1 / 14

The importance of Avogadro’s number

The importance of Avogadro’s number. 6.0221421 X 10 23. Molar Mass. The mass of a single atom of an element (in amu) is numerically equal to the mass (in grams) of 1 mol of that element. 1 atom of 12 C has a mass of 12 amu 1 mol of 12 C has a mass of 12.0 g

ostinmannual
Download Presentation

The importance of Avogadro’s number

An Image/Link below is provided (as is) to download presentation Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript

  1. The importance of Avogadro’s number 6.0221421 X 1023

  2. Molar Mass • The mass of a single atom of an element (in amu) is numerically equal to the mass (in grams) of 1 mol of that element. • 1 atom of 12C has a mass of 12 amu • 1 mol of 12C has a mass of 12.0 g • 1 molecule of H2O has a mass of 18 amu • 1 mol of H2O has a mass of 18 g

  3. Molar Mass (cont) • The mass in grams of one mole of a substance is call the molar mass of the substance. • The molar mass (g/mol) of any substance is always numerically equal to it formula weight (in amu). • Example • Molar Mass of glucose (C6H12O6) = 180.0 g

  4. Converting Mass to Moles and vice versa • Simply performed by dimensional analysis. • Example: • Calculate the number of moles of glucose in 5.380 g • Calculate the mass, in grams, of 0.433 mol of Ca(NO3)2

  5. Converting Mass into Number of Particles and vice versa • Example • How many atoms of Cu are in a penny, assume penny weighs 3 g and is 100% Cu? • How many glucose molecules are in 5.23 g of glucose?

  6. Empirical Formulas from Analyses • Remember the empirical formula for a substance tells us the relative number of atoms of each element it contains. • Also the ratio of the number of moles of each element in a compound gives the subscripts in a compound’s empirical formula. • Example • Mercury and chlorine combine to form a compound that is 73.9% mercury and 26.1% chlorine by mass. • Take 100g of the compound and determine the number of moles of each.

  7. Empirical formula • Ascorbic Acid (vitamin C) • 40.92% C, 4.58% H, 54.5% O • A sample of methylbenzoate weighs 5.325 g and contains 3.758 g of carbon, 0.316 g of H and 1.251 g of oxygen. What is the empirical formula?

  8. Molecular Formula from Empirical Formula • The subscripts in the molecular formula of a substance are always a whole number multiple of the corresponding subscripts in its empirical formula. • multiplier = molecular weight/empirical formula weight • Example • Mesitylene has empirical formula C3H4 and has an experimentally determined molecular weight of 121 amu. What is its formula weight?

  9. Combustion Analysis • One technique for determining empirical formulas is combustion analysis. • When a compound containing carbon and hydrogen is completely combusted the sample is converted to carbon dioxide and water. The amounts of carbon dioxide and water are measured and the results are weighed. • P 97

  10. Quantitative Information from Balanced Equations • The coefficients in a chemical equation represent the relative numbers of molecules in a reaction. • The mole concept allows us to convert this information to masses. • Table p 99

  11. Example • When 1.57 mol O2 reacts with H2 to form H2O, how many moles of H2 are consumed in the process? • 2H2 + O2 2 H2O • We know that for every mole of O2 we consume 2 moles of H2. Thus we have • 1.57 mol O2 * (2 mol H2/1 mol O2) • Giving us 3.14 mol of H2

  12. Study Sample Exercise 3.16 (p 100) • Sample Exercise 3.17 (p 100)

  13. Limiting reactants • Very rarely do we have exact quantities of each of the reactants so the reaction stops as soon as one of the reactants is used up. • The reactant that is completely consumed in the reaction is called either the limiting reactant or limiting reagent. • The other reactant(s) is (are) the excess reactant or excess reagent.

  14. Page 104 Sample Exercise • Exercise 3.18 • Exercise 3.19

More Related