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Moles and Molar Mass in Analytical Chemistry

Learn about moles, molar mass, and chemical formulas, calculate particles, masses, and molecular formulas, with practical examples, in the context of analytical chemistry.

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Moles and Molar Mass in Analytical Chemistry

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  1. Analytical Chemistry!!!!

  2. Recall • Mole—unit of measurement used to measure a quantity • 1 dozen = 12 doughnuts • 1 mole = 6.022*1023 atoms or molecules of a substance • Avogadro’s number—proportion that relates molar mass of a substance to the mass. • Used 12 g of Carbon-12 as a standard • 2 g of H2 has the same number of molecules as 12 g of C.

  3. Moles to Particles • Can find moles of a substance given the number of particles and visa versa. • Examples • Find the number of particles in 2 mol of Fluorine gas. • Find the number of moles of 5.63*1023 molecules of ammonia.

  4. Mass to Moles • 10 feathers has a different mass than 10 bowling balls • 1 mol of one substance has a different mass than 1 mol of another substance. • Molar mass—mass of 1 mol of a substance • equal to the atomic mass of an element • units are grams per mole or g/mol • when you weigh out that many grams of a substance you have 6.022 x 1023 atoms

  5. Mass to Moles • Use molar mass as a conversion factor between mass and moles. • Example: Find the number of moles of 3.00 g of C. • Example: Find the mass of 4.23 moles of C.

  6. Moles of Compounds Chemical Formula – indicates the types and number of atoms present in a compound—also represents a mol ratio of elements in the compound. • 1 mol of CH4 would consist of 6.022 x 1023 molecules of CH4 • or 1 mol C atoms, 4 mol of H atoms (remember 1 mole = atomic mass) Calculate the number of moles of each element in 5 mol of glucose (C6H12O6)

  7. Molar Mass of Compounds The mass of a briefcase could be obtained by the sum of the case + mass of each of the contents within. The mass of a mole of a compound equals the sum of the masses of each particle that makes it up. This is used to measure moles of a substance in laboratory experiments.

  8. Mole calculations • How many grams are in 6.7 mol of Aluminum Chloride? • How many moles are in 3.4*1024 molecules of Sodium Hydride? • How many molecules are in 2.35 g of Hydrogen gas? • How many grams are in 5.4*1023 molecules of Sodium Chloride? • How many Al3+ and Cl- ions are there in a 35.60 g sample of AlCl3? • What is the mass of 1 formula unit of AlCl3?

  9. Percent Composition The % composition of a compound is always the same, regardless of the size and source of the sample. • What is the percent composition of KMnO4?

  10. Aspartame is used in as an artificial sweetener in brands such as Equal and Nutrasweet. Its chemical formula is C14H18N2O5 • Calculate the mass percentage of each element in aspartame. • Calculate the mass of carbon in a 1.00 g packet of Equal, assuming it is pure aspartame.

  11. Empirical Formulas • Remember, empirical formula is the smallest whole number ratio of atoms (aka a simplified version of the molecular formulas). • 4 steps when given % composition. • Assume the sample is out of 100 g and convert percentages to masses. • Convert masses to moles. • Divide moles by lowest number of moles (if you need to convert these values to whole numbers, now is the time to do it.) • Use moles for each element as subscripts for that element.

  12. An unknown compound was found to have a percent composition as follows: 47.0 % potassium, 14.5 % carbon, and 38.5 % oxygen. What is its empirical formula?

  13. Formulas from Mass • Use masses to find moles of each substance and go from there. • A sample of indium chloride weighing 0.5000 g is found to contain 0.2404 g of chlorine. What is the empirical formula of the indium compound?

  14. Molecular formulas • Specifies that actual number of each atom in that molecule. • Remember many compounds can have the same empirical formula, but different molecular formulas.

  15. Find the empirical formula from the following molecular formulas: • C6H6 • C8H18 • WO2 • C2H6O2 • X39Y13

  16. Glucose undergoes chemical analysis and is found to have a molar mass of 180.16 g/mol and is composed of 39.90% C, 6.70% H, and 53.40% O. Find the empirical and molecular formulas.

  17. Hydrates A chemical compound containing water that is chemically combined with a substance and can usually be expelled without  changing the constitution of the substance.

  18. Uses for hydrates • Anhydrous compounds—without water • Used as dessicants—drying agents • Anhydrous compounds absorb water • Can be added to solvents to remove water

  19. Nomenclature for Hydrates Examples:CuSO4· 5H2O 1.) Use name of compound: Copper (II) Sulfate. 2.) Name hydrate with correct prefix: Di = 2 Tri = 3 Tetra = 4 Penta = 5 Hexa = 6 Hepta = 7 Octa = 8 Nona = 9 Deca = 10

  20. Practice • CuCl2∙2H2O • BeSO4∙4H2O • MgSO4∙7H2O • CaSO4∙6H2O

  21. More practice •  cobalt (II) fluoride tetrahydrate •  lithium nitrate trihydrate •  aluminum hypochlorite octahydrate •  cesium carbonate dihydrate

  22. Finding formulas for hydrates • Must use anhydrous form of compound. • 1.) Drive off water by heating. • 2.) Use mass of anhydrous compound to calculate moles. • 3.) Calculate mass of water and then find moles of water. • 4.) Use mole ratio of anhydrous compound:water to write formula.

  23. Empirical Formulas of Hydrates • What is the formula of a hydrate that is 86.7% Mo2S5 and 13.3% H2O? What is the name? • What is the formula for a hydrate that is 90.7g SrC2O4 and 9.30g H2O? • A sample of hydrated Li2SiF6 weighs 0.4813 grams. After heating, the anhydrous compound has a mass of 0.391 grams. Find the formula of the hydrate.

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