2. The atomic mass of an element is given in atomic mass units (amu) Found on the periodic table The mass of 1 mole of an element is equal to its atomic mass (periodic table) in grams

3. The sum of the atomic masses of its atoms Expressed in amu Example: 1 molecule of CO 2 =

4. Mass Examples Find the atomic or formula mass of the following substances: (To the nearest whole number) 1.C 2.Br 3.Cl 2 4.H 2 O 5.Ca 3 PO 4 6.Ca(OH) 2

5. The mass of 1 mole of the substance Measured in grams Example: 1.0 mole of CO 2 = 44g

6. Use the gram formula mass to convert between moles and mass 1 mole = gram formula mass Examples: 1. 1.5 moles of CO 2 =_______ g 2. 24 g of C =_______ moles 3. 3.0 moles of Br =_______ g 4. 5.0 g of Al =_______ moles

7. If you have an ideal gas at standard temperature and pressure (STP) – 0 o C and 1atm 1 mole = 22.4 L

8. Mole – Volume Examples **All are at STP 1. 1.5 moles of CO 2 (g) = _______ L 2. 38 L of Ne(g) = _______ moles 3. 3.0 moles of O 2 (g) = _______ L 4. 11 L of N 2 (g) = _______ moles

9. 1 dozen = 12 1 mole = 6.02 x 10 23 (Avagadro’s Number)

10. 602,000,000,000,000,000,000,000 A mole of marbles would spread over the surface of the earth, and produce a layer about 50 miles thick A mole of sand, spread over the United States, would produce a layer 3 inches deep A mole of dollars could not be spent at the rate of a billion dollars a day over a trillion years

11. He stated that equal volumes of all gases at the same temperature and pressure contain the same number of molecules (Avogadro’s Principle) Later his work led to the realization that a molecular mass in grams (mole) of any substance contains the same number of molecules (6.02 x 10 23 ) Amedeo Avogadro

12. 1 mole of atoms = 6.02 x 10 23 atoms 1 mole of particles = 6.02 x 10 23 particles 1 mole of molecules = 6.02 x 10 23 molecules 1 mole of compounds = 6.02 x 10 23 compounds

13. Mole – Number Examples 1.4.00 moles of NaCl = _______ molecules 2.0.50 mole of MgBr 2 = _______ molecules 3.3.01 x 10 23 NaCl molecules = _______ moles 4.6.02 x 10 21 Zn atoms = _______ moles

14. Empirical Formula Represents the simplest integer ratio in which the atoms combine to form a compound (the reduced form)

15. Molecular Formula The actual formula May be a multiple of the empirical formula Examples: What are the empirical formulas of the following molecules? a.H 2 O 2 b.C 6 H 12 O 6 c.CCl 4

16. Percent Composition Composition in terms of the percentage of each component present Example: H 2 SO 4 –Step 1: Calculate the formula mass of H 2 SO 4 –Step 2: Find the percent mass of each element

17. Percent Mass Example Calculate the percent composition of oxygen in CO 2 How many grams of Na are in 25.0g of NaCl?

18. Hydrates Hydrate – a compound that contains water Anhydrous – hydrate without the water Water of Hydration –Percentage of water in the crystal –Use percent composition formula

19. Water of Hydration Example Find the water of hydration in CuSO 4  5H 2 0 Step 1 - Find the formula mass Step 2 - Percentage of water

20. Water of Hydration Examples Calculate the water of hydration in Na 2 CO 3  10H 2 O 3.85g of a hydrate is heated and only 2.25g remains. Find the percentage of water in the hydrate.

21. Determining Mass Ratios from Formulas Example: Find the mass ratio for a compound with the empirical formula CH 2.

22. Determining the Molecular Formula from the mass and the empirical formula A compound has a molecular mass of 180amu and an empirical formula of CH 2 O. What is its molecular formula? Step 1: Determine the molecular mass of the empirical formula Step 2: Divide the molecular mass of the compound by the mass of the empirical formula. Step 3: Multiply the subscripts in the empirical formula by your answer to step 2

23. Determining the empirical formula from percent composition What is the empirical formula of a compound that consists of 58.80% barium, 13.75% sulfur, and 27.45% oxygen by mass? Step 1: Assume that the mass of the sample is 100g Step 2: Convert the masses into moles Step 3: Find the smallest whole numbers ratio (divide each number from step 2 by the smallest one)

24. Chemical Reactions Reactant – Substance that enters into a reaction, written to the left of the arrow, starting material Product – substance that is produced by the reaction, written to the right of the arrow, end material Example: HCl + NaOH  NaCl + H 2 O –Reactants: –Products:

25. Reaction that requires energy (heat) in order to occur – heat enters Heat is absorbed Heat is a reactant Surroundings will feel cold because heat has been absorbed from the surroundings Example: H 2 O(s) + heat  H 2 O(l)

26. Reaction that produce energy (heat) when they occur – heat exits Heat is released, given off Heat is a product Surroundings will feel hot because heat was released to the surroundings Example: H 2 O(l)  H 2 O(s) + heat

27. Synthesis Decomposition Single Replacement Double Replacement

28. Direct combination Substances combine to form a new compound Produces 1 product Examples: A + B  AB 2H 2 (g) + O 2 (g)  2H 2 O(l)

29. Break down of a compound into simpler parts Starts with 1 reactant Examples: AB  A + B 2H 2 O(l)  2H 2 (g) + O 2 (g)

30. One substances switches spots with another Element + compound makes a new element plus a new compound Examples: A + BC  B + AC Cu(s) + 2AgNO 3 (aq)  2Ag(s) + Cu(NO3) 2 *copper replaces silver

31. Exchange of ions Everything gets a new "partner" Compound + compound makes new compound + new compound Examples: AB + CD  AD + CB AgNO 3 (aq) + NaCl(aq)  AgCl(s) + NaNO 3 (aq)

32. Balancing Equations Coefficient - a number, placed before formulas to indicate the ratios of moles involved in a reaction Equations must be balanced in accordance with the Law of Conservation of Mass The mass of both sides of the arrow must be equal You must have an equal number of each type of atom on both sides of the equation Example: 2H 2 + O 2  2H 2 O

33. Balancing Examples 1.____ Na + ____ H 2 O  ____ NaOH + ____ H 2 2.____ CaO + ____ H 2 O  ____ Ca(OH) 2 3.____ Al + ____ O 2  ____ Al 2 O 3 4.___ PbCl 2 + ___ Al 2 (SO4) 3  __ PbSO 4 + __ AlCl 3 5.____ Na + ____ O 2  ____ Na 2 O

34. Determining the Missing Mass in Equations If 103.0g of potassium chlorate (KClO 3 ) are decomposed to form 62.7g of potassium chloride (KCl) and oxygen gas (O 2 ) according to the equation 2KClO 3  2KCl + 3O 2, how many grams of oxygen are formed? (Hint: Remember the mass of the products must equal the mass of the reactants)

35. How many grams of Fe are needed to react with 8.0g of O 2 to produce 28.9g of Fe 3 O 4 according to the equation 3Fe + 2O 2  Fe 3 O 4 ?

36. Equation Problems Using the balanced equations and mole conversions you can solve for variety of problems Remember that the coefficients used represent mole ratios

37. Given the following reaction answer the questions below: 2C 2 H 6 + 7O 2  4CO 2 + 6H 2 O 1.How many moles of CO 2 are produced when 2.0 moles of C 2 H 6 reacts? 2.How many moles of H 2 O are produced when 4.0 moles of C 2 H 6 reacts? 3.How many moles of H 2 O are produced when 5.0 moles of C 2 H 6 reacts?

38. Given the following reaction: ____ H 2 + ____ Cl 2  _____ HCl 1.The production of 37g of HCl would require how many moles of H 2 ? 2.If 20.L of H 2 completely reacts how many grams of HCl would be produced?

39. Use the balanced equation to answer the questions below: ____ C 2 H 4 + ____ O 2  ____ CO 2 + ____ H 2 O 1.How many liters of O 2 are used to produce 1.0 mole of H 2 O? 2.How many liters of CO 2 are produced when 9.00 liters of O 2 is consumed? 3.How many liters of C 2 H 4 are needed to produce 10.0 liters of CO 2 ?