4 To Get from Moles to Number of Particles Avogadro’s numberSimply multiply by 6.02 x 1023
5 Practice ProblemsHow many moles of magnesium is 1.25 x 1023 atoms of magnesium?How many atoms are in 2.12 mol of propane? (C3H8)
6 Moles to Mass (grams)Multiply by the molecular mass of the substance.Molar mass: mass (in grams) of one moleof the substance
7 Molar Mass The Mass of 1 mole (in grams) Molecular mass, molecular weight, formula mass, formula weightEqual to the numerical value of the average atomic mass (get from periodic table)1 mole of C atoms = g1 mole of Mg atoms = 24.3 g1 mole of Cu atoms = 63.5 g
8 Learning CheckFind the molar mass (usually we round to the tenths place)A. 1 mole of Br atoms = gB. 1 mole of Sn atoms = g
9 Molar Mass of Molecules and Compounds Mass in grams of 1 mole is numerically equal to the sum of the atomic masses1 mole of CaCl2 = g/mol1 mole Ca x 40.1 g/mol+ 2 moles Cl x 35.5 g/mol = g/mol CaCl2
10 Practice ProblemsHow many grams are in 9.45 mol of dinitrogen trioxide (N2O3)?How many moles in 92.2 g of iron (III) oxide (Fe2O3)?
11 Moles to Volume (Liters) The Special Number22.4
12 Mole – Volume Problemsvolume varies with changes in temperature and pressureusually measured at STP (0oC; kPa)at STP, one mole of any gas occupies a volume of 22.4 L
13 Practice ProblemsDetermine the volume, in liters, of 0.60 mol SO2 gas at STP.The density of a gaseous compound containing carbon and oxygen is g/L at STP. Determine the molar mass of the compound.
14 Number of particles / Avogadro’s number To go from mass, volume or particles to moles, you simply reverse the process.Mass / molecular massVolume / 22.4Number of particles / Avogadro’s number
16 Mixed Mole Practice How many molecules are in 5.0 g of N2O5? What is the volume of 10.0 g of CO2?What is the mass of 1.5 x 1025 molecules of C3H8?
17 10.4 Percent Composition Percent by mass of each element in a compound % mass of element =Grams of element XGrams of compoundX 100
18 Practice Problems9.03 g Mg combine completely with 3.48 g N to form a compound. What is the % composition of this compound?Calculate the % composition of ethane (C2H6).
19 Practice ProblemCalculate the mass of hydrogen in 350 g C3H8.
20 Empirical FormulaLowest whole number ratio of atoms of an element in a compound.May or may not be the same as the molecular formula!If not given an amount, you may assume it is a 100g.
21 Empirical Formula Find moles of each element Divide by the smallest number of molesIf not a whole number, multiply to obtain a whole number.
22 ExampleGive the empirical formula for a compound which is 25.9% Nitrogen and 74.1% Oxygen.Assume it is a 100g sample, so 25.9g N and 74.1g of O25.9/14 =1.85 moles of N74.1/16 = 4.63 moles of O1.85/1.85 = /1.85 = 2.5Double both numbers to get a whole number ratio…N2O5
24 Practice ProblemsCalculate the empirical formula of a compound that is 79.8% C and 20.2% H.Calculate the molecular formula of a compound whose molecular mass is 62 g/mol and empirical formula is CH3O.
25 10.5 Naming HydratesCompounds with specific numbers of water molecules bound to their atoms are called hydrates.In the formula, a dot is used to show that water is bonded. Use a prefix to name the hydrate.Na2CO3 ∙ 10H2OSodium carbonate decahydrate
26 HydratesIn order to analyze a hydrate, you must first find the number of moles of water associated with one mole of the hydrate.Heat the sample to drive off the water, then you can mass the anhydrous compound and determine the moles of water.
27 Practice ProblemA mass of 2.50 g of blue, hydrated copper sulfate (CuSO4· xH2O) is placed in a crucible and heated. After heating, 1.59 g of white anhydrous copper sulfate (CuSO4) remains. What is the formula for the hydrate? Name the hydrate.
28 Practice ProblemAn 11.75g sample of a common hydrate of cobalt (II) chloride is heated. After heating, mol of anhydrous cobalt chloride remains. What is the formula and the name of this hydrate?