Transcript Training Presentation

Warm Up
• What is a mole?
• What is molar mass?
• What is Avogadro’s number?
Chapter 7
The Mole and Chemical Compostion
Unit Essential Question:
HOW CAN CHEMICAL
COMPOSITION BE DETERMINED?
Lesson Essential Question:
HOW IS THE MOLE USED IN
CONVERSIONS?
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.
covalent
compounds
ionic
compounds
Recall that formula units = simplest ratio of ions
in an ionic compound.
Recall your mole map!
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!
Sample Problems 1 & 2: Moles
& Particles
• Find the number of molecules in 2.5 mol
of sulfur dioxide.
1.5 x 1024 molecules SO2
• A sample contains 3.01 x 1023 molecules
of sulfur dioxide. Determine the amount
in moles.
0.500mol SO2
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.
Sample Problems 3 & 4: Moles
& Mass
• What is the mass of 5.3mol Be?
48g Be
• If you have 27.0g of manganese, how
many moles do you have?
0.491mol Mn
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!
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
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
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
More Practice
• How many moles of iron (III) sulfate,
Fe2(SO4)3, are there in a 178g sample?
0.445mol
Lesson Essential Question:
HOW ARE MOLAR MASSES ON THE
PERIODIC TABLE DETERMINED?
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.
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?
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).
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?
Lesson Essential Questions:
WHAT INFORMATION CAN BE
DETERMINED FROM FORMULAS?
HOW CAN FORMULAS BE
DETERMINED?
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%.
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
Sample Problem #I
• Calculate the percent composition of
copper (I) sulfide.
Sample Problem #2
• Calculate the percent composition of
isopropyl alcohol, (CH3)2CHOH.
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  18.2g
Step 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.
Determining Empirical
Formulas Cont.
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.)
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.
Sample Problem #2
• A compound is found to contain 38.77%
Cl and 61.23% O. What is the empirical
formula?
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 formula
where n is a whole number.
− Ex: 6(CH2O)  C6H12O6
• Molecular and empirical formulas can be the
same!
Molecular Formulas Cont.
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.
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.
Sample Problem #2
• A brown gas has the empirical formula
NO2. Its experimental molar mass is
46g/mol. What is the molecular formula?
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
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.
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?
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.
Sample Problem #2
• What percentage, by mass, of water is found
in the hydrate CuSO4·5H2O?