Chemistry is the most fascinating science subject since it involves a large number of chemical compounds, their properties, and reactions. For FSc students, chemistry is the most important subject. Chemistry is extremely important in our daily life. If you practice guess papers as well, you can achieve full marks in chemistry. We are providing 1st-year chemistry guess paper 2022 to students in the 11th grade. This guess paper includes important long questions of chemistry 1st-year chapter wise 2022″ This guess paper of Chemistry 1st year is for all the boards of Punjab. You can download the Guess paper questions in pdf form.

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Chemistry 1st year guess paper 2022                                 

  1st Year                             CHEMISTRY     

GUESS PAPER 2022

FOR ALL BOARDS OF PUNJAB

MCQs from back exercise of each Chapter of Text Book.

SHORT QUESTIONS

CHAPTER 1

  1. Define molecular formula. Write down the relationship between molecular formula and empirical formula
  2. Why N2 and CO have same number of electrons, protons and neutrons?
  3. Why two moles of NaOH are required for complete neutralization of one mole of H2SO4.
  4. What are isotopes? Why they have the same chemical properties?
  5. Write two stoichiometric assumptions?
  6. Define Avogadro’s number with two examples/
  7. Differentiate between ion and molecular ion.
  8. Why law of conversation of mass has to be obeyed during stoichiometric calculations?
  9. Why do many chemical reactions taking place in our surroundings involve limiting reactants?
  10. Why no individual Neon atom in the sample of neon element has a mass of 20.18 amu?
  11. The removal of an electron from a neutral atom is an endothermic process. Explain with reason.
  12. One mg of K2CrO4 has thrice the number of ions than the number of formula units when ionized in water, justify.
  13. Actual yield is always less than the theoretical yield. Give two reasons.
  14. Mg atom is twice heavier than a carbon atom. Explain.
  15. What are Molecular Ions? How are they generated?
  16. Define gram Formula. Give two examples,
  17. Differentiate between atom and molecule.
  18. What is the function of the ionization chamber in a mass spectrometer?
  19. What is mass spectrum?
  20. Why do isotopes have the same chemical properties but different physical properties?

CHAPTER 2

  1. What is the basic principle of Crystallization?
  2. Write down the uses of chromatography
  3. Differentiate between the stationary phase and mobile phase in chromatographic technique?
  4. Define sublimation. Give one example.
  5. What is Rf value and why it has no unit?
  6. State distribution law.
  7. Why is there a need to crystallize the crude product?
  8. What is sublimation? Give one example of sublime solid.
  9. Differentiate between adsorption chromatography and partition chromatography.
  10. What is solvent extraction?
  11. Write down the name of eight solvent used in crystallization.
  12. What is difference between qualitative analysis and quantitative analysis?
  13. Why concentrated HCl and KMnO4 solutions cannot be filtered by Gooch’s crucible?
  14. Explain briefly two methods for drying of the crystallized substance?
  15. How undesirable colours are removed from the crystals?

CHAPTER 3

  1. Why water vapours do not behave ideally at 273K?
  2. Define Avogadro’s law with two suitable examples?
  3. Calculate the value of ‘R’ gas constant in SI units
  4. Define critical temperature and critical pressure giving one example in each case.
  5. What do you mean by natural plasma and artificial plasma?
  6. Pilots feel uncomfortable breathing at higher attitude. Give reason.
  7. Write down any two applications of plasma.
  8. Give four fundamental postulates of kinetic molecular theory of gases.
  9. Why the dipole moment of SO2 is 1.61 D but that of SO3 is zero?
  10. Gases deviate from ideal behaviour at low temperature and high pressure. Give reasons.
  11. SO2 is comparatively non-ideal at 273 K but behaves ideally at 317K. Explain.
  12. Write two faulty assumptions of kinetic molecular theory of gases.
  13. Explain Boyle’s law with the help of KMT.
  14. What is the difference between Diffusion and Effusion?
  15. Why deep sea divers take oxygen mixed with an inert gas like He?
  16. State Graham’s law of diffusion. Write its mathematical form.
  17. Derive Charles’s law by kinetic equation of gases..
  18. Derive molecular mass of a gas by general gas equation.
  19. Give quantitative definition of Charles’s law?
  20. Write characteristics of plasma?

CHAPTER 4

  1. Define Anisotropy and Allotropy
  2. Define (a)Dipole-dipole forces                   (b) Hydrogen bonding
  3. Why diamond is hard and electrically neutral?
  4. Ionic solids do not conduct electricity in solid state?
  5. Write two advantages of vacuum distillation?
  6. Define allotropy. Give its one example?
  7. Define transition temperature with two examples?
  8. Give reason that ice floats on water.
  9. Evaporation causes cooling. Explain.
  10. Vacuum distillation can be used to avoid decomposition of a sensitive liquid. Explain with reason.
  11. Explain why evaporation takes place at all temperatures.
  12. Why heat of sublimation of iodine is very high?
  13. Define transition temperature with an example.
  14. Describe that ionic crystals are highly brittle?
  15. Give two uses of liquid crystals.
  16. Why earthenware vessels keep water cool?
  17. Why the boiling point of water is higher than HF?
  18. Ice occupies more space then water. give reason.
  19. Water and ethanol can mix is all proportions. Give reason.
  20. Describe cleaning action of soaps and detergents on the basis of H-bonding.
  21. Transition temperature is exhibited by both elements and compounds. Explain.
  22. Why HF has less acidic strength than HI?
  23. Why molar heat of vaporization (Hv) is greater than the molar heat of fusion (∆Hf)?
  24. Cleavage is an anisotropic behavior. Explain it.
  25. Why one feels a sense of cooling under the fan after a bath?
  26. Define liquid crystal? Write down two applications of liquid crystals.
  27. Define isomorphism and polymorphism with example.
  28. Define liquid crystal with one example.
  29. Define lattice energy with an example
  30. Why the value of boiling point of noble gases increases from top to bottom within a group?
  31. Explain the term unit cell dimension.
  32. Why water is liquid at room temperature but H2S and H2Se are gases, comments.
  33. Give reason for the lowest boiling point of hydride of group IV-A elements.
  34. Define polarizability. How it affects London dispersion forces?

CHAPTER 5

  1. Why the positive rays are also called canal rays?
  2. Cathode rays possess momentum. Justify it.
  3. Give two defects of Bohr’s atomic model.
  4. State Hund’s rule.
  5. How neutrons are used in treatment of cancer?
  6. Why e/m value of the cathode rays is just equal to that of electron?
  7. Write the two drawbacks of Rutherford’s model of atoms.
  8. Define Heisenberg’s uncertainty principle and give its mathematical expression.
  9. What is orbital? Draw the shape of P orbital.
  10. Differentiate between line spectrum and continuous spectrum.
  11. Whichever gas is used in the discharge tube the nature of cathode rays remain same? Why?
  12. Distribute the electrons in the orbitals of 29 Cu and 24 Cr.
  13. Give any two properties of neutrons.
  14. Give reason for the production of positive rays.
  15. Calculate mass of an electron from its e/m value.
  16. What is Plank’s quantum theory?
  17. State Mosley’s law? Write two importance of Moseley law.
  18. Pressure can effect the production of Cathode Rays.
  19. Define Aufbau principle and Pauli’s Exclusion principal.
  20. Why pressure is reduced in discharge tube experiment?
  21. Differentiate between slow and fast neutron.
  22. Define Zeeman’s effect and Stark effect.
  23. Why the electrons move faster nearer to the nucleus in an orbit of a smaller radius?
  24. Describe (n + l) rule for distribution of electrons.
  25. Why the potential energy of bonded electron has a negative value?

CHAPTER 6

  1. What is “Octet rule”? Give two examples of the compounds which do not obey Octet rule.
  2. How Sigma and pi bonds are formed.
  3. The dipole moment of CO2 and CS2 are zero, but that of SO2 is 1.61D. justify it.
  4. Why the abnormality of bond strength in HI is less prominent than that of HCl.
  5. Define “Dipole moment Give its various units.
  6. Why is the second ionization energy greater than first one?
  7. Define bond energy and give one example?
  8. Rate of reaction of ionic compounds is faster than covalent compound. Why?
  9. Why molecular orbital theory is superior to valence bond theory?
  10. Write two main postulates of VSEPR theory?
  11. Why cationic radius is smaller than parent atom?
  12. Why CO is polar and CO2 is non-polar?
  13. Define electron affinity. And give an example.
  14. π- bonds are more diffused than δ-bonds. Give reason.
  15. What is bond order? Give an example.
  16. Define electronegativity. Give its trend in the periodic table.
  17. Define coordinate covalent bond and give one example.
  18. Represent the molecular orbitals of N2 molecule in the increasing order of energy.
  19. How does ionization energy vary in periodic table?
  20. Why the energy of anti-bonding molecular orbital is higher than corresponding bonding molecular orbitals?
  21. How the nature of a chemical bond is predicted with the help of electronegativity values of two bonded atoms?
  22. No bond in chemistry is 100% ionic. Justify it.
  23. The bond angles of H2O and NH3 are not 109.5° like that of CH4 although oxygen and nitrogen atoms are SP3 – hybridized. Why?
  24. Differentiate between atomic orbital and molecular orbital.
  25. How the type of bonding affects ‘solubility of compounds?
  26. Sketch the molecular orbital picture of O2.
  27. Explain that Helium is diamagnetic.
  28. Bond distance is the compromise distance between two atoms.
  29. Give two postulates of Bohr’s atomic mode!.
  30. Why anionic radius is greater than atomic radius? 

CHAPTER 7

  1. Define “System ” and Surroundings.
  2. The burning of candle is a spontaneous process. Justify it.
  3. Define enthalpy of formation with one example.
  4. Differentiate between internal energy change and enthalpy change?
  5. Is it true that a non-spontaneous processes never happens in the universe. Explain?
  6. Why it is necessary to mention the physical states of reactants and products in a thermochemical reaction?
  7. Differentiate between law of conservation of energy and Hess’s law?
  8. Differentiate between endothermic and exothermic reactions.
  9. Differentiate between internal energy and enthalpy?
  10. What do you mean by standard enthalpy of atomization?
  11. Define heat of neutralization with an example.
  12. What is State and State function? Explain with example.
  13. Prove that ∆E = qv
  14. Define enthalpy of solution?
  15. State the Hess’s law of constant heat summation.
  16. What are spontaneous and non-spontaneous processes? Give one example of each.

CHAPTER 8

  1. State Lechatlier’s Principle. What happens when pressure is increased on the reaction

N2 + 3H2 → 2NH3

  1. What is the effect of increase in temperature on the yield of the product for the reaction.

SO2 + O2 → SO3 + heat

  1. What is buffer solution? Give an example.
  2. How does a buffer act? Explain with an example.
  3. Define solubility product. Derive solubility product expression for sparingly soluble compound Ag2CrO4.
  4. Calculate the pH of 10-4 mole dm3 of Ba(OH)2.
  5. How buffer solutions are prepared? Discuss two methods.
  6. Define pH and pOH with mathematical expression.
  7. Define effect of catalyst on equilibrium constant.
  8. What is meant by conjugate acid and conjugate base?
  9. Write a note on buffer capacity?
  10. Write two applications of equilibrium constant.
  11. Write two uses of buffer solution.
  12. State law of mass action.
  13. How the values of equilibrium constant helps to predict the direction of a reversible reaction?
  14. What is common ion effect? Give an example.
  15. What is Henderson’s equation?
  16. Write two uses of buffer solution.
  17. Differentiate between reversible and irreversible reactions.
  18. Write down optimum conditions for the preparation of ammonia?
  19. What are Lowry-Bronsted acids and bases? Give an example.
  20. Why change of volume disturbs the equilibrium position for some of the gaseous phase reactions but not the equIribrium constant.?
  21. What are applications of buffer in daily life?
  22. Derive ionic product of water and what is its value at 25°C?.

CHAPTER 9

  1. Define Hydration energy of ions
  2. The concentration in terms of molality is independent of temperature but molarity depends upon temperature. Justify it.
  3. Justify that boiling point of solvents increases due to presence of solutes.
  4. Define hydration and hydrolysis?
  5. Calculate the percentage by weight of NaCl is dissolved in 20 g of water?
  6. What is part per million? Write its mathematical expression?
  7. Define solubility? What is solubility curve? Name its two types.
  8. Why colligative properties of solutions are obeyed when solute is non-electrolyte and also when the solutions are dilute?
  9. Define zeotropic mixtures. Give one example.
  10. What is upper consulate temperature and give one example?
  11. Define Molarity and MoIality.
  12. Differentiate between ideal and non-ideal solutions.
  13. One molal solution of glucose is dilute as compared to one molar solution of glucose. Justify it.
  14. What is meant by Water of Crystallization? Give examples.
  15. State colligative properties. Name important colligative properties.
  16. What is critical solution temperature? Give an example.
  17. Give two statements of Raoult’s law?
  18. Define upper consulate temperature with example.
  19. Define the ebullioscopic constant with an example.
  20. Define cryoscopic constant with an example?

CHAPTER 10

  1. Write reactions for the electrolysis of fused sodium chloride
  2. What are secondary cells. Give their two examples.
  3. Write down any four rules for balancing of reodx equations by ion Electron method?
  4. Define electrode potential. Give one example?
  5. Define electrolysis give one example.
  6. A salt bridge maintains etectrical neutrality in the cell. Why?
  7. Lead accumulators is a chargeable battery. Comment.
  8. Calculate the oxidation number of chromium in; (a) K2CrO4 (b) K2Cr2O7
  9. Define electrolytic cell? Give example.
  10. Write functions of the salt bridge?
  11. Find out the oxidation state of Mn in (a) KMnO4 (ii) K2MnO4
  12. How impure Cu can be purified by an electrolytic process.
  13. A voltaic cell is a reversible cell. State.
  14. Define electrochemical series.
  15. Write down reactions taking place at the electrodes on the discharging of Nickle-Cadmium Cell.
  16. Write two advantages of fuel cells.
  17. SHE acts as an anode when connected with a copper electrode but as a cathode with a Zinc electrode. Explain it?
  18. How anodized aluminum is prepared? Give the advantages of anodization of Al.
  19. Mg can displace hydrogen from acids but Pt cannot do so. Give reason.

CHAPTER 11

  1. Differentiate between “Instantaneous rate ” and average rate ” of a reaction.
  2. What is Zero order reaction? Give one example
  3. Write down four characteristics of a catalyst.
  4. How a catalyst is specific in its action?
  5. What is poisoning of catalysis? Give an example.
  6. The order of reactions may be in fractions. Justify with the help of an example.
  7. Name two physical methods used to determine the rate of a reaction.
  8. Differentiate between negative catalysis and autocatlysis?
  9. How surface area affects the rate of reaction’? Give one example.
  10. Define order of reaction with one example?
  11. Give two characteristics of enzyme catalysis.
  12. Define energy of activation. How is it affected by temperature?
  13. Differentiate between rate and rate constant of reaction.
  14. Define the rate of reaction and catalysis?
  15. Define the first order reaction with example?
  16. The rate of a reaction is an ever changing quantity. Comment.
  17. 50% of a hypothetical first order reaction completes in one hour. The remaining 50% requires more than one hour. Why?
  18. Define half life period. How is it used to determine the order of reaction?
  19. What is specific rate constant or velocity constant?
  20. Differentiate between Homogeneous and Heterogeneous Catalysis.
  21. The sum of the coefficients of a balanced chemical equation is not necessarily important to give the order of reaction justify.
  22. Radioactive reactions are always first order reactions. Give reason?

LONG QUESTIONS

Chapter 1

  1. Define limiting reactant. How limiting reactant can be determined? Explain it with an example?
  2. Write down the steps to calculate empirical formula.
  3. Explain the determine the empirical formula by combustion analysis?
  4. Define yield. How do we calculate the percentage yield of chemical reaction? Also mention the factors which are responsible for low yield of products.
  5. Define stoichiometry. Give its assumptions. Mention two important laws which help to perform the stoichiometric calculation.

Chapter 3

  • Numericals of Examples & Exercise of Chapter 3

Chapter 4

  1. Define vapour pressure. How vapour pressure is measured by manometric method?
  2. What are ionic solids? Write four properties of ionic solids?
  3. What are metallic crystals? Discuss the electron gas theory of metallic bond.
  4. Define liquid crystals and give their three uses.
  5. Define covalent solids with example. Write four properties of covalent solids.
  6. What are molecular solids? Give examples and explain their properties. 

Chapter 5

  1. Explain Rutherford’s model of the atom.
  2. Define quantum numbers. Discuss briefly all quantum numbers.
  3. What is spectrum? Explain atomic emission and atomic absorption spectrum?
  4. Give postulates of Bohr’s atomic model
  5. Write down the properties of cathode rays.
  6. Write down Millikan’s oil drop method for the measurement of the charge of an electron.
  7. Write the J.J. Thomson method to measure the e/m value of electron.
  8. Derive the equation for the radius of the nth orbit of the hydrogen atom of Bohr’s theory. 

Chapter 6

  1. Describe postulates of valence shell electron pair repulsion theory.
  2. Explain the molecular orbital structures of following molecules on the basis of the MOT:
  • N2 (Nitrogen) (ii) O2 (oxygen)
  1. Define electron affinity. Name the factors affecting it. How does it very in the periodic table?
  2. Write main postulates of molecular orbital theory.
  3. Give the defects of Bohr’s atomic model.
  4. Draw the shapes of following molecules according to VSEPR theory.
  5. BeCl2       ii) BF3       iii) H2O           iv)NH3
  6. Explain ionization energy by giving one example? Also discuss its periodic trends.
  7. Define hybridization. Explain the geometry of ethyne on the basis of sp hybridization.

OR Explain structure of CH4, NH3 and H2O on the basis of hybridization theory.

OR What is hybridization? Explain Sp2 hybridization with example.

Chapter 7

  1. Explain enthalpy of a reaction and prove ∆H =qp
  2. Define enthalpy and prove ∆H = qp
  3. What is molar heat of combustion? How it is measured by bomb calorimeter?
  4. Detine enthalpy of reaction. How is it measured by glass calorimeter?
  5. State first law of thermodynamics. Give its mathematical expression and apply this law to prove that ∆E = qv
  6. State Hess’s law of constant heat summation. Give two examples.

Chapter 8

  • Numericals of Examples & Exercise of Chapter 8.

Chapter 9

  1. Define the following terms

molarity,   molality, mole fraction, parts per million(PPM).

  1. Define Colligative Properties. How molecular mass of solute is determined by lowering of vapour pressure?
  2. Give three statements of Raoult’s law and also mention how Raoult’s law helps in determining the ideality of a solution.
  3. Differentiate between Hydration and Hydrolysis with examples.
  4. Write four differences between ideal solution and non ideal solutions.
  5. Define solubility and solubility curves and discuss continuous solubility curve.

Chapter 10

  1. Write four industrial applications of electrolysis?
  2. What is the oxidation state? Give rules for assigning it for example.
  3. Explain the construction and working of standard Hydrogen Electrode (SHE)?
  4. What is a standard hydrogen electrode (SHE)? How it is used to measure the electrode potential of zinc?
  5. Explain the construction and working of the fuel cell.
  6. Describe the galvanic cell explaining the functions of electrodes and salt bridge
  7. What is an electrochemical series? give two uses?
  8. Describe the electrolysis of molten sodium chloride and a concentrated solution of sodium chloride.
  9. Discuss the working and chemistry of Alkaline battery and fuel cells.
  10. Describe Nickel Cadmium Cell (rechargeable)?

Chapter 11

  • Describe half life method for determination of order of reaction.
  • How does Arrhenius equation helps us to calculate the energy of activation of a reaction?
  • Define order et reaction and explain 2nd and zero order reactions.
  • Define enzyme catalysis. Write three characteristics of enzyme catalysis.
  • What is catalysis? Differentiate between homogeneous and heterogenous catalysis give one example in each.
  • What is order of reaction? Describe two methods for finding the order of reaction. (Half life method and method of large excess)

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