1. Quantitative Chemistry Chapter 3

2. Objectives Learning objective 1.2 The student is able to select and apply mathematical routines to mass data to identify or infer the composition of pure substances and/or mixtures. Learning objective 1.3 The student is able to select and apply mathematical relationships to mass data in order to justify a claim regarding the identity and/or estimated purity of a substance. Learning objective 1.4 The student is able to connect the number of particles, moles, mass, and volume of substances to one another, both qualitatively and quantitatively. Learning objective 1.14 The student is able to use data from mass spectrometry to identify the elements and the masses of individual atoms of a specific element

3. Objectives, Part 2 Learning objective 3.1 Students can translate among macroscopic observations of change, chemical equations, and particle views. Learning objective 3.3 The student is able to use stoichiometric calculations to predict the results of performing a reaction in the laboratory and/or to analyze deviations from the expected results. Learning objective 3.4 The student is able to relate quantities (measured mass of substances, volumes of solutions, or volumes and pressures of gases) to identify stoichiometric relationships for a reaction, including situations involving limiting reactants and situations in which the reaction has not gone to completion. Learning objective 3.6 The student is able to use data from synthesis or decomposition of a compound to confirm the conservation of matter and the law of definite proportions.

4. Atomic Mass Atoms are so small, it is difficult to discuss how much they weigh in grams. Use atomic mass units. An atomic mass unit (amu) is 1/12 the mass of a carbon- 12 atom. (Adopted in 1961) The decimal numbers on the table are atomic masses in amu. Sometimes abbreviated as u.

5. Decimals?? Because they are based on averages of atoms and of isotopes. Can figure out the average atomic mass from the mass of the isotopes and their relative abundance. Add up the percent as decimals times the masses of the isotopes.

6. Mass Spectrometer

7. Mass Spec Data

8. The Mole The mole is a number. A very large number, but still, just a number. 6.022 x 10 23 of anything is a mole The number of atoms in exactly 12 grams of carbon-12.

9. Molar Mass Mass of 1 mole of a substance. Often called molecular weight. To determine the molar mass of an element, look on the table. To determine the molar mass of a compound, add up the molar masses of the elements that make it up.

10. Find the Molar Mass

11. Why not both?

12. Percent Composition Percent of each element a compound is composed of. Find the mass of each element, divide by the total mass, multiply by a 100.

13. Find the % Comp.

14. Working Backwards… From percent composition, you can determine the empirical formula. Empirical Formula - the lowest ratio of atoms in a molecule. Based on mole ratios of the constituent elements. A sample is 59.53% C, 5.38%H, 10.68%N, and 24.40%O, what is its empirical formula?

15. Pure O 2 in CO 2 is absorbed H 2 O is absorbed Sample is burned completely to form CO 2 and H 2 O

16. Try This! A 0.2000 gram sample of a compound (vitamin C) composed of only C, H, and O is burned completely with excess O 2. 0.2998 g of CO 2 and 0.0819 g of H 2 O are produced. What is the empirical formula?

17. Empirical To Molecular Formulas Empirical is lowest ratio. Molecular is actual molecule. Ratio of empirical to molar mass will tell you the molecular formula. Must be a whole number because...

18. Example A compound is made of only sulfur and oxygen. It is 69.6% S by mass. Its molar mass is 184 g/mol. What is its formula?