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11th CHEMISTRY Important Questions 2025
All Exercise MCQs
Chapter 1
SHORT QUESTIONS
1. Why isotopes have different physical but same
2. chemical properties?
3. Atomic masses may be in fractions. Justify.
4. No individual neon atom in the sample of the element has a mass of 20.18 amu.
5. Mg atom is twice heavier than that of carbon atom.
6. Calculate moles of O atoms in 9.00 g of Mg(NO3)2.
7. Calculate mass in grams of 2.74 moles of KMnO4.
8. Define mole and Avogadro number.
9. 180 g of glucose and 342 g of sucrose have the same number of molecules but different number of atoms present in them.
10. N2 and CO have the same number of electrons, protons and neutrons.
11. 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 these gases are very
12. different from each other.
13. What is stoichiometry? Write down assumptions of
14. stoichiometry.
15. Define limiting reactant with example. How does limiting
16. reactant control the amount of product formed?
17. Concept of limiting reactant is not applicable to the reversible reactions. Explain it.
18. Many chemical reactions taking place in our surrounding involve the limiting reactant.
19. Why actual yield is less than theoretical yield?
20. Define yield. How efficiency of reaction is measured?
21. Define and explain the molecular ion.
22. Differentiate between empirical and molecular formula.
23. The molecular formula is multiple of the empirical formula. Give an example.
24. What is molar volume?
25. Law of conservation of mass has to be obeyed during stoichiometric calculations.
26. One mole of H2SO4 should completely react with two moles of NaOH. How does Avogadro’s number help to explain it?
Long Questions
1. Define the following terms with examples:​​​
(i) Relative atomic mass​(ii) Molecular ion (iii) Isotope​(iv) Molar volume.
2. Define types of yields. How do we calculate the percentage yield of a chemical reaction?
3. What is limiting reactant? Give two examples. How limiting reactant can be identified?
4. Explain combustion analysis with diagram and write formulas for percentage of carbon, hydrogen and oxygen.
5. Explain evidence of atoms with the help of diagram.
6. Define stoichiometry. Give its assumption and mention laws obeyed during stoichiometric calculation.
7. Write the steps involved for the determination of empirical formula.
8. Explain Isotopes. Also describe relative abundance of isotopes.
Chapter 2
SHORT QUESTIONS
1. Define sublimation with two examples.
2. Explain the process of sublimation.
3. Define solvent extraction. State distribution law or partition law.
4. Differentiate​between adsorption and partition chromatography.
5. Define chromatography with its principle. Describe importance of chromatography.
6. What is Rf value? Why it has no units?
7. What are two phases in chromatography?
8. Define quantitative and qualitative analysis.
9. Describe the major steps involved in quantitative analysis.
10. Describe Gooch crucible and sintered glass crucible?
11. How will you decolorize the crude product?
12. How is a vacuum desiccator used for drying crystals?
13. Write four properties of a good solvent.
14. Why sintered glass crucible is preferred over Gooch crucible?
Chapter 3
SHORT QUESTIONS
1. What are isotherms? What happens to the positions of isotherms when they are plotted at high temperature for a particular gas?
2. What is absolute zero? Justify that volume of one mole of gas becomes zero theoretically at –273.16°C.
3. Write down faulty assumptions of KMT? (Causes of deviation from ideal behaviour).
4. Convert ⁰C to ⁰F and vice versa.
5. List postulates of kinetic molecular theory.
6. What is plasma? Where is plasma found?
7. What are characteristics of plasma? What do you mean by natural plasma and artificial plasma?
8. Write down four applications of plasma.
9. State Boyle’s law. Give its two mathematical forms
10. Define pressure. Give its SI units.
11. Calculate the value of gas constant.
12. What is Avogadro’s law of gases?
13. How Dalton’s law applies to the respiration process?
14. Define partial pressure. State Dalton’s law of partial pressure. Give an example.
15. Differentiate between diffusion and effusion.
16. Lighter gases diffuse more rapidly than heavier gases.
17. What is the Joule-Thomson effect?
18. Can we determine the molecular mass of an unknown gas if we know the pressure, temperature and volume along with the mass of the gas?
19. Define critical temperature and critical pressure with example.
20. Water vapours don’t behave ideally at 273K. Justify.
21. H2 and He are ideal at room temperature, but SO2 and Cl2 are non-ideal. Justify
22. Why SO2 is comparatively non-ideal at 273K but behaves ideally at 373K?
23. Gases show non-ideal behavior at low temperature and high pressure.
24. Write units of a and b (Van der Waals constants).
Long Questions
1. A sample of nitrogen gas is enclosed in a vessel of volume 380 cm3 at 120°C and pressure of 101325 Nm-2. This gas is transferred to a 10 dm3 flask and cooled to 27°C. Calculate the pressure in Nm-2 exerted by the gas at 27°C.
2. Calculate the mass of 1 dm3 of NH3 gas at 30°C and 1000 mmHg pressure considering that NH3 is behaving ideally?
3. What pressure is exerted by a mixture of 2.00 g of H2 and 8.00 g of N2 at 273K in a 10 dm3 vessel?
4. 250cm3 of the sample of hydrogen effuses four times as rapidly as 250cm3 of an unknown gas. Calculate the molar mass of unknown gas.
5. 250 cm3 of hydrogen is cooled from 127°C to -27°C by maintaining the pressure constant. Calculate the new volume of the gas at Low temperature.
6. Calculate the number of atoms in 20cm3 of CH4 at 0°C and pressure of 700 mm of Hg.
7. 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?
8. There is a mixture of hydrogen, helium and methane occupying a vessel of volume 13dm3 at 37°C and pressure of 1 atmosphere. The masses of hydrogen and helium are 0.8g and 0.12g respectively. Calculate the partial pressure in torr of each gas in mixture.
9. A sample of nitrogen gas is enclosed in a vessel of volume 380 cm3 at 120°C and pressure of 101325 Nm-2. This gas is transferred to a 10 dm3 flask and cooled to 27°C. calculate the pressure in Nm-2 entered by the gas at 27°C.
Chapter 4
SHORT QUESTIONS
1. What are dipole-dipole and Debye forces (dipole- induced dipole forces)?
2. Why boiling points of halogens and noble gases increases down the group?
3. Why HF is weaker acid than HI?
4. Define hydrogen bonding. Why H2S is gas while H2O is liquid?
5. Why ice floats on water? (Why ice has less density than water?)
6. Define liquid crystal with example. Write applications of liquid crystals.
7. Describe isomorphism, polymorphism, and symmetry with two examples.
8. What is anisotropy and allotropy? Give one example.
9. What is difference between crystalline and amorphous solids?
10. Cleavage of the crystals is itself an isotropic behaviour.
11. Describe transition temperature with examples.
12. Define unit cell. What are crystallographic elements?
13. Define the boiling point. How it depends upon external pressure?
14. Define molar heat of v***rization with one example.
15. Define lattice energy. Give an example.
16. Graphite is a conductor but a diamond is an insulator. Why?
17. What are the advantages of vacuum distillation?
18. Ev***ration causes cooling.
19. Ev***ration takes place at all temperatures.
20. Earthenware vessels keep water cool.
21. One feels a sense of cooling under the fan after the bath.
22. The boiling point of water is different at Murree Hills and Mount Everest.
23. The heat of sublimation of iodine is very high.
24. The vapour pressures of the solids are far less than those of liquids.
25. Cleavage of the crystals is itself anisotropic behaviour.
26. The electrical conductivity of the metals decreases by increasing temperature.
27. Ionic crystals don’t conduct electricity in the solid-state.
28. Ionic crystals are highly brittle.
Long Questions
1. Describe four properties of the crystalline solids.​
2. What are metallic solids? Describe their properties.
3. What are covalent solids? Explain their properties.
4. What are London forces? Explain factors affecting London forces.
5. Describe the measurement of vapour pressure by manometric method with diagram.
6. Explain molecular solids in detail.
7. Define boiling point. What is the effect of external pressure on boiling point? Give two examples.
8. Explain following types of inter Molecular forces at least with one example each i. Dipole-Dipole forces​ii. Dipole-Induced Dipole forces
9. Define ionic solids. Discuss properties of ionic solids in detail.
10. Define Hydrogen Bonding and explain its any three applications.
Chapter 5
SHORT QUESTIONS
1. Why is it necessary to decrease the pressure in the discharge tube to get the cathode rays?
2. Whichever gas is used in the discharge tube the cathode rays remains the same. Why?
3. Justify that cathode rays are material particles with a negative charge.
4. What are X-rays? State Moseley’s law.
5. Why e/m value of the cathode rays is just equal to that of electron?
6. How the bending of the cathode rays in the electric and magnetic fields shows that they are negatively charged?
7. What is the reason of production of positive rays?
8. (How positive rays are produced?)
9. Why the positive rays are also called canal rays? Give its reason.
10. Write balanced equations for two nuclear reactions.
11. Give two postulates of Planck’s quantum theory.
12. What are products of decay of neutron?
13. Calculate mass of an electron from e/m value.
14. What are drawbacks of Rutherford’s atomic model?
15. Give two postulates of Bohr’s atomic model.
16. Differentiate between continuous and line spectrum.
17. Differentiate between atomic emission and atomic absorption spectrum.
18. What are defects of Bohr’s atomic model?
19. What is Zeeman Effect and Stark effect?
20. Give de-Broglie equation and Describe Heisenberg uncertainty principle.
21. State Pauli-Exclusion principle, Hund’s rule and Auf bau principle.
22. Describe Heisenberg uncertainty principle.
23. Give electronic configuration of elements with atomic numbers 7, 11, 15, 17, 19, 21, 24, 26, 29, 35, 37, 55, 57.
24. What is the (n+l) rule?
25. How the slow neutrons prove to be more effective than the fast neutrons? (Write down reactions for the conversion of Cu into Zn.)
Long Questions
1. Derive the equation to calculate radius of electron in nth orbit hydrogen atom by using Bohr’s model.​
2. Write down measurement of charge on electron by the Millikan’s Oil Drop Method.
3. Define and explain:​(i) Atomic emission spectrum​(ii) Atomic absorption spectrum
4. Write down any four properties of cathode rays.
5. How neutron was discovered? Explain with the help of an experiment also write four properties of neutron.
6. Give four defects of Bohr’s atomic model
7. What is the Importance of principle and magnetic quantum number?
8. How nucleus was discovered? Explain with the help of an experiment also write four properties of neutron.
Chapter 6
SHORT QUESTIONS
1. The radius of an atom cannot be determined precisely. Write two reasons.
2. What is octet rule? Give two examples of compounds
3. which deviate from it.
4. Bond distance is the compromise distance
5. between two atoms.
6. Why size of cation is smaller and Size of anion is greater
7. than parent atom.
8. How ionization energy, electron affinity and electronegativity vary across periods and groups?
9. 2nd ionization energy is greater than 1st.Why?
10. Define ionization energy and electron affinity with example.
11. Why 2nd electron affinity is shown with positive sign?
12. Define atomic and covalent radii with examples
13. Define polar and non-polar covalent bond.
14. Define coordinate covalent bond with example. (How coordinate covalent bond is formed between
15. NH3 and BF3?).
16. Write down two postulates of VSEPR theory?
17. Why lone pair occupies more space than bond pair?
18. Differentiate between sigma and π-bond. Why π -bonds are more diffused than σ- bonds?
19. Define bond order. What is bond order of O2+2 andN2?
20. Differentiate between bonding molecular orbital and anti- bonding molecular orbital.
21. Why MOT is superior to VBT?
22. Ionization energy is an index to the metallic character. Why?
23. The bond angles of H2O and NH3 are not 109.5° like that of CH4, although O and N-atoms are sp3 hybridized.
24. How does electronegativity difference decide the nature of bond?
25. There is no bond 100 % ionic in chemistry. Justify.
26. What is paramagnetism? Give the reason for the paramagnetic character of oxygen.
27. He2 molecule is not formed, why?
28. Define dipole moment. Give its units.
29. The dipole moment of CO2 and CS2 is zero but that of SO2 is 1.61 Debye. Justify.
30. The abnormality of bond length and bond strength in Hl is less prominent than that of HCl.
Long Questions
1. Define atomic orbital hybridization. How can we describe the geometry of NH3 on its basis?
2. Draw energy level diagram of N2 with reference to molecular orbital theory and explain magnetic behaviour.
3. Define atomic orbital hybridization. Explain sp2 hybridization by giving example of BF3.​
4. Write the four postulates of “VSEPR” theory.
5. What is orbital hybridization? Explain the structure of CH4 molecule on the basis of hybridization theory.
6. Explain paramagnetic nature of oxygen on the basis of MOT.
7. Define ionization energy, name the factors influencing the ionization energies of elements. What is a trend of ionization energy in the periodic table?
8. Explain atomic orbital hybridization with reference to the structure of C2H2 and C2H4.
9. What is meant by VSEPR theory? Explain in detail, also discuss structures of BF3 and CH4 in the light of VSEPR theory.
10. Discuss the shape and geometry of H2O with reference to sp3 Hybridization.
11. Define electron affinity, name the factors influencing the electron affinity of elements. What is a trend of electron affinity in the periodic table?
Chapter 7
SHORT QUESTIONS
1. State 1st law of thermodynamics.
2. Define enthalpy of formation and enthalpy of atomization.
3. What is enthalpy of neutralization?
4. State Hess’s law of constant heat summation.
5. What is a thermochemical equation?
6. Why is it necessary to mention the physical states of
7. reactants and products in a thermo-chemical reaction? Apply Hess’s law to justify your answer.
8. Define lattice energy. Give example.
9. Define internal energy, system, surrounding, state and
10. state function.
11. Differentiate between exothermic and endothermic reactions.
12. Differentiate between spontaneous and non-spontaneous reactions.
13. Define enthalpy of atomization.
14. What is standard enthalpy of combustion?
15. Define standard enthalpy of solution.
16. Explain that burning of a candle is a spontaneous process.
Long Questions
1. Explain spontaneous and non-spontaneous reactions. Describe four points which differentiate them.
2. Explain the Born Haber cycle for the measurement of the lattice energy of common salt (NaCl).
3. State first law of thermodynamics. Also prove that ∆E = qv and ∆E = qp
4. Describe the measurement of enthalpy of a reaction by bomb calorimeter with diagram.
5. What is Enthalpy of a reaction? How ∆H of a reaction is measured in Laboratory by glass calorimeter?
6. State and explain Hess’s law of constant Heat summation with two examples.
Chapter 8
SHORT QUESTIONS
1. What is state of dynamic equilibrium? Differentiate
2. between reversible and irreversible reactions.
3. State law of mass action.
4. Give two applications of Kc.
5. What will be the effect of volume (pressure) change and temperature on reactions?
6. How does catalyst affect reversible reaction?
7. State Le-Chatelier principle.
8. Define pH and pOH.
9. What is common ion effect? (How is NaCl purified by common ion effect?)
10. Define buffers. How acidic and basic buffers are prepared or classify buffers?
11. What is Henderson equation for acids and buffers?
12. Define buffer capacity.
13. What are uses of buffers?
14. Why equilibrium constant value has its units for some reversible reactions but has no units for some other reactions?
15. A catalyst does not affect equilibrium constant. Comment on it.
16. In what way yield of ammonia can be increased in Haber’s process?
17. Define the solubility product constant. Derive solubility product expression of Ag2CrO4.
18. The solubility of glucose in water is increased by increasing the temperature.
19. The change of temperature disturbs the equilibrium position and the equilibrium constant of the reaction.
Long Questions
1. The solubility of PbF2 at 25°C is 0.64gdm-3. Calculate Ksp of PbF2.​
At. Mass of Pb = 207
At. Mass of F = 19
2. The solubility of CaF2 in water at 25°C is found to be 2.05 × 10-4mol dm-3. What is the value of Ksp at this temperature?
3. The solubility product of Ca(OH)2 is 6.5×10-6. Calculate the solubility of Ca(OH)2.
4. N2(g) and H2(g) combine to give NH3(g). The value of Kc in this reaction at 500°C is 6.0×10-2. Calculate the value of Kp for this reaction.
5. The solubility product of Ag2CrO4 is 2.6×10-2 at 25°C. Calculate the solubility of compound. Atomic mass of Ag = 108, Cr = 52, O = 16.
6. What is the percentage ionization of acetic acid in a solution in which 0.1 Mole of it has been dissolved per dm3 of the solution?
7. The equilibrium constant for the reaction between acetic acid and ethyl alcohol is 4.0. A mixture of 3 moles of acetic acid and one mole C2H5OH is allowed to come to equilibrium. Calculate the amount of ethyl acetate at equilibrium state in no of moles and grams. Also calculate mass of reactants left behind.
8. Calculate the pH of a buffer solution in which 0.11 molar CH3COONa and 0.09 molar acetic acid solutions are present. Kc for CH3COOH is 1.85×10-5.
9. Benzoic acid C6H5COOH, is a weak mono-basic acid (Ka=6.4×10-5mol dm-3) What is the pH of a solution containing 7.2 g of sodium benzoate in one dm3 of 0.02 mol dm-3 benzoic acid?
Chapter 9
SHORT QUESTIONS
1. Differentiate between ideal and non-ideal solution.
2. State Raoult’s law in two different ways.
3. Define molarity and molality.
4. Define mole fraction and ppm.
5. Relative lowering of v***r pressure is independent of the temperature.
6. Define solubility and solubility curves.
7. What are colligative properties? When they are obeyed? Give two applications
8. Describe ebullioscopic constant and cryoscopic constant with one example.
9. Freezing points are depressed and boiling points are elevated due to the presence of solutes.
10. The boiling point of one molal urea solution is 100.52 °C but the boiling point of two molal urea solution is less than
11. 101.04 °C.
12. Beckmann thermometer is used to note the depression in freezing point.
13. What is water of crystallization? Give two examples.
14. Boiling points of the solvents increase due to the presence of solutes.
15. Define hydration. On what factors it depends?
16. What is conjugate solution?
17. What is upper consulate temperature?
18. The aqueous solution of CH3COONa is basic in nature. Why?
19. Explain why CuSO4 give acidic solution when put in water?
20. Relative lowering of v***r pressure is independent of the temperature.
21. Colligative properties are obeyed when the solute is non–electrolyte, and also when the solutions are dilute.
22. Non-ideal solutions do not obey Raoult’s law.
23. Beckmann thermometer is used to note the depression in freezing point.
24. In summer the antifreeze solutions protect the liquid of the radiator from boiling over.
25. NaCl and KNO3 are used to lower the melting point of ice.
Long Questions
1. Discuss Raoult’s law when one component is volatile other is non-volatile.​
2. Describe Beckmann’s method for the measurement of freezing point depression with the help of diagram.
3. Discuss in detail any two examples of solutions of partially miscible liquid.
4. Differentiate between hydration and hydrolysis. Describe with two examples in each case.
5. Explain continuous and discontinuous solubility curves.
6. Explain phenol-water system in detail.
7. Differentiate between ideal and non-ideal solutions.
8. Enlist colligative properties and why some properties are colligative? Also give conditions for observing colligative properties.
9. Describe in detail the Elevation of Boiling point.
10. Write down measurement of elevation of boiling point by Landsberger’s method with diagram
11. What are colligative properties? Explain lowering of vapour pressure.
Chapter 10
SHORT QUESTIONS
1. Define electrochemistry, electrolysis, oxidation and reduction, voltaic and electrolytic cell.
2. Define oxidation state. Determine the oxidation number of element in given compound.
3. Impure Cu can be purified by electrolytic process.
4. How anodized Al can be prepared?
5. What are the products of electrolysis of aqueous and molten sodium chloride?
6. What is the function of salt bridge?
7. Define electrode potential and standard electrode potential.
8. What is electrochemical series? Give two applications of electrochemical series.
9. The standard oxidation potential of Zn is 0.76 V and its reduction potential is -0.76 V.
10. Na and K can displace hydrogen from acids but Pt, Pd and Cu cannot.
11. SHE acts as anode when connected with Cu electrode but as cathode with Zn electrode.
12. What is meant by standard hydrogen electrode?
13. Why Zn can displace hydrogen from a dilute solution of acids but Cu cannot?
14. Differentiate between primary and secondary cells.
15. A porous plate or a salt bridge is not required in lead storage cell.
16. Lead accumulator is a chargeable battery.
17. What are electrode reactions of a dry cell (alkaline battery) and silver oxide battery?
18. How power is generated by using fuel cells?
Long Questions
1. What is lead accumulator battery? Discuss its discharging process.​
2. Give four applications of electro-chemical series.
3. Describe fuel cell in detail with diagram.​
4. Explain four industrial applications of electrolysis.
5. Write note on alkaline battery.
6. Describe the construction and working of standard hydrogen electrode (SHE).
7. What is meant by Lead Accumulator explain it in detail, Give chemical equations of discharging and recharging.​
8. What is a Galvanic cell? Draw diagram. Explain its electrodes with reactions occurring on electrodes.
Chapter 11
SHORT QUESTIONS
1. Define chemical kinetics, rate of reaction, order of reaction, velocity constant, average and
2. instantaneous rate.
3. How does increase of temperature increase rate of reaction?
4. Rate of chemical reaction is an ever changing parameter under the given conditions.
5. The radioactive decay is always first order reaction.
6. The units of rate constant of second order reaction is dm3 mol1s1, but the unit rate of reaction is mole
7. dm3 s1.
8. The sum of the coefficients of a balanced chemical equation is not necessarily important to give the order of a reaction.
9. Define transition state, activation energy, half life and zero order reaction.
10. How half-life method is used to find the order of reaction?
11. How is a catalyst specific in its action?
12. How is rate of reaction affected by surface area?
13. Differentiate between homogeneous and heterogeneous catalysis.
14. Define activation of catalyst.
15. Define autocatalyst.
16. What is a negative catalyst?
17. What are enzymes? Give two examples in which enzymes act as catalyst.
18. What is catalytic poisoning? Give two examples.
Long Questions
1. Describe half-life method and method of large excess for finding the order of reaction.
2. What is enzyme catalysis? Give one example. Also give any four characteristic enzyme catalysis.​
3. Differentiate between homogeneous catalysis and heterogeneous catalysis with one example in each.
4. How does the Arrhenius equation help us to calculate the energy of activation of a reaction?
5. Describe energy of activation in detail.
6. Write down any four characteristics of catalyst.​
7. Discuss how surface area and nature of reactants affect rate of a chemical reaction.
8. How order of reaction is measured using half-life method and method of large excess?
9. Discuss any two factors affecting rate of reactions.
10. Write down any four physical methods for the determination of rate of reaction.
11. What is chemical kinetics? How do you compare chemical kinetics with chemical equilibrium?

11th ENGLISH Important topics 2025
Bk-1 Imp for SQ
Ch # 1,3,5,7,9,10,11
Bk-1 Imp for Translation
Ch # 1,2,3,4, 7,10,12,13
Imp Punctuation
Ch # 1,2,3,4,14
Bk-3 Plays SQ
Play-1 Ex. Q No. a, d, e, h
Play-2 Ex. Q No. ii, viii, ix, x.
Play-3 Ex. Q No. 4,9,12,14, 26,31,38,39
Imp Poems for SQ
Poem # 4,7,9,10,12,14,15
Imp Poems for Stanza
Poem # 1,2,3,8,13,19,20
Imp Letters
To friend: congratulation, condoling, inviting. To Father: health & progress in studies, hostel life, causes of failure.
Imp Applications
To Principal: fee concession, character certificate, readmission, fine remittance.
Imp Stories
Union, Pride, Friend in need, Necessity, greed is curse, all that glitters.
Pairs of Words from PTB book.
a, b, c, m, p, s
For Objective
prepared 5years papers

1st Year 2024

No1
Define radian and steradian? 2) Differentiate between random and systematic error? 3) Find the dimension of ŋ used in equation F= 6πŋrv.? 4) Write down the dimension of Pressure and density? 5) Find out the dimension and unit of G used in equation F=G m_(1m_2 )/r^2 6) Show that E=mc2 dimensionally correct? 7) The period of simple pendulum is measured by stop watch .what type of errors are possible in the time period? 8) Why do we find it useful to have two units for the amount of the substance kilogram and mole? 9) Give draws backs to use the period of simple pendulum as time standard? 10) How many meters is a light year? 11) How many second in a year and how many years in nano second?
No. 2
1) Two vectors have unequal magnitude can their sum be equal to zero? 2) Can a vector have component greater than vector magnitude? 3) Can a body rotate about its center of gravity under the action of its weight? 4) If one of the rectangular components of the vector is not zero, can its magnitude by zero, explain? 5) Name Three condition which makes →┬A X →┬A =0? 6) Can you add zero to null vector? 7) If →┬A +→┬B =→┬0what can you say about its components? 8) Define position vector and Unit vector?
No.3
1) Show that the rate of change of momentum is equal to force? 2) Derive expression for the range of projectile? 3) Can the velocity of an object reverse the direction when acceleration is constant? If so give an example? 4) An object is thrown vertically upward .Discuss the sign of acceleration due to gravity, relative to velocity, while the object is in air? 5) Explain the circumstance when A) a and v are parallel B) v is zero but a is not zero? 6) Differentiate between elastic and inelastic collision?
No.4
1) Calculate work done in kilo joule in lifting a mass of 10kg through a vertical height of 10m? 2) When rocket reenter the atmosphere, its nose cone become very hot? Why 3) A girl drop a cup from a certain height which breaks in to pieces. What energy changes are involved? 4) A boy use catapult to throw stone which accidently smashes a green hose window .What energy changes are involved? 5) Show that 1kWh = 3.6 x106J.
No5
1) what is centripetal force give its significance? 2) Define moment of inertia give its significance? 3) Why does diver change his body position before diving in the pool? 4) Differentiate between Tangential and angular velocity? 5) Describe what should be the minimum velocity for a satellite to orbit close to the earth around it? 6) When Mud flies off the tyre of moving bicycle, in what direction does it fly. Explain?
N0.6
1) Define Viscosity? 2) Why fog droplets appear to be suspended in air? 3) Differentiate between Streamline and turbulent flow? 4) A person is standing near a fast moving trains is there any danger the he will fall toward it? 5) Explain the working of carburetor? 6) How swing is produced in a fast moving cricket ball?
no 7
1) Write down the characteristics of S.H.M? 2) Can we realize an ideal Simple Pendulum? 3) What happens to the period of a simple pendulum if its length is doubled? And if it’s mass doubled? 4) What is meant by Phase Angle? 5) Describe some common Phenomena in which resonance plays an important role? 6) Does frequency depends up on amplitude for harmonic oscillator?
no 8
1) Explain why sound travels faster in warm air than in cold air? 2) Explain term node antinode, crest, Trough? 3) How are beats are useful in tuning musical instruments? 4) As result of distant explosion an observer senses a ground tremor and then heat the explosions. Explain the time difference? 5) What is the effect of variation of density on the speed of sound in a gas? 6) What are common features in longitudinal and transvers wave? 7) Why does sound travel faster in solid as compared to gasses?
no 9
1) Under what conditions two sources of light behave as coherent? 2) Can visbibel light produce Interference fringes? Explain? 3) An oil film spread over a wet footpath shows colours. Explain How does it happen? 4) How would distinguish between Un-polarized and plane polarized light? 5) Why Sun glasses are better than ordinary sun glasses?
NO 11
1) Does entropy of a system increase or decrease due to friction explain? 2) why the average velocity of the molecule in a gas zero but average of square of velocity is not? 3) Is It possible to convert internal energy into mechanical energy? Explain with example? 4) State 1st law of thermodynamic? How it is applicable to human body? 5) Write four postulate of kinetic theory of gases? 6) A thermo flask containing milk as a system is shaken rapidly. Does the temperature of
milk rise? 7) Can mechanical energy be converted completely in to heat energy? If so give example? 8) Prove that CP-CV=R
SAJID HUSSAIN SSTsc
Govt Central model school samnabad Lahore