Physical Chemistry Introduction

Section 1: Basic Concepts, Laws, and Measurements (Questions 1-10)

  1. Define matter and classify it into pure substances and mixtures with examples. Distinguish between elements and compounds.
  2. State and explain the Law of Conservation of Mass with a suitable example. How does it apply to chemical reactions?
  3. State the Law of Definite Proportions and Law of Multiple Proportions. Give one example for each.
  4. Explain Gay-Lussac’s Law of Gaseous Volumes and Avogadro’s Law. How do they support the concept of molecules?
  5. State Dalton’s Atomic Theory. Mention two postulates that were later modified and why.
  6. Define precision and accuracy in measurements. How do they differ? Give an example where a measurement is precise but not accurate.
  7. What are significant figures? Round off 2.3456 to three significant figures and explain the rules for addition/subtraction and multiplication/division.
  8. Convert 1.5 km into meters using dimensional analysis. Also, express 0.0000456 in scientific notation with appropriate significant figures.
  9. Define SI units for mass, length, time, and temperature. What is the difference between mass and weight?
  10. A sample of a compound contains 1.8 g of carbon and 0.3 g of hydrogen. Does this data support the Law of Definite Proportions? Explain.

Section 2: Atomic/Molecular Masses and Mole Concept (Questions 11-25)

  1. Define atomic mass and molecular mass. How is atomic mass expressed relative to the carbon-12 isotope?
  2. Calculate the molecular mass of glucose (C₆H₁₂O₆) and sulphuric acid (H₂SO₄). (Atomic masses: C=12, H=1, O=16, S=32)
  3. What is a mole? Calculate the number of moles in 88 g of CO₂.
  4. How many molecules are present in 18 g of water? (Avogadro’s number = 6.022 × 10²³)
  5. Calculate the mass of 2.5 moles of oxygen atoms and 0.5 moles of oxygen molecules (O₂).
  6. Find the number of atoms in 196 g of H₂SO₄. (Atomic masses: H=1, S=32, O=16)
  7. A gas occupies 11.2 L at STP. Calculate the number of moles and the mass if it is nitrogen (N₂).
  8. Calculate the number of moles of atoms in 5.6 g of iron (atomic mass = 56).
  9. What is the mass of one atom of carbon-12? Express it in grams.
  10. How many moles of electrons are present in 1 mole of Na⁺ ions?
  11. Calculate the total number of protons, electrons, and neutrons in 1 mole of ¹²C atoms.
  12. Define molar mass. Calculate the molar mass of Na₂CO₃·10H₂O.
  13. 4.5 g of a compound (molecular weight 90) is dissolved to make 250 mL solution. Find the molarity.
  14. Find the percentage composition of nitrogen in NH₃ (atomic masses: N=14, H=1).
  15. An organic compound contains 40% carbon, 6.67% hydrogen, and 53.33% oxygen by mass. Find its empirical formula.

Section 3: Empirical/Molecular Formula and Percentage Composition (Questions 26-35)

  1. A compound has empirical formula CH₂O and molecular mass 180. Find its molecular formula.
  2. 0.5 g of a hydrocarbon on combustion gives 1.1 g CO₂ and 0.9 g H₂O. Find the empirical formula.
  3. An oxide of metal contains 60% metal by mass. If the equivalent weight of the metal is 12, find the formula of the oxide.
  4. Calculate the percentage of water of crystallization in CuSO₄·5H₂O.
  5. A compound on analysis gave C=18.5%, H=1.55%, Cl=55.04%, and O=24.81%. Find the empirical formula.
  6. The ratio of masses of C and H in an organic compound is 6:1. If the molecular mass is 180, find the molecular formula (assuming it is CₓHᵧO₂).
  7. Find the mass percent of oxygen in Al₂(SO₄)₃.
  8. A sample of 5.25 g of an organic compound on combustion gave 6.08 g CO₂ and 2.49 g H₂O. Find the empirical formula.
  9. Vitamin C has 40.92% C, 4.58% H, and 54.50% O. Its molecular mass is 176. Find the molecular formula.
  10. Calculate the percentage by mass of nitrogen in urea (NH₂CONH₂).

Section 4: Stoichiometry, Limiting Reagent, and Concentration Terms (Questions 36-50)

  1. Balance the equation: Fe + H₂O → Fe₃O₄ + H₂. Calculate the moles of H₂ produced from 1 mole of Fe.
  2. In the reaction 2Al + 6HCl → 2AlCl₃ + 3H₂, calculate the volume of H₂ gas at STP produced from 27 g of Al.
  3. What is a limiting reagent? In a reaction, 2 g H₂ and 32 g O₂ are mixed. Identify the limiting reagent and calculate the mass of water formed.
  4. 10 g of CaCO₃ on heating gives 4.4 g CO₂. Calculate the percentage yield.
  5. Define molarity and molality. How do they differ? Calculate the molality of a solution containing 18 g glucose in 500 g water.
  6. A solution is 2 M H₂SO₄ with density 1.2 g/mL. Calculate its molality and mass percentage.
  7. Calculate the normality of a 500 mL solution containing 49 g H₂SO₄.
  8. What is ppm? Express 5 mg of solute in 1 kg solution in ppm.
  9. In a reaction, 1584 g (NH₄)₂SO₄ reacts with Ca(OH)₂. Calculate the mass of products formed (assume complete reaction; use standard atomic masses).
  10. 6.977 g of a mixture containing Ba²⁺ ions gives 0.4123 g BaSO₄ precipitate. Find the mass percentage of Ba in the sample.
  11. Calculate the number of moles of methane required to produce 11 g CO₂ on complete combustion.
  12. Find the mass of 1 mole of electrons in grams.
  13. A 1.8 g sample of water contains how many water molecules? Compare with 18 moles of water.
  14. In the reaction KMnO₄ + KI (weakly basic), calculate moles of I₂ released per mole of KMnO₄ (balanced equation required).
  15. Explain stoichiometry with an example. Calculate the amount of CaO produced when 100 g CaCO₃ is heated (assuming 90% purity and 80% yield).