1. Stoichiometry Quantitative nature of chemical formulas and chemical reactions Chapter 3 (Sections 3.3 - 3.7)

2. Atomic Mass Scale Atomic mass units (amu) are convenient units to use when dealing with extremely small masses of individual atoms 1 amu = 1.66054 x 10 -24 g 1 g = 6.02214 x 10 23 amu By definition, the mass of C-12 is exactly 12 amu

3. Average Atomic Mass (Atomic Weight) We average the masses of isotopes using their masses and relative abundances to give the average atomic mass of an element. Naturally occurring C consists of 98.892% C-12 (12.00 amu) and 1.108% C-13 (13.00335 amu) The average mass of C is (0.98892)(12.00 amu) + (0.01108)(13.00335) = 12.01 amu Atomic weights are listed on the periodic table

4. Calculating % Abundance Chlorine is made up of two isotopes, Cl-35 atomic mass = 34.969 amu and Cl- 37 atomic mass = 36.966 amu. Given chlorine’s atomic weight of 35.453, what is the % abundance of each isotope? 34.969(x) + (36.966) (1-x) = 35.453 34.969x + 36.966 – 36.966x = 35.453 x =.7576 75.76% Cl-35 and 24.24% Cl-37

5. Formula & Molecular Weights Formula Weight is the sum of atomic weights for the atoms present in the chemical formula Molecular Weight is the sum of atomic weights of the atoms in a molecule as shown in the molecular formula Sample Exercise 3.5 page 80

6. Percentage Composition from Formulas Obtained by dividing the mass contributed by each element (number of atoms times atomic weight) by the formula weight of the compound and multiplying by 100 Sample Exercise 3.6 page 80

7. The Mass Spectrometer page 81 Mass spectrometers are pieces of equipment designed to measure atomic and molecular masses accurately. The sample is converted to positive ions by collisions with a stream of high-energy electrons upon entering the spectrometer. The charged sample is accelerated using an applied voltage. The ions are then passed into an evacuated tube through a magnetic field. The magnetic field causes the ions to be deflected by different amounts depending on their mass – more mass, less deflection. The ions are then detected.

8. The Mole The mole is a convenient measure of chemical quantities. 1 mole of something is 6.0221421 x 10 23 of that thing. This number is called Avagadro’s number. Thus 1 mole of carbon atoms = 6.0221421 x 10 23 carbon atoms

9. Molar Mass mass in grams of 1 mole of a substance expressed in units of g/mol formula weights are numerically equal to the molar mass Sample Exercise 3.8 page 84

10. Interconverting Masses, Moles, & Number of Particles To convert between grams and moles, we use the molar mass To convert between moles and particles (atoms, molecules, or ions) we use Avogadro’s number Sample Exercises 3.7, 3.9, 3.10, 3.11 pages 82-86

11. Empirical Formula Gives the relative number of atoms of each element in the substance Can be calculated from mass percent data Sample Exercise 3.12 page 87

12. Molecular Formula Actual number of atoms of each element in one molecule of the substance Whole number multiple of empirical formula Sample Exercise 3.13 page 88

13. Combustion Analysis A sample containing C, H, and O is combusted in excess oxygen to produce CO 2 & H 2 O Can be used to determine empirical formula or percent composition

14. Quantitative Information from Balanced Equations Coefficients can be interpreted as the relative numbers of molecules or formula units in the reaction as well as the relative number of moles See Figure 3.14 page 92 Sample Exercises 3.14, 13.15 pages 92 & 93

15. Limiting Reactant(s) The reactant(s) that is completely consumed in a reaction Limits or determines the amount of product that will be formed The other reactant(s) that is left over is called the excess reactant Sample Exercise 3.16 and 3.17 pages 95-96 How much of the excess reactant is left over?

16. Percent Yield (actual yield / theoretical yield) X 100 Actual is the amount of product recovered in the lab Theoretical is the amount predicted from stoichiometry Sample Exercise 3.18 page 97