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Atomic Masses

Atomic Masses. Balanced equations tell us the relative numbers of molecules of reactants and products. C + O 2  CO 2 1 atom of C reacts with 1 molecule of O 2 to make 1 molecule of CO 2. Atomic Masses (cont.).

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Atomic Masses

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  1. Atomic Masses • Balanced equations tell us the relative numbers of molecules of reactants and products. C + O2 CO2 1 atom of C reacts with 1 molecule of O2 to make 1 molecule of CO2

  2. Atomic Masses (cont.) • If you want to know how many O2 molecules you will need, or how many CO2 molecules you can make, you will need to know how many C atoms are in the sample. • You can either count the atoms (not possible) or weigh the carbon and calculate the number of atoms. You can do this if you know the mass of one atom of carbon. • Atomic masses allow us to convert weights into numbers of atoms. • The unit is the amu (atomic mass unit) 1 amu = 1.66 x 10-24g

  3. Atomic Masses (cont.)

  4. Atomic Masses (cont.) • If our sample of carbon weighs 3.00 x 1020 amu, how many carbon atoms are present? Since our equation tells us that 1 C atom reacts with 1 O2 molecule, if we have 2.50 x 1019 C atoms, how many molecules of O2 will we need?

  5. Moles • The mass of 1.0 mole of an element is equal to its atomic mass in grams. • One mole = 6.022 x 1023 units. This number is called Avogadro’s number. • One mole of carbon-12 atoms weighs 12.0 g and has6.02 x 1023 atoms. • One atom of carbon-12 weighs 12.0 amu.

  6. Compute the number of moles and the number of atoms in 10.0 g of Al. • Use the periodic table to determine the mass of 1 mole of Al. • Use this as a conversion factor for convertinggrams to moles.

  7. Use Avogadro’s Number to determine the number of atoms in 1 mole. • Use this as a conversion factor for converting moles to atoms.

  8. Compute the number of moles and the mass in g of 2.23 x 1023 atoms of Al. • Use Avogadro’s Number to determine the number of atoms in 1 mole. • Use this as a conversion factor for converting atoms to moles.

  9. Use the periodic table to determine the mass of 1 mole of Al. • Use this as a conversion factor for converting moles to grams.

  10. Molar Mass • Molar mass: the mass in grams of one mole of a compound • The relative weights of molecules can be calculated from atomic masses. water = H2O = 2(1.008 amu) + 16.00 amu = 18.02 amu • 1 mole of H2O will weigh 18.02 g, therefore the molar mass of H2O is 18.02 g. • 1 mole of H2O will contain 16.00 g of oxygen and 2.02 g of hydrogen.

  11. A 12 oz Pepsi contains approximately 20 g of sucrose, whose molecular formula is C12H22O11. How many molecules of sucrose are present?

  12. Percent Composition • Percentage by mass of each element in a compound • Can be determined from the formula of the compound. • The percentages may not always total to 100% due to rounding.

  13. Determine the percent composition of each element in the compound C2H5OH. • Determine the mass of each element in 1 mole of the compound. • Determine the molar mass of the compound by adding the masses of the elements.

  14. Divide the mass of each element by the molar mass of the compound and multiply by 100%

  15. Calculate the mass % of each element in Ca(NO3)2.

  16. Calculate the mass of N in grams present in 5g of glycine (C2H5O2N)

  17. Empirical Formulas • Empirical formula: the simplest, whole-number ratio of atoms in a molecule • Molecular formula: a multiple of the empirical formula • Examples: Compound Empirical Formula Molecular Formula Benzene CH C6H6 Hydrogen peroxide HO H2O2 Water H2O H2O

  18. Determine the empirical formula of benzopyrene, C20H12 • Find the greatest common factor (GCF) of the subscripts. • Divide each subscript by the GCF to get the empirical formula. • Empirical Formula =

  19. Determining the empirical formula from grams • Obtain the mass of each element in g. • Determine the number of moles of each element. • Divide the number of moles of each element by the smallest number of moles present. • If a non-integer results, multiply all the numbers by the smallest number integer that converts all of the numbers into whole number integers.

  20. When 3.269 g of Zn is heated in pure oxygen, the sample gains 0.8 g of oxygen in forming the corresponding oxide. Calculate its empirical formula.

  21. When 1 g of chromium metal is heated with chlorine gas, 3.045g of a chromium chloride salt results. Calculate the empirical formula of the salt.

  22. Determine the empirical formula of acetic anhydride if its percent composition is 47% carbon, 6.0% hydrogen, and 47% oxygen. • Convert the percentages to grams by assuming you have 100 g of the compound. • Convert the grams to moles. • Divide each by the smallest number of moles.

  23. If any of the ratios is not a whole number, multiply all the ratios by a factor to make it a whole number. • Use the ratios as the subscripts in the empirical formula.

  24. Calculate the empirical formula of a compound that contains 40% carbon, 6.7% hydrogen, and 53.3% oxygen.

  25. Molecular Formulas • The molecular formula is a multiple of the empirical formula. • To determine the molecular formula you must know the empirical formula and the molar mass of the compound.

  26. Determine the molecular formula of benzopyrene if it has a molar mass of 252 g and an empirical formula of C5H3 Determine the molar mass of the empirical formula. Divide the given molar mass of the compound by the molar mass of the empirical formula and round to the nearest whole number. Multiply the empirical formula by the calculated factor to give the molecular formula.

  27. If a compound with an empirical formula CH2O has a molar mass of 180.16, what is its molecular formula?

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