Mole Concept slides

Chemistry Note

Page 1

  • Measurement in chemistry

    • Mass, volume, and counting pieces

  • Units of measurement

    • Mass is measured in grams

    • Volume is measured in liters

    • Counting pieces is measured in moles

Page 2

  • Definition of moles

    • Number of carbon atoms in 12 grams of carbon-12

    • 1 mole = 6.02 x 10^23 particles

    • Avogadro's number = 6.02 x 10^23

Page 3

  • Representative particles

    • Smallest pieces of a substance

    • Molecular compound: molecule

    • Ionic compound: formula unit

    • Element: atom

Page 4

  • Definition of a mole

    • Amount of a substance containing Avogadro's number of atoms, ions, molecules, or any other chemical unit

    • 1 mole of C-12 atoms = 12.00 g

Page 5

  • Mole calculations

    • 1 mole = 6.02 x 10^23 objects (atoms, molecules)

    • Examples:

      • 1 mol 12C atoms = 6.02 x 10^23 12C atoms

      • 1 mol H2O molecules = 6.02 x 10^23 molecules

      • 1 mole NO3- ions = 6.02 x 10^23 NO3- ions

Page 6

  • The mole as a practical unit

    • Number of atoms in 12.000 grams of 12C

    • 1 mol = grams / formula weight

    • Moles provide a practical unit for measuring atoms, ions, and molecules

Page 7

  • Weighing moles

    • Moles cannot be directly weighed

    • Grams are used to measure moles

    • Different number of grams needed to have the same number of molecules

Page 8

  • Moles and masses

    • Atoms have different sizes and masses

    • Mole of atoms of one type has a different mass than a mole of atoms of another type

    • Examples of atomic masses: H - 1.008 u, O - 16.00 u, Mo - 95.94 u, Pb - 207.2 u

Page 9

  • Masses of atoms and molecules

    • Atomic mass: average, relative mass of an atom in an element

    • Atomic mass unit (u): arbitrary mass unit for atoms

    • Molecular or formula mass: total mass for all atoms in a compound

Page 10

  • Molar masses

    • Mass of one unit: use u

    • Mass of one mole of units: use grams/mole

    • Numbers don't change, only the units

Page 11

  • Molar mass

    • Atomic mass of a substance expressed in grams corresponds to 1 mol of the substance

    • Molar mass of a diatomic substance is twice its atomic mass

Page 12

  • Generic term for the mass of one mole

    • Same as gram molecular mass, gram formula mass, and gram atomic mass

Page 13

  • Gram Atomic Mass

    • Mass of 1 mole of an element in grams

    • Example: 12.01 g C = 1 mole

Page 14

  • Mole calculations example

    • Calculate the number of sodium atoms in 0.120 mol Na

    • Answer: 7.22 x 10^22 atoms Na

Page 15

  • Mole calculations example

    • Calculate the number of moles of potassium in 1.25 x 10^21 atoms of K

    • Answer: 2.08 x 10^-3 mol K

Page 16

  • Mole calculations example

    • Calculate the mass in grams of 2.01 x 10^22 atoms of sulfur

    • Answer: 1.07 g S

Page 17

  • Mole calculations example

    • Calculate the number of O2 molecules in 0.470 g of oxygen gas

    • Answer: 8.84 x 10^21 molecules O2

Page 18

  • Mole calculations for compounds

    • In 1 mole of H2O molecules, there are two moles of H atoms and 1 mole of O atoms

    • To find the mass of one mole of a compound, determine the moles of the elements and add them up

Page 19

  • Molar Mass

    • Laboratory-sized sample: 1 molecule H2O = 1 mol H2O (18.0 amu) (18.0 g)

Page 20

  • Mole calculations for compounds

    • Calculate the number of Magnesium and Chlorine ions present in 0.450 mol of MgCl2

    • Answer: # Mg2+ ions = 2.71 x 10^23, # Cl- ions = 5.42 x 10^23

Page 21

  • The mole concept applied to compounds

    • Formula weight: sum of atomic masses of the atoms in a species

    • Example: Molecular weight of NH3, formula weight of MgF2

Page 22

  • Mass of one mole, one formula unit, and Avogadro's number of formula units

Page 23

  • Gram Formula Mass

    • Mass of one mole of an ionic compound

    • Example: GFM of Fe2O3 = 159.70 g

Page 24

  • Molar Mass Table

    • Mole relationships, formula weight, molar mass, and particles in one mole

Page 25

  • Molar mass of compounds

    • Example: Mass of one mole of CH4 = 16.05 g

Chemistry Note

Page 26

  • Gram atomic weight: mass in grams of one mole of an element

    • Numerically equal to its atomic weight

  • Gram formula weight: mass in grams of one mole of a compound

    • Numerically equal to its formula weight

  • Gram formula mass: sum total of all individual atomic weights in the formula

  • Gram molecular weight: gram formula weight of a molecular compound

Page 27

  • Molar mass of Ag (silver) = atomic mass = 107.87 g/mol

  • Molar mass of magnesium nitrate (Mg(NO3)2) = 24.31 + 2(14.01 + 3 x 16.00) = 148.33 g/mol

Page 28

  • Example: Calculate the number of moles in 5.69 g of NaOH

    • 1 mole Na = 22.99 g

    • 1 mole O = 16.00 g

    • 1 mole H = 1.01 g

    • 1 mole NaOH = 40.00 g

Page 29

  • Example: Calculate the mass in grams of a single molecule of carbon dioxide (CO2)

    • 44.01 g CO2 x 1 mol CO2 = 7.31 x 10^-23 g/molecule

    • 1 mol CO2 = 6.02 x 10^23 molecules

Page 30

  • Molar mass: number of grams of 1 mole of atoms, ions, or molecules

  • Conversion factors can be made from molar mass to convert grams to moles of a compound

Page 31

  • Examples:

    • Calculate the number of moles in 4.56 g of CO2

    • Calculate the mass in grams of 9.87 moles of H2O

    • Calculate the number of molecules in 6.8 g of CH4

    • Calculate the weight of 49 molecules of C6H12O6

Page 32

  • Calculate the molar mass of Ca(NO3)2

  • Calculate the molar mass of a compound if 0.372 moles of it has a mass of 152 g

Page 33

  • Calculate the grams required to have 0.1 mol of:

    • A. NaOH

    • B. H2SO4

    • C. C2H5OH

    • D. Ca3(PO4)2

Page 34

  • Calculate the number of moles in 50.0 g of:

    • A. CS2

    • B. Al2(CO3)3

    • C. Sr(OH)2

    • D. LiNO3

Page 35

  • Calculate the number of C, H, and O atoms in 1.50 g of glucose (C6H12O6)

  • Calculate the average mass of one C3H8 molecule

  • Calculate the mass of 5.00 x 10^24 molecules of NH3

Page 36

  • Types of questions:

    • Calculate the number of molecules of CO2 in 4.56 moles of CO2

    • Calculate the number of moles of water in 5.87 x 10^22 molecules

    • Calculate the number of atoms of carbon in 1.23 moles of C6H12O6

    • Calculate the number of moles in 7.78 x 10^24 formula units of MgCl2

Page 37

  • Gases are difficult to weigh, so we need to know the number of moles of gas

  • Two factors that affect the volume of a gas: temperature and pressure

  • Comparison should be made at the same temperature and pressure

Page 38

  • Standard Temperature and Pressure (STP): 0ºC and 1 atm pressure

  • At STP, 1 mole of gas occupies 22.7 L (molar volume)

  • Avogadro's Hypothesis: equal volumes of gas at the same temperature and pressure have the same number of particles

Page 39

  • Molar Volume at STP = 22.7 L

  • Avogadro's Theory: two gases containing equal numbers of molecules occupy equal volumes under similar conditions

  • STP: 0°C and 1 atm

Page 40

  • Examples:

    • Calculate the volume of 4.59 moles of CO2 gas at STP

    • Calculate the number of moles in 5.67 L of O2 gas at STP

    • Calculate the volume of 8.8 g of CH4 gas at STP

Page 41

  • Mass, Volume, PT (Periodic Table), Moles, 6.02 x 10^23 Representative Particles, Ions, Atoms

Page 42

  • Example: Calculate the mass of 3.36 L of ozone gas (O3) at STP

    • 3.36 L O3 x 1 mol O3 x 48.00 g O3 = 7.20 g O3

    • 22.7 L O3 x 1 mol O3

Page 43

  • Example: Calculate the number of molecules of hydrogen gas (H2) occupying 0.500 L at STP

    • 0.500 L H2 x 1 mol x 6.02 x 10^23 molecules H2 = 1.34 x 10^22 molecules H2

    • 22.7 L x 1 mol

Page 44

  • Mole Calculations:

    • (a) Use N as a unit factor: multiply by 1 mol/6.02 x 10^23

    • (b) Use N as a unit factor: multiply by 6.02 x 10^23/1 mol

    • (c) Use molar mass as a unit factor: multiply by 1 mol/g

    • (d) Use molar mass as a unit factor: multiply by g/1 mol

    • (e) Use molar volume as a unit factor: multiply by 1 mol/22.4 L

    • (f) Use molar volume as a unit factor: multiply by 22.4 L/1 mol