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Ch. 10 Chemical Quantities. 3 Methods of Measuring. Counting Mass Volume. Example 1. If 0.20 bushel is 1 dozen apples, and a dozen apples has a mass of 2.0 kg, what is the mass of .050 bushel of apples?. Example 1. Count: 1 dozen apples = 12 apples Mass: 1 dozen apples = 2.0 kg apples - PowerPoint PPT Presentation
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Ch. 10 Chemical Quantities
3 Methods of Measuring
• Counting• Mass• Volume
Example 1
• If 0.20 bushel is 1 dozen apples, and a dozen apples has a mass of 2.0 kg, what is the mass of .050 bushel of apples?
Example 1
• Count: 1 dozen apples = 12 apples• Mass: 1 dozen apples = 2.0 kg apples• Volume: 1 dozen apples = 0.20 bushels applesConversion Factors:• 1 dozen 2.0 k.g 1 dozen
12 apples 1 dozen 0.20 bushels
Example 1
• 0.50 bushel x 1 dozen x 2.0 kg = 0.20 bushel 1 dozen
= 5.0 kg
Avogadro’s Number
• Named after the Italian scientist Amedo Avogadro di Quaregna
• 6.02 x 10 23
Mole (mol)
• 1 mol = 6.02 x 10 23 representative particles
• Representative particles: atoms, molecules ions, or formula units (ionic compound)
Mole (mol)
• Moles= representative x 1 mol particles 6.02 x 10 23
Example 2 (atoms mol)
• How many moles is 2.80 x 10 24 atoms of silicon?
Example 2
• 2.80 x 10 24 atoms Si x 1 mol Si 6.02 x 10 23 atoms Si
= 4.65 mol Si
Example 3 (mol molecule)
• How many molecules of water is 0.360 moles?
Example 3
• 0.360 mol H2O x 6.02 x 10 23 molecules H2O1 mol H2O
=2.17 molecules H2O
The Mass of a Mole of an Element• The atomic mass of an element expressed
in grams = 1 mol of that element = molar mass
Molar mass S
Molar mass C
Molar mass Hg
Molar mass Fe
6.02 x 10 23 atoms S
6.02 x 10 23 atoms C
6.02 x 10 23 atoms Hg
6.02 x 10 23 atoms Fe
Example 4 (mol gram)
• If you have 4.5 mols of sodium, how much does it weigh?
Example 4
• .45 mol Na x 23 g Na = 10.35 g Na = 1.0 x 10 2 g Na 1 mol Na
Example 5 (grams atoms)
• If you have 34.3 g of Iron, how many atoms are present?
Example 5
• 34.3 g Fe x 1 mol Fe x 6.02 x 10 23 atoms 55.8 g Fe 1 mol Fe=3.70 x 10 23 atoms Fe
The Mass of a Mole of a Compound
• To find the mass of a mole of a compound you must know the formula of the compound
• H2O H= 1 g x 2
O= 16 g 18 g = 1 mole = 6.02 x 10 23
molecules
Example 6 (gram mol)
• What is the mass of 1 mole of sodium hydrogen carbonate?
Example 6
• Sodium Hydrogen Carbonate = NaHCO3
• Na=23 g• H=1 g• C=12 g• O=16 g x3• 84 g NaHCO3 = 1 mol NaHCO3
Mole-Volume Relationship
• Unlike liquids and solids the volumes of moles of gases at the same temperature and pressure will be identical
Avogadro’s Hypothesis
• States that equal volumes of gases at the same temperature and pressure contain the same number of particles
• Even though the particles of different gases are not the same size, since the gas particles are spread out so far the size difference is negligible
Standard Temperature and Pressure (STP)
• Volume of a gas changes depending on temperature and pressure
• STP= 0oC (273 K) 101.3 kPa (1 atm)
Standard Temperature and Pressure (STP)
• At STP, 1 mol = 6.02 X 1023 particles = 22.4 L of ANY gas= molar volume
Conversion Factors
• AT STP• 1 mol gas 22.4 L gas 22.4 L gas 1 mol gas
Example 7
• At STP, what volume does 1.25 mol He occupy?
Example 7
• 1.25 mol He x 22.4 L He = 28.0 L He 1 mol He
Example 8
• If a tank contains 100. L of O2 gas, how many moles are present?
Example 8
• 100. L O2 X 1 mol O2 = 4.46 mol O2
22.4 L O2
Calculating Molar Mass from Density
• The density of a gas at STP is measured in g/L
• This value can be sued to determine the molar mass of gas present
Example 9
• A gaseous compound of sulfur and oxygen has a density of 3.58 g/L at STP. Calculate the molar mass.
Example 9
• 1 mol gas x 22.4 L gas X 3.58 g gas = 1 mol gas 1 L gas
Molar Mass= 80.2 g
Percent Composition
• The relative amounts of the elements in a compound
• These percentages must equal 100
Percent Composition
• %element = mass of element x 100 mass of compound
Example 10
• Find the percentage of each element present in Al2 (CO3)3
Example 10
• Al2(CO3)3
• Al= 27 g x 2 = 54 g / 234 g x 100=23%• C= 12 g x 3 = 36 g/ 234 g x 100= 15%• O = 16 g x 9 = 144 g / 234 g x 100=62% 234 g Al2(CO3)3