Short Q&A for Some Basic Concepts of Chemistry (Class 11 NCERT)

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Some Basic Concepts of Chemistry

112 Short Q&A • NCERT Class 11 Chemistry

Class 11 JEE / NEET

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112 Short Q&A for Basic Concepts of Chemistry

Part 1: Nature of Matter and Laws of Chemical Combination
1. Q1: What is the scientific study of matter's composition, structure, properties, and reactions?
   A: Chemistry.
2. Q2: What are the three common states of matter?
   A: Solid, Liquid, and Gas.
3. Q3: How is matter classified at the macroscopic level?
   A: Mixtures and Pure Substances.
4. Q4: Define a Pure Substance.
   A: Substance having a fixed composition and single set of properties.
5. Q5: What is the difference between an element and a compound?
   A: Element cannot be broken down; Compound can be broken down chemically into elements.
6. Q6: Give an example of a Homogeneous Mixture.
   A: Salt solution or Air.
7. Q7: Give an example of a Heterogeneous Mixture.
   A: Sand and water or Gunpowder.
8. Q8: State the Law of Conservation of Mass.
   A: Mass can neither be created nor destroyed in a chemical reaction.
9. Q9: Who proposed the Law of Conservation of Mass?
   A: Antoine Lavoisier.
10. Q10: State the Law of Definite Proportions.
    A: A pure chemical compound always contains the same elements combined in a fixed proportion by mass.
11. Q11: Who stated the Law of Definite Proportions?
    A: Joseph Proust.
12. Q12: State the Law of Multiple Proportions.
    A: If two elements combine to form more than one compound, the masses of one element that combine with a fixed mass of the other are in a simple whole-number ratio.
13. Q13: Who proposed the Law of Multiple Proportions?
    A: John Dalton.
14. Q14: What is the basis of Gay-Lussac's Law of Gaseous Volumes?
    A: Volumes of reacting gases and products are in a simple whole-number ratio (at constant T and P).
15. Q15: State Avogadro's Law.
    A: Equal volumes of all gases (at the same T and P) contain an equal number of molecules.
16. Q16: What is the fundamental idea of Dalton's Atomic Theory?
    A: Matter consists of indivisible atoms.
17. Q17: Which of Dalton's postulates was proved wrong by later discoveries?
    A: Atoms are indivisible (we now know of subatomic particles).
Part 2: Concepts of Mass, Moles, and Stoichiometry
18. Q18: What is the SI unit of mass?
    A: Kilogram (kg).
19. Q19: What is the SI unit of amount of substance?
    A: Mole (mol).
20. Q20: Define Atomic Mass Unit (amu) or u.
    A: 1/12th of the mass of a C-12 atom.
21. Q21: Define Average Atomic Mass.
    A: The weighted average of the atomic masses of all naturally occurring isotopes of an element.
22. Q22: Define Molecular Mass.
    A: The sum of the atomic masses of all atoms in a molecule.
23. Q23: Define Formula Mass.
    A: The sum of the atomic masses of ions in a formula unit of an ionic compound (e.g., NaCl).
24. Q24: What is the mass of one mole of a substance called?
    A: Molar Mass (g mol^-1).
25. Q25: Define a Mole.
    A: The amount of substance that contains as many particles as there are atoms in 12 g of C-12 isotope.
26. Q26: What is the value of Avogadro's constant (N_A)?
    A: 6.022 × 10²³ particles/mol.
27. Q27: How many molecules are in 1 mole of water (H₂O)?
    A: 6.022 × 10²³ molecules.
28. Q28: What is the volume occupied by 1 mole of an ideal gas at STP?
    A: 22.7 L (Standard Temperature and Pressure, 273.15 K, 1 bar).
29. Q29: What is the molar mass of CO₂?
    A: 12.01 + 2(16.00) = 44.01 g/mol.
30. Q30: Define Percentage Composition.
    A: The percentage of the mass of each element in a compound.
31. Q31: Define Empirical Formula.
    A: The simplest whole-number ratio of atoms of different elements present in a compound.
32. Q32: Define Molecular Formula.
    A: The actual number of atoms of different elements present in a molecule of a compound.
33. Q33: What is the empirical formula of glucose (C₆H₁₂O₆)?
    A: CH₂O.
34. Q34: How are molecular mass and empirical formula mass related?
    A: Molecular Mass = n × Empirical Formula Mass (n is a simple integer).
35. Q35: What is Stoichiometry?
    A: The calculation of the mass/moles/volumes of reactants and products in a balanced chemical reaction.
36. Q36: What is the Limiting Reagent (or Reactant)?
    A: The reactant that is completely consumed during the reaction and limits the amount of product formed.
37. Q37: What is the Excess Reagent (or Reactant)?
    A: The reactant present in an amount greater than required by the limiting reagent.
Part 3: Solutions and Concentration Terms
38. Q38: Define Mass Percentage of a component in a solution.
    A: (Mass of component / Total mass of solution) × 100.
39. Q39: Define Molarity (M).
    A: Moles of solute dissolved per liter of solution (mol L^-1).
40. Q40: Write the formula for Molarity.
    A: M = n/V (where n is moles, V is volume in L).
41. Q41: Define Molality (m).
    A: Moles of solute dissolved per kilogram of solvent (mol kg^-1).
42. Q42: Which concentration term is independent of temperature?
    A: Molality (since it involves only masses, which don't change with T).
43. Q43: Define Mole Fraction (χ).
    A: The ratio of the number of moles of one component to the total number of moles of all components.
44. Q44: What is the sum of the mole fractions of all components in a solution?
    A: Always 1 (χ_A + χ_B + ⋯ = 1).
45. Q45: What is the relationship between Molarity and Volume of solution upon dilution?
    A: M₁V₁ = M₂V₂ (Dilution Equation).
46. Q46: If you dissolve 10 g of NaOH (molar mass ≈ 40) in 1 L of water, what is its Molarity?
    A: 0.25 M (0.25 moles / 1 L).
47. Q47: If you dissolve 1 mole of glucose in 1 kg of water, what is its Molality?
    A: 1 mol/kg or 1 m.
Part 4: Scientific Notation, Units, and Measurements
48. Q48: What is the standard notation for expressing very large or very small numbers?
    A: Scientific Notation (N × 10ⁿ).
49. Q49: Convert 0.00045 into scientific notation.
    A: 4.5 × 10^-4.
50. Q50: Convert 602200000000000000000000 into scientific notation.
    A: 6.022 × 10²³.
51. Q51: What does Precision refer to in measurements?
    A: The closeness of a set of measurements to each other.
52. Q52: What does Accuracy refer to in measurements?
    A: The closeness of a single measurement to the true value.
53. Q53: What are Significant Figures?
    A: All certain digits plus one uncertain digit.
54. Q54: How many significant figures are in 0.0025?
    A: Two (leading zeros are not significant).
55. Q55: How many significant figures are in 200.0?
    A: Four (trailing zeros after the decimal point are significant).
56. Q56: How many significant figures are in 500?
    A: One (unless specified with a decimal point).
57. Q57: What is the result of 4.5 + 2.15 to the correct number of significant figures?
    A: 6.7 (result limited by 4.5 which has one digit after decimal).
58. Q58: What are the seven Base SI Units?
    A: metre, kilogram, second, ampere, kelvin, mole, candela.
59. Q59: What is the SI unit of temperature?
    A: Kelvin (K).
60. Q60: What is the relationship between Celsius (t_C) and Kelvin (T_K)?
    A: T_K = t_C + 273.15.
61. Q61: What is the SI unit of density?
    A: kg m^-3 (or g cm^-3).
62. Q62: What is the conversion factor between 1 L and cm³?
    A: 1 L = 1000 cm³ (1 dm³).
63. Q63: What method is used to convert units by tracking dimensions?
    A: Dimensional Analysis (or Factor Label Method).
Part 5: Stoichiometry Calculations (Numerical Focus)
64. Q64: What is the formula to calculate the number of moles (n) from mass (m)?
    A: n = m / Molar Mass (MM).
65. Q65: How many moles are in 4.0 g of NaOH (MM = 40 g/mol)?
    A: 0.1 mole.
66. Q66: What is the mass of 0.5 mole of H₂O (MM = 18 g/mol)?
    A: 9 g.
67. Q67: What is the volume of 2 moles of O₂ gas at STP?
    A: 2 × 22.7 = 45.4 L.
68. Q68: Calculate the number of atoms in 1 mole of He.
    A: 6.022 × 10²³ atoms.
69. Q69: Calculate the number of O atoms in 1 molecule of H₂SO₄.
    A: 4 atoms.
70. Q70: Calculate the number of H atoms in 1 mole of H₂O.
    A: 2 × N_A = 1.2044 × 10²⁴ atoms.
71. Q71: If 1 g of H₂ reacts with O₂ to form H₂O, how much H₂O is formed? (Balanced Eqn: 2H₂ + O₂ → 2H₂O).
    A: 9 g (H₂ MM=2; 1 g H₂ = 0.5 mole. 0.5 × 18 g/mol = 9 g).
72. Q72: What does 2 L of N₂ reacting with 6 L of H₂ yield (volume basis)? (N₂ + 3H₂ → 2NH₃).
    A: 4 L of NH₃.
73. Q73: If H₂ and O₂ react (2H₂ + O₂ → 2H₂O) and H₂ is the limiting reactant, which reactant is in excess?
    A: Oxygen (O₂).
Part 6: Important Definitions and Miscellaneous Concepts
74. Q74: What is the property of CO₂ being 44.01 g/mol?
    A: Its Molar Mass.
75. Q75: What is the freezing point of water in Kelvin?
    A: 273.15 K (0°C).
76. Q76: What is the boiling point of water in Kelvin?
    A: 373.15 K (100°C).
77. Q77: Is the number of atoms in 12 g of C-12 exactly 6.022 × 10²³?
    A: Yes, by definition of the mole.
78. Q78: What is the SI unit of volume?
    A: m³ (cubic meter).
79. Q79: Which unit of pressure is equal to 10⁵ Pa?
    A: 1 bar.
80. Q80: What is the standard temperature for STP in Celsius?
    A: 0°C.
81. Q81: What is the term for the number of valence shell electrons?
    A: Valency (for main group elements).
82. Q82: How is the mass of a single atom of H calculated?
    A: Molar Mass / N_A (1.008 g/mol / 6.022 × 10²³).
83. Q83: What are the two types of uncertainties in measurements?
    A: Random and Systematic errors.
84. Q84: What is the term for matter that is uniform in composition?
    A: Homogeneous (can be a solution or a pure substance).
85. Q85: What is the term for matter that has components visually separable?
    A: Heterogeneous (a mixture).
86. Q86: What did Dalton say about atoms of the same element?
    A: They are identical in all respects (later proved incorrect due to isotopes).
87. Q87: What is the percentage of O in H₂O? (MM=18.02 g/mol).
    A: 88.81% ((16.00 / 18.02) × 100).
88. Q88: What is the relationship between the empirical formula and the simplest ratio of moles in a compound?
    A: They are the same; the EF is derived from the mole ratio.
89. Q89: How does Molarity change if the solution volume is doubled by adding water?
    A: Molarity is halved.
90. Q90: What is the minimum number of significant figures in the result of any calculation?
    A: It is limited by the term with the fewest significant figures used in the calculation.
91. Q91: Define Standard Conditions of Temperature and Pressure (STP).
    A: 273.15 K and 1 bar.
92. Q92: What is the maximum number of decimal places allowed in the final answer of an addition problem?
    A: It is limited by the term with the fewest decimal places.
93. Q93: Name the chemical principle behind balancing a chemical equation.
    A: Law of Conservation of Mass.
94. Q94: What is the definition of 1 Pascal (Pa)?
    A: 1 Newton per meter² (1 N/m²).
95. Q95: What is the mass of one electron?
    A: 9.109 × 10^-31 kg (negligible in chemical mass calculations).
96. Q96: What is the SI unit for luminous intensity?
    A: Candela (cd).
97. Q97: What is the SI unit for electric current?
    A: Ampere (A).
98. Q98: Give an example of a derived SI unit.
    A: Density (kg/m³), Volume (m³), or Force (Newton).
99. Q99: What is the conversion factor for 1 L to m³?
    A: 1 L = 10^-3 m³.
100. Q100: What is the difference between 1 bar and 1 atm?
     A: 1 atm (≈ 1.01325 bar) is slightly greater than 1 bar.
101. Q101: What is the term for the degree of exactness of a measurement?
     A: Precision.
102. Q102: When do the empirical and molecular formulas become the same?
     A: When the simplest whole-number ratio of atoms is also the actual number of atoms (e.g., CH₄).
103. Q103: What is the factor 10^-6 represented by?
     A: Micro (μ).
104. Q104: What is the factor 10⁹ represented by?
     A: Giga (G).
105. Q105: Which Law explains why 1 g of H₂ reacts completely with 8 g of O₂ to form 9 g of H₂O?
     A: Law of Definite Proportions (fixed mass ratio 1:8).
106. Q106: If 10 g of CaCO₃ produces 4.4 g of CO₂, how much CaO is produced?
     A: 10 - 4.4 = 5.6 g (Law of Conservation of Mass).
107. Q107: Give the scientific notation for the speed of light (300,000,000 m/s).
     A: 3.0 × 10⁸ m/s.
108. Q108: Is Molarity preferred over Molality in technical applications?
     A: Yes, Molarity is easier to measure (volume-based).
109. Q109: What is the percentage purity of a sample that contains 8 g of pure NaCl in a 10 g total sample?
     A: 80%.
110. Q110: What is the mass of 1 molecule of O₂?
     A: 32 / 6.022 × 10²³ grams (Molar Mass divided by N_A).
111. Q111: What is the number of significant figures in the constant 2 in the formula Circumference = 2πr?
     A: Infinite (it is an exact number).
112. Q112: What is the density of water at 4°C?
     A: 1.0 g/cm³ (or 1000 kg/m³).

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