses by fermenting the same with yeast cells. Yeast cells contains two types of enzymes. CO2 is a byproduct of alcohol industry.
13. Amino acids contain amino (NH2) and Carboxylic Acid (CaOH) groups is a molecule.
14. Amino acids have salt like or zwitter ionic structure.
15. Amino acids are the building blocks of proteins.
16. Human body cannot prepare nine amino acids which are called as essential amino acids.
17. The – CO – NH – bond is called peptide bond.
18. The result product obtained by the joint of a large number of amino acids, is called polypeptide. Hemoglobin which carries oxygen is a protein consists 574 amino acids. By changing one amino acid in Hemoglobin this will be converted to sickle cell hemoglobin.
19. Each protein has its own number and sequence of amino acids and performs a specific function.

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Chemistry of Carbon Compounds

1. Valency of carbon is four. Carbon exhibits allotropy. Allotropic forms of carbon are Diamond, graphic and C60.
2. In diamond carbon atoms are arranged in tetrahedral structure.
3. In graphite carbon atoms are arranged in hexagonal structure.
4. C60 having a football like structure contains 12 pentagons and 20 hexagon rings.

5. Solid CO2 is known as dry ice.
6. Carbon exhibits catenation and Isomerism.
7. Dry distillation of coal given many important carbon compounds.
8. The compounds of carbon and hydrogen are called as Hydrocarbons.
          a) Alkanes are saturated hydrocarbons. They under go substitution reactions.
          b) Alkenes, Alkynes and Benzenes are unsaturated hydrocarbons.
          c) Some unsaturated hydrocarbons undergo polymerisation.
9.  A group of atoms in carbon compounds showing characteristic properties is called functional group.
10. Alcohols (– OH), aldehydes (– CHO), Ketones acids (– COOH), esters (– COOR) and ...... (– NH2) are some important esters (– COOR) and amines (– NH2) are some important functional groups.
11. Aldehydes are detected with Tollen's test.
12. A piece of sodium metal reacts with alcohol immediately and liberates Hydrogen. This is
due to C – OH group.

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Acids,Bases And Salts

1. Acids turn blue litmus to red and bases turn red litmus to blue.
2. Acids are formed when metallic oxides are dissolved in water.
3. Bases are formed when non-metallic oxides are dissolved in water.
4. Low volatile acids are prepared from high volatile acids.
5. Hydrogen is released When acids react with metals.
6. Water and salt are formed when acids react with bases.
7. According to Arrhenius theory acids produce H+ ions, and bases produce OH– ions in aqueous media.
8. Combination of H+ and OH– ions is called neutralisation.
9. All acids and bases are not falling under Arrhenius theory of Acids and bases.
10. [H+] × [OH–] is called ionic product of water.
11. pH scale was introduced by Sorenson. pH = –log[H+]
12. pH value of acids is between '0' and '7' and that of bases is between 7 and 14.
13. Neutral solution has a pH of 7.
14. For a reaction between strong acid and strong base the heat neutralisation is 13.7 K.Cal mol–1.
15. For a reaction involving weak acid and weak base, the heat of neutralization is less than 13.7 K.Cal.

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Soltions

Homogeneous mixture: Substances which dissolve with each other thoroughly to form a uniform mixture is called homogeneous mixture.

Eg: Water + Salt.

Solutions: A homogeneous mixture formed with two or more substances is called as solution.

Components: Substances present in solution are called as components. Solvent: A component which is larger in quantity in a solution is called solvent.

Eg: Water in salt water.

Solute: A component which is less in quantity in a solution is called solute.

Eg: Salt in salt water

Aqueous Solution: If the solvent in a solution is water, the solution is called as aqueous solution.

Types of Solution: Solution are divided into unsaturated, saturated and super saturated solutions based on the solubility of the solute.

Note: Super saturated solution is more unstable.

Solubility: The maximum amount of solute by weight in grams in 100 grams of solvent at constant temperature is called the solubility of that substance.

Eg: Solubility of certain compounds at 30°C.

Factors affecting solubility of substances.
(a) Nature of solute & solvent.
(b) Temperature.
             Generally polar substances dissolve in polar solvents, but not in a non-polar solvents. In the same manner non-polar substances dissolve in non-polar solvents but not in polar solvents.
Eg (1): Copper Sulphate is a polar solute. Hence it dissolves in water which is also a polar solvent. But the same copper sulphate doesn't dissolve in kerosene as kerosene is a non-polar solvent.

Eg (2): Naphthalene is a non-polar solute. Hence it dissolves in kerosene which is also a non-polar solvent. But the same Naphthalene doesn't dissolve in water which is a polar solvent.

            Solubility of KOH, NaNO3, KNO3, NH4Cl increase with the increase in temperature.
            Solubility of Ce2(SO4)3 and all gases decrease as the temperature increases.
            Effect of temperature is negligible with respect to NaCl solubility.

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Alkaline Earth Metals

♦ Beryllium (Be), Magnesium (Mg), Calcium (Ca), Strontium (Sr), Barium (Ba) and Radium (Ra) are the elements belong to IIA group of the periodic table.
♦ Their general electronic configuration is ns2. (where n = 2, 3...)
♦ Radium is a Radioactive element.
♦ Atomic size of elements from Beryllium to Radium increaser.
♦ Ionization energies and electronegativity of IIA elements decrease turn Be to Ra.
♦ All elements react with water to form their respective hydroxides and liberate Hydrogen.
♦ All elements react with oxygen to form their respective oxides. But Barium forms even peroxides besides oxides.
♦ All elements except Beryllium form Hydrides when they react with hydrogen.
♦ All elements react with chlorine and form their respective metallic chlorides.

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Periodic Classification of Elements

♦ The first classification of elements is due to dobereiner in 1817.

♦ Dobereiner Triad Theory: "The atomic weight of the middle element is the arithmetic mean of first and third elements (or) the atomic weights of all the three elements are approximately the same"
            Eg: i) Li, Na, K ii) Fe, Co, Ni iii) Cl, Br, I.
♦ Newland's Concept of Octaves: Every eighth element is having the similar properties to that of the first element like in octave of music.
♦ Mendeleef, Lother Meyer used atomic property. Atomic weight for classification
elements.
♦ Mendeleef's Periodic Law: The properties of the elements are the periodic functions of their atomic weights".
♦ Modern Periodic Law: The properties of the elements are the periodic functions of their atomic number (or) electronic configuration.
♦ Modern Periodic Table is called long form of periodic table.
♦ Long form of periodic table is divided into seven horizontal rows called "Periods" and eighteen vertical columns called "Groups"
♦ The first period has only two elements.
♦ The second and third period have 8 elements each.
♦ Fourth, Fifth and sixth periods have 18 elements each.
♦ Lanthanides and actinides are placed at the bottom of the periodic table.
♦ Seventh period is incomplete.
♦ Based on electronic configuration, the elements are classified into 4 types. They are inert gases, representative elements, transition elements representative elements and inner transition elements.
♦ Atomic radius: It is the distance between the centre of the nucleus and outermost orbital.
♦ It is expressed in units of Angstroms.
♦ Ionization energy: It is the energy required to remove an electron from the outermost orbital in the gaseous state.
♦ IE is expressed in ev or KCal.mol–1 or K.J.mol–1.
♦ Electronegativity: The ability of the bonded atom to attract the electron density of the shared electrons.
♦ Electronegativity is expressed ... pauling scale.
♦ In a period, Atomic radius decreases and ionisation energy and electronegativity increases.
♦ In a group, atomic radius increases and IE and EN value decreases.

Questions 

Section-I:

Short Answer Questions (2 Marks each)

1. How does the atomic radius (size) vary in a period and a group?
2. Why does atomic size decreases from left to right?
3. How does the ionsation energy vary in a period and a group?
4. Distinguish between oxidation and Reduction?
5. Which group elements can be used as oxidising and reducing agents?
6. What is Newland's concept of Octaves?

Section-II

Very Short Answer Questions (1 Mark Each)

1. Which group elements can be used as strong reducing agents?
2. On which atomic property is the Mendeleef's periodic table based?
3. What is oxidation?
4. Explain the relationship between ionisation energy and size of the molecule?
5. What are inner transition elements?
6. Write the general electronic configuration of inert gases?
7. Define atomic radius?
8. Which group of elements have the highest electropositive character?
9. State Mendeleef's periodic law?
10. Define electropositive character?

Section III

Long Answer Questions (4 Marks each)

1. How does atomic size and ionisation energy vary in a period and in a group?
2. How does the following properties vary in a period and in a group?
                  a) Atomic Size (or) Atomic radius
                  b) Oxidising Property
                  c) Electronegativity
                  d) Electropositivity
                  e) Ionisation energy
3. Explain the main features of long form of the periodic table?
4. Explain the electronic configuration of inert gases?
5. Answer the following
                 a) Define Ionisation Energy
                 b) Newland's Concept of Octaves
                 c) Inner Transition Elements
                 d) Transition Elements

Section IV

- Diagrams

Answers 

Part - A, Section - I

Short Answers

1A. In a period: The atomic radius decreases from left to right. This is because, as the atomic number increases in a period, the nuclear attraction over the valence electrons increases and the atomic radius decreases.
In a group: The atomic radius increases from to to bottom in a group. This is due to the addition of one extra shell from one element to another element.

2A. 1. Atomic size is expressed in °A units.
       2. Atomic size decreases from left to right in the periodic table.
       3. This is because, as the atomic number increases in a period, the nuclear attraction over the electron charge cloud increases and as a result the radius decreases.

3A. Period:
1. In a period, from left to right, the ionisation energy do not follow any regular trend. In the first period it increases form hydrogen to helium.
2. But, in second period, it increases from Li to Be decreases at Boron.
3. From boron to nitrogen it again increases and then decreases at oxygen.
4. This is because nitrogen has half filled electron configuration (2p3) which is stable.
5. Therefore, Ionisation energy do not follow any regular trend in a period.
Groups:
1. In a group, Ionisation energy decreases from top to bottom.
2. As we go from top to bottom, the size of the atom increases and attractive power of the nucleus on the outermost electron decreases.
3. Hence, ionisation energy decreases in a group.

4A. OXIDATION                                                                              REDUCTION
1.Addition of oxygen to a given Reduction.                     1. Addition to hydrogen to a compound is                                                                                                         compound is called oxidation
2.Removal of hydrogen from a                                         2. Removal of oxygen atom from a called
  compound is called oxidation                                             compound is called Reduction

5A. 1. Group - IA and IIA elements can be used as reducing agents.
       2. Group - VII A elements can be used as oxidising agents.

6A. Newland's concept of octaves:
If the elements are arranged sequentially in the increasing order of their atomic weights, every eighth element is having the similar properties to that of the first element like in the octave of music. This is called "Newland's Concept of Octaves".
Eg:   1   2   3   4   5   6   7   8
         Li  Be B  C  N  O  F  Na
Here, properties of Lithium and sodium are the same.

Very Short Answers:

1.A Group - IA and Group - IIA elements can be used as strong reducing agents.
2A. Mendeleef's periodic table is based on atomic weights of the elements.
3A. Oxidation: Addition of oxygen atom to a compound (or) Removal of hydrogen from a compound is called oxidation.
4A. In a period, as the atomic size decreases, hence he ionisation energy increases. In a group, as the atomic size increases, hence the ionisation energy decreases.
5A. The elements belonging to f-block are called "inner-transition elements"
6A. The general electronic configuration of inert gases is "ns2 np6"
7A. Atomic radius: "The distance between the centre of the nucleus and the outer most orbital of an atom" is defined as atomic radius.
8A. Group - IA and Group-IIA elements are having the highest electropositive character.
9A. Mendeleef's Periodic law: "The physical and chemical properties of the elements are the periodic functions of their atomic weights"
10A. Electropositive Character:
- The ability of a bonded atom in a molecule to attract the bonded pair of electrons
towards itself is called ''electropositive character".

Section III

Long Answer Questions (4 Marks Each)

1A. Atomic Radius (or) Atomic Size:
1. "The distance between the nucleus and outermost orbital of an atom" is called "Atomic radius.
2. Atomic radius (or) Atomic size is expressed in Angstrom units (1 °A = 10–8 cm)
3. In a period, the atomic radius decreases from left to right of the periodic table.
4. This is because, as the atomic number increases in a period, the nuclear attraction over the electron charge cloud increases and as a result the radius decreases.
5. In a group, the atomic radius increases front top to bottom.
6. This is due to the addition of one extra shell from one element to another. Ionization Energy:
1. The minimum energy required to remove an electron from the outermost orbital of an atom in the gaseous state" is called ionisation energy.
2. Its units are electron volt (or) kilo joules/mole.
3. In a period from left to right, ionisation energy do not follow any regular trend. In the first period ionisation energy increases from hydrogen to helium, as it has stable
configuration 1s2
4. In the second period. I.E. increases from Li to Be and decreases at Boron.
5. From Boron to nitrogen it again increases and then decreases at oxygen. This is because nitrogen has half - filled stable electronic configuration (2p3)
6. In a group the I.E. decreases from top to bottom. This is because as we go from top to bottom the size increases, so it is easy to remove the electron.

2A. a) Atomic Radius:
1. In a period, the atomic radius decreases from left to right.
2. In a group, the atomic radius increases from top to bottom.
b) Oxidising Property:
1. In a period, oxidising property increases from left to right.
2. In a group, the oxidising property decreases from top to bottom.
c) Electronegativity:
1. In a period, the electroneagativity increases due to the decrease in the size of the atom.
2. In a group, the electronegativity decreases due to increase in the size of the atom.
d) Electropositivity Character:
1. In a period, the electropositive character decreases.
2. In a group, the electropositive character increases.

3A. Main Features of Long form of the Periodic Table:
The long form of the periodic table relates the properties of elements to their electronic configuration. This is otherwise called "Modern Periodic Table".
1. It consist of 7 periods and 18 groups.
2. Every period starts with alkali metal and ends with inert gas.
3. First period has 2 elements.
4. 2nd and 3rd periods have 8 elements.
5. 4th and 5th periods have 18 elements each.
6. Sixth period consists of 32 elements.
7. Seventh period is incomplete period.
8. All the elements are classified into four blocks like s,p,d and f.
9. IA and IIA group elements are called s-block elements.
10. IIIA and IIIA group elements are called p-block elements.
11. The elements between s and p - blocks are known as d-block elements.
12. The elements lie at the bottom are called f-block elements.
13. Based on electronic configuration, elements are classified into four types like representative, Transition, inert gases and inner transition elements.
14. s and p block elements together known as representative elements.
15. d - block elements are called transition elements.
16. Zero group elements are known as inert gases.
17. f - block elements are called inner transition elements.
18. f - block elements with atomic numbers 58 to 71 are called lanthanides and 90 to 103 called actinides.

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Chemical Bond

♦ Atoms combine to form molecules.
♦ Formation of chemical bond involves redistribution of electrons.

♦ Ionic Bond: "The chemical bond formed by the complete transfer of one or more electrons from one atom to other".
                Eg: NaCl, MgO
♦ Covalent Bond: "The bond formed by the sharing of electrons between two atoms".
                Eg: H2, Cl2, F2, H2O
♦ Octet configuration: "The filling up of valence orbital with 8 electrons".
♦ As the atoms approach each other, the attractive forces between electrons of one atom and the nucleus of other atom increases, results in the decrease of combined potential energy of combined atoms.
♦ At the critical internuclear distance, the energy of the molecule (or) combined atoms is less than the sum of the energies of individual atoms and thus a chemical bond is formed.

Potential- Energy Level Diagram:

♦ Zero group elements have stable electronic configuration having 2 (or) 8 electrons.

♦ A covalent bond is formed by the overlap of two pure atomic orbitals (or) two hybridised orbitals (or) one hybridized and one atomic orbital.

♦ Strong bond is formed by the maximum overlap of orbitals.

♦ Sigma-bond: End-on-end overlap of orbitals lead to the formation of sigma bond. It is formed by the overlap of s-s orbitals, p-p orbitals and s-p orbital overlap.

♦ Pi-bond: Partial overlap of orbitals lead to the formation of pi (π) bond.

♦ Sigma bond is stronger than pi-bond, sigma bond exists independently.

       Eg: H2, HCl, Cl2, CH4

♦ π-bond has no independent existence. It exists only after the formation of σ bond.

♦ In a double bond, one sigma and one π bond are present.

        Eg: C2H4, O2, CO2, etc.,

♦ In a triple bond, one sigma and two pi-bonds are present.

        Eg: N2, C2H2, HCN etc.,

♦ In a co-ordinate covalent bond, both the electrons are supplied by one atom and shared between two atoms.

♦ The primary requirement for the formation of a covalent bond is that one atom should have a lone pair of electrons (with octet configuration) to donate and another atom should have an empty orbital to receive it.

♦ Molecules having linear structure are

      Eg: CO2, BeCl2, C2H2, HCN etc.,

♦ Molecules having pyramidal structure are

      Eg: NH3, PCl3, PH3, ....

♦ Molecules having double bond are

      Eg: O2, CO2, C2H4, ....

♦ Molecules having triple bond are

      Eg: N2, HCN, C2H2...

♦ Water molecule have "V" - shape.

♦ PCl5 has trigonal bi-pyramidal shape.

Questions

SECTION - II

Very Short Answer Questions (1 Mark Each)

1. Draw the shape of water molecule.

2. Name the two molecules having pyramidal shape?

3. What is the shape of ammonia molecule? Draw it and explain.

4. Give three examples of molecules having a double bond?

5. Give two examples of molecules having a triple bond?

6. Name two molecules having p-p overlapping?

7. Which orbitals can form pi (π) bond?

8. What is the shape of carbon dioxide molecule?

9. Show the electron dot structure of H2 molecule?

10. What are multiple covalent bonds? Give examples?

SECTION - III

Long Answer Questions (4 Marks Each)

1. Explain the formation of a double-bond with example?

2. Explain the formation of co-ordinate covalent bond?

3. Explain s–p overlap with examples?

4. Explain the formation of triple bond in Nitrogen molecule?

(or)

Explain the formation of multiple bond with an example?

5. Discuss the type of overlaps that are possible with s and p orbitals.

SECTION - IV

DIAGRAMS (5 Marks Each)

1. Draw the bond formation of Hydrogen Chloride and Nitrogen molecules in a diagram.

2. Draw the diagram showing the formation of a double bond (oxygen molecule)?

3. Draw the diagram showing the formation of p–p overlap in fluorine molecule.

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Atomic Structure

♦ Rutherford's model of atom and its draw backs

♦ Bohr's model of atom and its draw backs

♦ Sommerfelds elliptical model

♦ Sub Energy levels - Quantum Numbers

♦ Atomic orbitals

♦ Relative energies of the atomic orbitals

♦ Electronic configuration of elements

♦ Some Physical Quantities of Atoms

♦ 1) Ionization potential

   2) Atomic radius

   3) Electron affinity

♦ Some principles to write electronic configuration

         Aufbau principle

         Hund's rule of

         Pauli's exclusion principle

         Maximum multiplicity

Unit-1

Important Points:

1. Electrons, protons and neutrons are called sub-atomic particles.

2. J.J. Thomson's atomic model could not explain the atomic spectra.

3. Rutherford proposed "Planetary atomic model" based on α-ray scattering 

experiments.

4. According to this model, electrons are subjected to two types of forces that is

       1) The force of attraction between electrons and the nucleus

       2) The centrifugal force

           These two forces are equal and opposite.

5. Max Planck proposed the Quantum theory of radiation.

6. E = hν Planck's constant h = 6.625 × 10–27 erg.sec. (or) 6.625 × 10–34 J.sec.

7. Bohr proposed his atomic model based on Planck's Quantum theory.

8. Bohr model explains stationary orbits, angular momentum and radiation of energy.

9. Bohr model could not explain the atomic spectra of higher elements, Zeeman effect, 

quantization of angular momentum.

10. Sommerfeld proposed the elliptical orbits and sub-stationary states.

11. Orbits are indicated by "Principal Quantum number" ie 'n', sub-stationary states are 

indicated by "Azimuthal Quantum number i.e 'l', orientation of orbitals is indicated by 

"magnetic Quantum number" i.e, 'm' and spinning of electrons of "Spin Quantum 

number" 's'.

12. l value = (n – 1), the number of 'm' values is equal to (2l + 1) and ranges from

                   –l....0....+l. s is +1/2 or –1/2.

13. "The region in space where there is finite probability of finding electron" is called "atomic orbital".

14. s, p, d and f orbitals are present in the orbits. 1, 2, 3 and 4 onwards (or) K, L, M and N onwards successively.

15. Orbitals having same energy are called degenerate orbitals.

16. "Electron occupies the orbital whose (n + l) value is minimum. If (n + l) value is same, it occupies the orbital having less 'n' value" ____ Aufbau Principle.

17. "No two electrons will have all the four Quantum numbers same" ___ Pauli exclusion principle.

18. "Pairing of electrons takes place when all the degenerate orbitals are occupied by one electron each" ____ Hund's rule.

19. Atomic radius is the distance between the nucleus and outermost orbital.

                 Units: A° ; 1A° = 10–8 cm or 10–10m.

20. Ionization potential is "the minimum energy required to remove an electron from

outermost orbital in the gaseous state".

                 Units: ev (or) K.J. mol–1 (or) K.Cal.mol–1.

21. Electron affinity is "the energy released when an electron is added to a neutral gaseous atom in the ground state".

                 Units: ev (or) K.J. mol–1 (or) K.Cal.mol–1.

Questions

SECTION - I

Short Answer Questions (2 Marks Each)

1. Explain why electrons enter into '4s' orbital but not '3d' after filling '3p' orbital?

2. Write the electronic configuration of (a) Nitrogen (b) Chromium (c) Copper?

3. Distinguish between orbit and orbital?

4. Write a short notes on Quantum theory of radiation?

5. Distinguish between Principal Quantum number and sub-level Quantum number.

SECTION - II

Very Short Answer Questions (1 Mark Each)

1. What is Planck's Quantum equation? and what is the value of Planck's constant?

2. What is a stationary orbit?

3. What are degenerate orbitals?

4. Who proposed the principal Quantum number?

5. What are the limits of azimuthal Quantum number for a given 'n'?

6. What are upper and lower limits of m for l = 4?

7. How many 'm' values are possible for l = 3

8. What is the shape of '1s' orbital?

9. What is nodal plane?

10. Write the designations of any two d-orbitals?

11. How are the three p-orbitals designated?

12. What is the unit of electron affinity?

SECTION - III

Long Answer Questions (4 Marks Each)

1. State The Postulates of Bohr's model. What are the defects of Bohr's model?

2. Explain the important features of Rutherford's model of atom. Discuss its draw-backs?

3. State and explain with one example The Hund's rule of maximum multiplicity?

4. Define the ionization energy and mention the factors that influence it.

5. Explain the four Quantum numbers briefly.

6. State and explain Pauli's exclusion principle with example.

7. Discuss the features of modern atomic structure.

SECTION - IV

DIAGRAMS (5 Marks Each)

1. Draw the shapes of s and p orbitals?

2. Draw the shapes of five 'd' orbitals?

3. Draw the diagram showing the sequence of filling of various atomic orbitals?

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