3Mole Facts6.02 X 1023 Pennies: Would make at least 7 stacks that would reach the moon.6.02 X 1023 Watermelon Seeds: Would be found inside a melon slightly larger than the moon.6.02 X 1023 Blood Cells: Would be more than the total number of blood cells found in every human on earth.1 Liter bottle of Water contains 55.5 moles H20
4Definition of MoleThe amount of atoms in 12.0 grams of Carbon 12 (6.02 x 1023 atoms known as Avogadro’s number).A sample of any element with a mass equal to that element's atomic weight (in grams) will contain precisely one mole of atoms (6.02 x atoms).
5The sum of the masses of all the atoms in a molecule of a substance Molecular MassThe sum of the masses of all the atoms in a molecule of a substanceCaCO31 atom of Ca = amu1 atom of C = amu3 atoms of O = amuamu}Add these
6Formula Mass (Molar Mass) The mass of 1 mole (in grams)Equal to average atomic mass but the unit is grams1 mole of C atoms = g1 mole of Na atoms = g1 mole of Cu atoms = g
7Example Problem Find the mass of 1 mole of KAl(SO4)2 ● 12H2O 2 (SO4) 2( ((16.00 x 4))=192.1212 H2O 12( ) =216.24Mass of 1 mole = g/mol
8Try These: Find the molecular mass for these : HNO3CO2Find the molar mass for these compounds:C6H10O5H2SO4
9The Mole1 mole of gas always contains 6.02 x molecules of that gas1 mole Cl2 gas = 6.02 x 1023 molecules of Cl21 mole NO2 gas = 6.02 x 1023molecules of NO21 mole CO gas = 6.02 x 1023 molecules of CO1 mole CO2 gas = 6.02 x 1023molecules of CO2
10The MoleAlso applies to other particles! (not only molecules in a gas)1 mole C = x 1023 C atoms1 mole H2O = x 1023 H2O molecules1 mole NaCl = x 1023 NaCl formula units1 mole of Na+ = 6.02 x 1023 Na+ ions1 mole of Cl- = x 1023 Cl– ions
11Avogadro’s Number We can use Avogadro’s # as a conversion factor: 1 mole6.02 x 1023 particlesOrNote that a particle could be an atom, molecule, formula unit, or ion !
12Example Problems How many molecules are in 3.5 moles of H2O? How many moles are present in molecules of NO2?
13Mass and Mole Relationship 1 mole of any substance = the molar mass of that substance (in grams)Find the number of moles present in 56.7 g of HNO3.56.7 g HNO mole HNO363.01 g HNO3
14Example ProblemsFind the number of grams present in 4.5 moles of C6H10O5.Find the number of moles present in g of H2SO4.How many molecules are in 4.5 grams of NaCl?
15Gas Volumes and Molar Mass Avogadro’s LawEqual volumes of gases under the same conditions of temperature and pressure contain equal numbers of molecules1 mole of any gas at standard temperature and pressure (STP) occupies 22.4 litersStandard temperature: 0ºC or 273KStandard pressure: 1 atm or kPa
16Gas Volumes and Molar Mass 32.00 g O2 = 1 mole = 22.4 L2.02 g H2 = 1 mole = 22.4 L44.01 g CO2 = 1 mole = 22.4 L
17Example Problems How many liters are present in 5.9 moles of O2? How many liters are present in 3.67 moles of CO2?How many atoms of O are present in 78.1 g of O2?
18Total mass of the element in the compound Percent CompositionFinding what percent of the total weight of a compound is made up of a particular elementFormula for calculating % composition:Total mass of the element in the compoundTotal formula massX 100
19Example Problem:Calculate the % composition of calcium in Ca(OH)2.
20Example ProblemsFind the percentage composition of a compound that contains 1.45 g of carbon, g of sulfur, and 1.00 g of hydrogen in a g sample.21.7% carbon63.3% sulfur15.0% hydrogen
21Example ProblemA sample of an unknown compound with a mass of 5.00 grams is made up of 75% carbon and 25% hydrogen. What is the mass of each element?3.75 g of carbon1.25 g of hydrogen
22FormulasEmpirical Formula - expresses the smallest whole number ratio of atoms presentE.g. CH2OIonic formulas are always empirical formulasMolecular Formula - states the actual number of each kind of atom found in one molecule of the compound.E.g. C6H12O6
23Empirical Formula Determine mass in grams of each element Calculate the number of moles of eachDivide each by the smallest number of moles to obtain the simplest whole number ratioIf whole numbers are not obtained in step 3, multiply all by the smallest number that will give whole numbers
24Empirical Formula Remember this: Percent to mass Mass to mole Divide by smallMultiply ‘till whole
25Example ProblemGiven that a compound is composed of 60.0% Mg and 40.0% O, find the empirical formula.
26Example ProblemA compound is found to contain 68.5% carbon, 8.63% hydrogen, and 22.8% oxygen. The molecular weight of this compound is known to be approximately g/mol. Find the empirical and molecular formulas.
27Hydrates Ionic compounds Water is bonded to the crystal structure Ex: CuSO3 • 7H2OThe percentage of water in a hydrate can easily be calculated using the formula:% Water = Mass of water x 100Mass of hydrate
28Example Problem What is the percentage of water in CuSO3 • 7H2O? A 3.5 g sample of a hydrate is heated and only 1.7 g of the anhydrous salt remain. What is the percentage of water?
29Law of Definite Proportions Formulas give the numbers of atoms or moles of each elementAlways a whole number ratio1 molecule NO2 : 2 atoms of O for every 1 atom of N1 mole of NO2 : 2 moles of O atoms to every 1 mole of N atoms
30Law of Multiple Proportions When any two elements, A and B, combine to form more than one compound, the different masses of B that unite with a fixed mass of A have a small whole-number ratioExample:In H2O, the proportion of H:O = 2:16 or 1:8In H2O2, H:O is 2:32 or 1:16
31How Do We Determine Concentration? MolarityMolality
32How do we make solutions? M1 = m1/V rearrange to M1V1 = m1M2 = m2/V rearrange to M2V2 = m2If m1=m2then, M1V1 = M2V2
33M1V1 = M2V2M1 = concentration of the first solution V1 = volume of the first solution M2 = concentration of the second solution V2 = volume of the second solutionLet's consider a sample problem:You have 1 L of a M aqueous solution of table sugar. You want to dilute the solution to M. What do you do?
34DilutionTo solve the problem, you simply plug in the three numbers you know:(0.125 M) (1 L) = (0.05 M) V22.5 L = V2Using the equation, you determine that the volume of the diluted solution should be 2.5 L. So we simply add enough water to the first solution so that the solution's volume becomes 2.5 L.
35What is Saturation?A solution is saturated if it contains as much solute as can possibly be dissolved under the existing conditions of temperature and pressureUnsaturated: Has less than maximum amount of solute that can be dissolvedSupersaturated: Contains more than maximum (How can this happen?)