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1.2 MOLE CONCEPT

1.2 MOLE CONCEPT. Learning Outcome. At the end of this topic, students should be able : (a) Define mole in terms of mass of carbon-12 and Avogadro constant, N A . (b) Interconvert between moles, mass, number of particles, molar volume of gas at s.t.p. and room temperature.

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1.2 MOLE CONCEPT

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  1. 1.2 MOLE CONCEPT MATTER

  2. Learning Outcome At the end of this topic, students should be able : (a)Define mole in terms of mass of carbon-12 and Avogadro constant, NA. (b) Interconvert between moles, mass, number of particles, molar volume of gas at s.t.p. and room temperature. MATTER

  3. (c)Determine empirical and molecular formulae from mass composition or combustion data. MATTER

  4. (d) Define and perform calculation for each of the following concentration measurements : i) molarity (M) ii) molality (m) iii) mole fraction, X iv) percentage by mass, % w/w v) percentage by volume, %V/V MATTER

  5. (e) Determine the oxidation numberof an element in a chemical formula. (f) Write and balance : i) chemical equation by inspection method ii) redox equation by ion-electron method MATTER

  6. (g) Define limiting reactant and percentage yield. (h) Perform stoichiometric calculations using mole concept including reactant and percentage yield. MATTER

  7. Pair = 2 1.2 Mole Concept • A mole is defined as the amount of substance which contains equal number of particles (atoms / molecules / ions) as there are atoms in exactly 12.000g of carbon-12. MATTER

  8. Dozen = 12 • One mole of carbon-12 atom has a mass of exactly 12.000 grams and contains 6.02 x 1023 atoms. • The value 6.02 x 1023 is known as Avogadro Constant. NA = 6.02 x 1023 mol-1 MATTER

  9. Example MATTER

  10. Molar Mass • The mass of one mole of an element or one mole of compound is referred as molar mass. • Unit : g mol-1 • Example: - molar mass of Mg = 24 g mol-1 - molar mass of CH4 = (12 + 4)gmol-1 = 16 g mol-1 MATTER

  11. Number of Mole MATTER

  12. Example 1 MATTER

  13. MATTER

  14. Example 1 (cont…) MATTER

  15. Example 2 MATTER

  16. Example 3 MATTER

  17. 1.2.1 Mole Concept of Gases • Molar volume of any gas at STP = 22.4 dm3 mol-1 s.t.p. = Standard Temperature and Pressure Where, T = 273.15 K P = 1 atm MATTER

  18. 1 mole of gas has a volume of 22.4 dm3 at s.t.p At s.t.p, volume of gas (dm3) = number of mole X 22.4 dm3 mol-1 • 1 mole of gas has a volume of 24.0 dm3 at room temperature At room temperature, volume of gas (dm3) = number of mole X 24.0 dm3 mol-1 MATTER

  19. Example 1 MATTER

  20. Cont… from example 1 MATTER

  21. Exercise A sample of CO2 has a volume of 56 cm3 at STP. Calculate: • The number of moles of gas molecules 0.0025 mol • The number of molecular 1.506 x 1021 molecules • The number of oxygen atoms in the sample 3.011x1021atoms Note: 1 dm3 = 1000 cm3 1 dm3 = 1 L MATTER

  22. Empirical And Molecular Formulae • Empirical formula is a chemical formula that shows the simplest ratio of all elements in a molecule. • Molecular formula is a formula that show the actual number of atoms of each element in a molecule. MATTER

  23. The relationship between empirical formula and molecular formula is : Molecular formula = n ( empirical formula ) Where ; MATTER

  24. Example A sample of hydrocarbon contains 85.7% carbon and 14.3% hydrogen by mass. Its molar mass is 56. Determine the empirical formula and molecular formula of the compound. MATTER

  25. Solution : Empirical formula = CH2 MATTER

  26. n = 56 14 = 4 molecular formula = C4H8 MATTER

  27. 1.2.2 Concentration of Solution Solution • When an amount of solute dissolved completely in a solvent and it will form a homogeneous mixture. MATTER

  28. Exercise A combustion of 0.202 g of an organic sample that contains carbon, hydrogen and oxygen produce 0.361g carbon dioxide and 0.147 g water. If the relative molecular mass of the sample is 148, what is the molecular formula. Ans : C6H12O4 MATTER

  29. Units of concentration of a solution: A. Molarity B. Molality C. Mole Fraction D. Percentage by Mass E. Percentage by Volume MATTER

  30. A. Molarity (M) • The number of moles of solute per cubic decimetre (dm3) or litre (L) of solution. Note: 1 dm3 = 1000 cm3 1 L = 1000 mL MATTER

  31. Example MATTER

  32. Cont… MATTER

  33. Exercises MATTER

  34. B. Molality (m) • Molality is the number of moles of solute dissolved in 1 kg of solvent • Note: • Mass of solution = mass of solute + mass of solvent • Volume of solution ≠ volume of solvent MATTER

  35. Example 1 MATTER

  36. MATTER

  37. Example 2 MATTER

  38. MATTER

  39. Exercises MATTER

  40. MATTER

  41. C. Mole Fraction (X) • Mole fraction is the ratio of the number of moles of one component to the total number of moles of all component present. MATTER

  42. It is always smaller than 1 • The total mol fraction in a mixture (solution) is equal to one. XA + XB + XC = 1 MATTER

  43. Example 1 MATTER

  44. MATTER

  45. MATTER

  46. MATTER

  47. Example 2 MATTER

  48. MATTER

  49. D. Percentage by Mass (%w/w) • Percentage by mass is defined as the percentage of the mass of solute per mass of solution. MATTER

  50. Example 1 MATTER

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