warm up what is a mole? what is molar mass? what is avogadro’s number?
TRANSCRIPT
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Warm Up
• What is a mole?
• What is molar mass?
• What is Avogadro’s number?
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Chapter 7
The Mole and Chemical
Compostion
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How can chemical composition be determined?
Unit Essential Question:
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How is the mole used in conversions?
Lesson Essential Question:
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Section 1: Avogadro’s Number and Molar Conversions• 1 mole = 6.022 x 1023 particles
−SI unit for amount of substance.•It’s a counting unit (like a dozen).
−Remember that the unit of particles can be: ions, molecules (mcs.), atoms, formula units (f.u.), etc.
Recall that formula units = simplest ratio of ions in an ionic compound.
covalent compounds
ionic compounds
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Recall your mole map!
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Converting moles particles• Same as Chapter 3, but it will involve
molecules, formula units, or ions instead
of just atoms.
• Steps:1) Need 1mol = 6.022 x1023 molecules, etc.2) Use dimensional analysis- turn this into a fraction!*Be sure to place the correct units on the top and bottom so they cancel!
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Sample Problems 1 & 2: Moles & Particles
• Find the number of molecules in 2.5 mol of sulfur dioxide.
• A sample contains 3.01 x 1023 molecules of sulfur dioxide. Determine the amount in moles.
1.5 x 1024 molecules SO2
0.500mol SO2
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Molar Mass• Amount of mass (in grams) in 1 mole of a
substance.
• Use molar masses from the periodic table.−Round to 2 decimal places!−Use units of g/mol.−Example:
• C: 12.01g/mol means that 1 mol C = 12.01 g• Cl: 35.45g/mol means that 1 mol Cl = 35.45g
• Use to convert between moles and mass.
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Sample Problems 3 & 4: Moles & Mass
• What is the mass of 5.3mol Be?
• If you have 27.0g of manganese, how many moles do you have?
48g Be
0.491mol Mn
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Molar Masses of Compounds• Add together the molar masses of all
elements or ions present.−Ex: CH4
−C: 12.01g/mol H: 1.01g/mol−12.01g/mol + 4(1.01g/mol) = 16.05g/mol
−This means that 1 mole of CH4 has a mass of 16.05g.
• You will need to calculate the molar mass
of a compound whenever you are
converting between mass and moles!
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Additional Molar Mass Examples:• Element
−Ag = 107.87 g/mol
• Diatomic Element/molecule−Br2 = 79.90 x 2 = 159.80 g/mol
• Molecule (Covalent compound)−H2O = (1.01 x 2) + 16.00 = 18.02 g/mol
• Formula unit (Ionic compound)−Ca(NO3)2 = 40.08 + (2 x 14.01) + (6 x
16.00) = 164.10 g/mol
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Sample Problem 5: Mass to Moles with a Compound
• Find the number of moles present in 47.5 g of glycerol, C3H8O3.
• Hint: you will need to calculate the molar mass of glycerol!
Glycerol’s molar mass: 92.11g/mol
0.516mol C3H8O3
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Sample Problem 6: Number of Particles to Mass
• Remember- you can’t go directly between mass (g) and the number of particles! You must convert to moles first!
• Find the mass in grams of 2.44 x 1024 atoms of carbon.
48.7 g C
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More Practice
• How many moles of iron (III) sulfate,
Fe2(SO4)3, are there in a 178g
sample?
0.445mol
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How are molar masses on the periodic table
determined?
Lesson Essential Question:
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Mole Ratios in Chemical Formulas• Ratios can be formed between amounts of elements or ions within
a compound.−Look at the subscripts.
• Example #1: CaCl2
−For every 1mol of CaCl2 there is 1mol of Ca+2 ions and 2mol of Cl- ions.
• Example #2: Na2CO3
−For every 1mol of Na2CO3, there are 2mol of Na+ ions and 1mol of CO3
-2 ions.
• Example #3: N2O3
−For every 1mol of N2O3 there are 2mol of N atoms and 3mol of O atoms.
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Practice
• If you have one mole of strontium
cyanide, Sr(CN)2, how many moles of
strontium ions are there? How many
moles of cyanide ions are there?
• Given the compound P2O5 what is the
mole ratio of P atoms to O atoms?
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Section 2: Relative Atomic Mass and Chemical
Formulas• Periodic table masses are averages of
all isotopes present. −Recall that we said a weighted average is
used- takes into account the amount of each isotope.
−Average atomic mass: (% x atomic mass)+(% x atomic mass)+…
100−Note: % is the percent abundance (how
often the element is found as that isotope in nature).
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Sample Problem• The mass of a Cu-63
atom is 62.94 amu, and that of a Cu-65 atom is 64.93 amu. If the abundance of Cu-63 is 69.17% and the abundance of Cu-65 is 30.83%, what is the average atomic mass of copper?
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What information can be determined from formulas?
How can formulas be determined?
Lesson Essential Questions:
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Calculating Percent Composition
• Tells you the percent each element
makes up of the whole compound.Step 1: Determine the molar mass of the entire compound.Step 2: Divide each element’s total molar mass by the molar mass of the compound.Step 3: Multiply by 100 to get percent.Step 4: Check your answer by adding up the percentages to makes sure they equal 100%.
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Percent Composition Cont.
• Calculating the percent composition of
a compound can be helpful in
determining the formula/identity.
• Example:−Iron and oxygen form two compounds:
•Fe2O3 and FeO
−Fe2O3 = 69.9% Fe and 30.1% O
−FeO = 77.7% Fe and 22.3% O
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Sample Problem #I• Calculate the percent composition of
copper (I) sulfide.
• Calculate the percent composition of isopropyl alcohol, (CH3)2CHOH.
Sample Problem #2
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Determining Empirical Formulas
• The empirical formula shows the simplest ratio of elements/ions in the compound.−Ionic compounds are represented with empirical
formulas.
• Given percent composition data, you can determine the empirical formula of a compound.Step 1: Assume 100 g of the sample- put ‘g ’ in
for ‘% ’. Ex: 18.2% O 18.2gStep 2: Convert grams to moles.Step 3: Divide each mole value by the smallest
mole value. This will tell you the number of each element that appears in the formula.
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Step 4: If you get a decimal, multiply ALL numbers by a whole number to turn the decimal into a whole number.
•The numbers you will need to multiply by should be relatively small (2, 3, etc.)
Determining Empirical Formulas Cont.
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Sample Problem #1• Chemical analysis of a liquid shows
that it is 60.0% C, 13.4% H, and 26.6% O by mass. Calculate the empirical formula of this substance.
• A compound is found to contain 38.77% Cl and 61.23% O. What is the empirical formula?
Sample Problem #2
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Molecular Formulas• Show the actual numbers of elements in the
compound- not necessarily the simplest
formula.−Often seen for covalent compounds.
• They will be a whole number multiple of the
empirical formula (can’t be a decimal).−In other words: n(empirical formula) = molecular formulawhere n is a whole number.
−Ex: 6(CH2O) C6H12O6
• Molecular and empirical formulas can be the
same!
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Molecular Formulas Cont.
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Molecular Formulas Cont.• The molecular formula can be determined from
the empirical formula and experimental molar
mass of a compound.Step 1: Determine the molar mass of the given empirical formula.Step 2: Solve for n by dividing the experimental molar mass by the molar mass of the empirical formula.*Remember: n(empirical formula) = molecular formula
Step 3: Multiply the subscripts in the empirical formula by n.
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Sample Problem #1• The empirical formula for a compound is
P2O5. Its experimental molar mass is 284g/mol. Determine the molecular formula of the compound.
• A brown gas has the empirical formula NO2. Its experimental molar mass is 46g/mol. What is the molecular formula?
Sample Problem #2
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Hydrates- Honors Only
• Not in the textbook.
• Hydrates – ionic compounds that
contain water molecules within the
crystal structure.−Example: CuSO4•5H2O
−Anhydrous – without the water = CuSO4
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Determining Hydrate Formulas
• Formula can be determined if given: the
mass of the hydrate, the anhydrous
mass, and the formula of the ionic
compound.Step 1: Determine the mass of water in the hydrate (subtract anhydrous mass).Step 2: Convert the anhydrous ionic compound mass and water mass to moles.Step 3: Divide both molar amounts by the smallest number. This gives you the number of water molecules in the hydrate.
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Sample Problem #1
• A 5.82 g sample of Mg(NO3)2· XH2O in an evaporating dish is heated until it is dry. The mass of the anhydrous sample is 2.63 g Mg(NO3)2. What is the formula for the hydrate?
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Determining % Water in a Hydrate
• Formula can be determined if given
the formula of the hydrate.Step 1: Calculate the mass of the entire hydrate and the mass of just the water. Step 2: Divide the mass of the water by the mass of the entire hydrate and multiply by 100 to get a percent.
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• What percentage, by mass, of water is found in the hydrate CuSO4·5H2O?
Sample Problem #2