Some Basic Concepts of Chemistry Class 11 Notes

• Law of conservation of mass — mass is neither created nor destroyed in a chemical reaction.
• Law of constant proportions — a compound always contains the same elements in the same ratio by mass.
• Law of multiple proportions — when two elements form more than one compound, the masses of one that combine with a fixed mass of the other are in a simple whole-number ratio.
• Law of reciprocal proportions — relates combining ratios of elements.
• Gay-Lussac's law — volumes of reacting gases (at constant T, P) are in simple whole-number ratios.

A balanced chemical equation gives the mole ratios of reactants and products. Stoichiometry uses these ratios to convert between amounts of substances — from moles to mass, volume or number of particles.

• $\text{Molarity (M)}:$
• moles per litre of solution.
• $\text{Molality (m)}:$
• moles per kg of solvent.
• $\text{Mole fraction (x)}:$
• ratio of moles of one component to total moles.
• $\text{Mass percent:}$
• mass of solute per 100 g of solution.
• $\text{ppm:}$
• parts per million, for very dilute solutions.

Example: What is the mass of 0.5 moles of ${\text{CO}}_2$? (Molar masses: C = 12, O = 16)

Solution: $M({\text{CO}}_2)=12+2\times 16=44\text{g/mol}$. So $m=nM=0.5\times 44=22\text{g}$.

Exercises (NCERT Q1.1–Q1.36)

• 1.1) Calculate the molar mass of the following:- (i) H₂O (ii) CO₂ (iii) CH₄
• 1.2) Calculate the mass per cent of different elements present in sodium sulphate (Na₂SO₄).
• 1.3) Determine the empirical formula of an oxide of iron, which has 69.9% iron and 30.1% dioxygen by mass.
• 1.4) Calculate the amount of carbon dioxide that could be produced when (i) 1 mole of carbon is burnt in air, (ii) 1 mole of carbon is burnt in 16 g of dioxygen, (iii) 2 moles of carbon are burnt in 16 g of dioxygen.
• 1.5) Calculate the mass of sodium acetate (CH₃COONa) required to make 500 mL of 0.375 molar aqueous solution. Molar mass of sodium acetate is 82.0245 g mol⁻¹.
• 1.6) Calculate the concentration of nitric acid in moles per litre in a sample which has a density, 1.41 g mL⁻¹ and the mass per cent of nitric acid in it being 69%.
• 1.7) How much copper can be obtained from 100 g of copper sulphate (CuSO₄)?
• 1.8) Determine the molecular formula of an oxide of iron, in which the mass per cent of iron and oxygen are 69.9 and 30.1, respectively. Given that the molar mass of the oxide is 159.69 g mol⁻¹ (Fe = 55.85, O = 16.00).
• 1.9) Calculate the atomic mass (average) of chlorine using the following data: ³⁵Cl has natural abundance 75.77% and molar mass 34.9689; ³⁷Cl has natural abundance 24.23% and molar mass 36.9659.
• 1.10) In three moles of ethane (C₂H₆), calculate the following: (i) Number of moles of carbon atoms, (ii) Number of moles of hydrogen atoms, (iii) Number of molecules of ethane.
• 1.11) What is the concentration of sugar (C₁₂H₂₂O₁₁) in mol L⁻¹ if its 20 g are dissolved in enough water to make a final volume up to 2L?
• 1.12) If the density of methanol is 0.793 kg L⁻¹, what is its volume needed for making 2.5 L of its 0.25 M solution?
• 1.13) Pressure is determined as force per unit area of the surface. The SI unit of pressure, pascal, is as shown below: 1 Pa = 1 N m⁻². If mass of air at sea level is 1034 g cm⁻², calculate the pressure in pascal.
• 1.14) What is the SI unit of mass? How is it defined?
• 1.15) Match the following prefixes with their multiples: (i) micro, (ii) deca, (iii) mega, (iv) giga, (v) femto — against the multiples 10⁶, 10⁹, 10⁻⁶, 10⁻¹⁵, 10.
• 1.16) What do you mean by significant figures?
• 1.17) A sample of drinking water was found to be severely contaminated with chloroform, CHCl₃, supposed to be carcinogenic in nature. The level of contamination was 15 ppm (by mass). (i) Express this in per cent by mass. (ii) Determine the molality of chloroform in the water sample.
• 1.18) Express the following in the scientific notation: (i) 0.0048 (ii) 234,000 (iii) 8008 (iv) 500.0 (v) 6.0012
• 1.19) How many significant figures are present in the following? (i) 0.0025 (ii) 208 (iii) 5005 (iv) 126,000 (v) 500.0 (vi) 2.0034
• 1.20) Round up the following upto three significant figures: (i) 34.216 (ii) 10.4107 (iii) 0.04597 (iv) 2808
• 1.21) The following data are obtained when dinitrogen and dioxygen react together to form different compounds: (i) 14 g dinitrogen with 16 g dioxygen (ii) 14 g dinitrogen with 32 g dioxygen (iii) 28 g dinitrogen with 32 g dioxygen (iv) 28 g dinitrogen with 80 g dioxygen. (a) Which law of chemical combination is obeyed by the above experimental data? Give its statement. (b) Fill in the blanks in the following conversions: (i) 1 km = ____ mm = ____ pm (ii) 1 mg = ____ kg = ____ ng (iii) 1 mL = ____ L = ____ dm³
• 1.22) If the speed of light is 3.0 × 10⁸ m s⁻¹, calculate the distance covered by light in 2.00 ns.
• 1.23) In a reaction A + B₂ → AB₂, identify the limiting reagent, if any, in the following reaction mixtures: (i) 300 atoms of A + 200 molecules of B₂ (ii) 2 mol A + 3 mol B₂ (iii) 100 atoms of A + 100 molecules of B₂ (iv) 5 mol A + 2.5 mol B₂ (v) 2.5 mol A + 5 mol B₂
• 1.24) Dinitrogen and dihydrogen react with each other to produce ammonia according to the following chemical equation: N₂(g) + H₂(g) → 2NH₃(g) (i) Calculate the mass of ammonia produced if 2.00 × 10³ g dinitrogen reacts with 1.00 × 10³ g of dihydrogen. (ii) Will any of the two reactants remain unreacted? (iii) If yes, which one and what would be its mass?
• 1.25) How are 0.50 mol Na₂CO₃ and 0.50 M Na₂CO₃ different?
• 1.26) If 10 volumes of dihydrogen gas reacts with five volumes of dioxygen gas, how many volumes of water vapour would be produced?
• 1.27) Convert the following into basic units: (i) 28.7 pm (ii) 15.15 pm (iii) 25365 mg
• 1.28) Which one of the following will have the largest number of atoms? (i) 1 g Au(s) (ii) 1 g Na(s) (iii) 1 g Li(s) (iv) 1 g of Cl₂(g) (Atomic masses: Au = 197, Na = 23, Li = 7, Cl = 35.5)
• 1.29) Calculate the molarity of a solution of ethanol in water, in which the mole fraction of ethanol is 0.040 (assume the density of water to be one).
• 1.30) What will be the mass of one ¹²C atom in g?
• 1.31) How many significant figures should be present in the answer of the following calculations? (i) 0.02856 × 298.15 × 0.112 / 0.5785 (ii) 5 × 5.364 (iii) 0.0125 + 0.7864 + 0.0215
• 1.32) Use the data given in the following table to calculate the molar mass of naturally occurring argon isotopes: ³⁶Ar, isotopic molar mass 35.96755 g mol⁻¹, abundance 0.337%; ³⁸Ar, isotopic molar mass 37.96272 g mol⁻¹, abundance 0.063%; ⁴⁰Ar, isotopic molar mass 39.9624 g mol⁻¹, abundance 99.600%.
• 1.33) Calculate the number of atoms in each of the following (i) 52 moles of Ar (ii) 52 u of He (iii) 52 g of He.
• 1.34) A welding fuel gas contains carbon and hydrogen only. Burning a small sample of it in oxygen gives 3.38 g carbon dioxide, 0.690 g of water and no other products. A volume of 10.0 L (measured at STP) of this welding gas is found to weigh 11.6 g. Calculate (i) empirical formula, (ii) molar mass of the gas, and (iii) molecular formula.
• 1.35) Calcium carbonate reacts with aqueous HCl to give CaCl₂ and CO₂ according to the reaction, CaCO₃(s) + 2HCl(aq) → CaCl₂(aq) + CO₂(g) + H₂O(l). What mass of CaCO₃ is required to react completely with 25 mL of 0.75 M HCl?
• 1.36) Chlorine is prepared in the laboratory by treating manganese dioxide (MnO₂) with aqueous hydrochloric acid according to the reaction 4HCl(aq) + MnO₂(s) → 2H₂O(l) + MnCl₂(aq) + Cl₂(g). How many grams of HCl react with 5.0 g of manganese dioxide?

Answers at a glance

1. 1
   Q1.1 → (i) 18.02 u (ii) 44.01 u (iii) 16.043 u.
2. 2
   Q1.2 → Na = 32.4%, S = 22.6%, O = 45.05%.
3. 3
   Q1.3 → Empirical formula = Fe₂O₃.
4. 4
   Q1.4 → (i) 44 g (ii) 22 g (iii) 22 g of carbon dioxide.
5. 5
   Q1.5 → 15.38 g of sodium acetate is required.
6. 6
   Q1.6 → Concentration of nitric acid = 15.44 mol L⁻¹.
7. 7
   Q1.7 → 39.81 g of copper can be obtained from 100 g of copper sulphate.
8. 8
   Q1.8 → Empirical formula = Fe₂O₃, n = 1, so the molecular formula of the oxide is Fe₂O₃.
9. 9
   Q1.9 → Average atomic mass of chlorine = 35.45 u ≈ 35.4527 u.
10. 10
    Q1.10 → (i) 6 moles of carbon atoms (ii) 18 moles of hydrogen atoms (iii) 1.807 × 10²⁴ molecules (18.07 × 10²³ molecules).
11. 11
    Q1.11 → Molar concentration of sugar = 0.02925 mol L⁻¹.
12. 12
    Q1.12 → Volume of methanol required = 25.2 mL (≈ 25.22 mL).
13. 13
    Q1.13 → Pressure of air at sea level = 1.01 × 10⁵ Pa (1.01332 × 10⁵ Pa).
14. 14
    Q1.14 → SI unit of mass = kilogram (kg); 1 kg is the mass of the international prototype kilogram.
15. 15
    Q1.15 → (i) micro = 10⁻⁶ (ii) deca = 10 (iii) mega = 10⁶ (iv) giga = 10⁹ (v) femto = 10⁻¹⁵.
16. 16
    Q1.16 → Significant figures are the total number of digits in a number including the last digit whose value is uncertain (the first uncertain digit).
17. 17
    Q1.17 → (i) 1.5 × 10⁻³ % by mass (ii) molality = 1.26 × 10⁻⁴ m.
18. 18
    Q1.18 → (i) 4.8 × 10⁻³ (ii) 2.34 × 10⁵ (iii) 8.008 × 10³ (iv) 5.000 × 10² (v) 6.0012 × 10⁰.
19. 19
    Q1.19 → (i) 2 (ii) 3 (iii) 4 (iv) 3 (v) 4 (vi) 5 significant figures.
20. 20
    Q1.20 → (i) 34.2 (ii) 10.4 (iii) 0.0460 (iv) 2810.
21. 21
    Q1.21 → (a) Law of multiple proportions. (b)(i) 1 km = 10⁶ mm = 10¹⁵ pm (ii) 1 mg = 10⁻⁶ kg = 10⁶ ng (iii) 1 mL = 10⁻³ L = 10⁻³ dm³.
22. 22
    Q1.22 → Distance covered by light in 2.00 ns = 0.600 m.
23. 23
    Q1.23 → (i) B₂ (ii) A (iii) none (stoichiometric) (iv) B₂ (v) A.
24. 24
    Q1.24 → (i) 2.43 × 10³ g of NH₃ (ii) Yes, dihydrogen (H₂) remains (iii) 571.4 g of H₂ unreacted.
25. 25
    Q1.25 → 0.50 mol Na₂CO₃ is simply 53 g of the substance; 0.50 M Na₂CO₃ is 53 g dissolved in 1 L of solution (0.50 mol per litre).
26. 26
    Q1.26 → 10 volumes of water vapour would be produced.
27. 27
    Q1.27 → (i) 2.87 × 10⁻¹¹ m (ii) 1.515 × 10⁻¹¹ m (iii) 2.5365 × 10⁻² kg.
28. 28
    Q1.28 → 1 g of Li(s) has the largest number of atoms (8.6 × 10²²).
29. 29
    Q1.29 → Molarity of the ethanol solution = 2.31 M ≈ 2.314 M.
30. 30
    Q1.30 → Mass of one ¹²C atom = 1.99 × 10⁻²³ g ≈ 1.993 × 10⁻²³ g.
31. 31
    Q1.31 → (i) 3 (ii) 4 (iii) 4 significant figures.
32. 32
    Q1.32 → Molar mass of naturally occurring argon = 39.948 g mol⁻¹ (≈ 39.95 g mol⁻¹).
33. 33
    Q1.33 → (i) 3.13 × 10²⁵ atoms of Ar (ii) 13 atoms of He (iii) 7.83 × 10²⁴ atoms of He.
34. 34
    Q1.34 → (i) Empirical formula = CH (ii) Molar mass ≈ 26 g mol⁻¹ (iii) Molecular formula = C₂H₂.
35. 35
    Q1.35 → 0.96 g of CaCO₃ is required (≈ 0.964 g).
36. 36
    Q1.36 → 8.4 g of HCl react completely with 5.0 g of manganese dioxide.

MCQ Questions

1 Mark Questions

1. 1
   Define the mole. (Ans: 6.022 × 10²³ particles (Avogadro's number) — the amount of substance containing as many entities as there are atoms in 12 g of C-12)
2. 2
   State Avogadro's law. (Ans: equal volumes of all gases at the same temperature and pressure contain the same number of molecules)
3. 3
   What is meant by empirical formula? (Ans: the simplest whole-number ratio of atoms of different elements in a compound)
4. 4
   Define molarity. (Ans: the number of moles of solute dissolved per litre of solution; unit mol L⁻¹)

2 Mark Questions

1. 1
   Calculate the mass per cent of each element in Na₂SO₄ (Na = 23, S = 32, O = 16). (Ans: Na = 32.4%, S = 22.5%, O = 45.1%)
2. 2
   State the law of multiple proportions with an example. (Ans: two elements combine in more than one simple mass ratio — e.g. C and O form CO and CO₂, masses of O being 16/32 = 1:2)
3. 3
   How many significant figures are present in 0.00509, 500.0 and 5.00 × 10³? (Ans: 3, 4, 3)

3 Mark Questions

1. 1
   A compound contains 69.9% Fe and 30.1% O. Find its empirical formula, and its molecular formula if its molar mass is 159.69 g mol⁻¹. (Ans: empirical Fe₂O₃, n = 1, hence molecular formula Fe₂O₃)
2. 2
   Calculate the number of atoms in (i) 52 moles of Ar (ii) 52 u of He (iii) 52 g of He. (Ans: 3.13 × 10²⁵; 13 atoms; 7.83 × 10²⁴ atoms)
3. 3
   In a reaction A + B₂ → AB₂, identify the limiting reagent when 2 mol A is mixed with 3 mol B₂, and 5 mol A with 2.5 mol B₂. (Ans: for 2:3, A is limiting (needs 2 mol B₂ but has 3); for 5:2.5, B₂ is limiting)

5 Mark Questions

1. 1
   Dinitrogen and dihydrogen react to give ammonia: N₂ + 3H₂ → 2NH₃. If 2.00 × 10³ g of N₂ reacts with 1.00 × 10³ g of H₂, calculate the mass of ammonia formed and the mass of the reactant left unreacted. (Ans: 2.43 × 10³ g NH₃; H₂ left unreacted, 571.4 g)
2. 2
   Define the four laws of chemical combination. Calculate the molarity of an ethanol–water solution in which the mole fraction of ethanol is 0.040 (density of water = 1 g mL⁻¹). (Ans: molarity ≈ 2.31 M; laws — conservation of mass, definite/constant proportions, multiple proportions, gaseous volumes)