1. Stoichiometry

2. Dalton’s Model (1803) All matter is ultimately composed of atoms. Atoms can be neither subdivided nor changed into one another. Atoms can be neither created or destroyed. Atoms of a particular substance are identical. Chemical change is the union or separation of atoms.

3. Mole concept Basis of stoichiometry A mole is a count – 6.02 x 10 23 particles (Avogadro’s number) Why this strange number? If you take the molar mass of any compound in grams, this quantity will contain Avogadro’s number of particles, 1 mole.

4. Formulae Show the number of atoms of each element in a covalent molecule or relative number of units in an ionic compound. Take the molar mass in grams. You get one mole of molecules. Chemical substances react by count, not mass. However, in lab, we measure quantities by mass. The mole concept gives us a way of relating count to mass.

5. Example Which of the following weighs most: 50 g of iron, 5 moles of nitrogen gas, 0.10 mole of silver, 1 x 10 23 atoms of radium ?

7. Example A compound was analyzed to give the following percentage composition: potassium (26.57%), chromium (35.36%), and oxygen (38.07%). What is the formula for the compound?

9. Example Compound Q contains 3 atoms of L and 2 atoms of M for every atom of X. In a given reaction starting with 3.6 x 10 22 atoms of X and 0.12 mole of L, 3.6 g of M is required. What is the atomic mass of M?

11. Chemical Equations Atoms are neither destroyed or created in a chemical reaction. We need the same number of atoms of each element on both sides of the equation. Chemical equations show counts of atoms involved, not masses. Need to convert moles to masses and vice versa.

12. Types of Chemical Reactions Combustion C 2 H 6 O + 3 O 2  2 CO 2 + 3 H 2 O Replacement (displacement) Zn + CuSO 4  Cu + ZnSO 4 Double displacement AgNO 3 + NaCl  NaNO 3 + AgCl (s)

13. Types of Chemical Reactions Acid-base neutralization HCl + NaOH  NaCl + H 2 O Combination (single product) 2 SO 2 + O 2  2 SO 3 Decomposition 2 HgO  2 Hg + O 2

14. Example Balance FeS 2 + O 2  Fe 2 O 3 + SO 2

15. Example Balance C 7 H 6 O 2 + O 2  CO 2 + H 2 O

16. example Balance Ba(NO 3 ) 2 + Na 3 PO 4  Ba 3 (PO 4 ) 2 + NaNO 3

17. Example Before the public became serious about pollution, it was common to improve the performance of gasoline by the addition of lead compounds. A particular 100-octane aviation gasoline used 1.00 cm 3 of tetraethyl lead, (C 2 H 5 ) 4 Pb, per liter of product. The pure tetraethyl lead is made as follows: 4 C 2 H 5 Cl + 4 NaPb  (C 2 H 5 ) 4 Pb + 4 NaCl + 3 Pb How many grams of ethyl chloride, C 2 H 5 Cl, are used to make enough tetraethyl lead for 1 liter of gasoline?

19. Example A solution containing 2.00 g of mercury (II) nitrate was added to a solution containing 2.00 g of sodium sulfide. Calculate the mass of mercury (II) sulfide that was formed.

21. Reactions and Energy Reactions will either produce energy or require energy. Exothermic – reactions liberate energy. Products have lower energy content than reactants Endothermic – reactions consume energy. Products have higher energy content than reactants.

22. Activation Energy Energy in endothermic reactions is NOT the same as activation energy. Activation energy relates only to the rate of reactions. Catalysts provide an alternative reaction path to avoid need for high activation energy. Even though activation energy may well be required, a reaction can still be very exothermic.