Stoichiometry and The Mole 6.02 X 10 23 The word stoichiometry derives from two Greek words:...
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Chapter 10 Stoichiometry and The Mole 6.02 X 10 23
Stoichiometry and The Mole 6.02 X 10 23 The word stoichiometry derives from two Greek words: stoicheion (meaning "element") and metron (meaning "measure")
The word stoichiometry derives from two Greek words: stoicheion
(meaning "element") and metron (meaning "measure"). Stoichiometry
deals with calculations about the masses (sometimes volumes) of
reactants and products involved in a chemical reaction. It is a
very mathematical part of chemistry, so be prepared for lots of
calculator use.
Slide 4
What do the terms ream, gross, dozen, pair have in common? They
are all counting units, designed to make counting objects easier.
Today, we will look at a counting unit for chemistry.
Slide 5
Chemists need a convenient method for counting atoms,
molecules, and formula units in a sample of substance. This
counting unit is called the Mole. What is a Mole???? Its Just like
a Dozen only bigger.
Slide 6
SI base unit used to measure the amount of substance. It is the
number of carbon atoms in exactly 12 grams of Carbon-12. Has a
value of 6.02 x 10 23 Called Avogadros Number Named after Amedeo
Avogadro Commonly abbreviated mol. Just like a dozen only
bigger
Slide 7
Enough soft drink cans to cover the surface of the earth to a
depth of over 200 miles. If you had Avogadro's number of unpopped
popcorn kernels, and spread them across the United States of
America, the country would be covered in popcorn to a depth of over
9 miles. If we were able to count atoms at the rate of 10 million
per second, it would take about 2 billion years to count the atoms
in one mole. 6.02 X 10 23
Slide 8
To appreciate the magnitude of the mole To practice dimensional
analysis
Slide 9
Molecules for covalently bonded substances Ex. A water molecule
Formula Units for Ionic Substances Ex. A formula unit of sodium
chloride Atoms for elements Ex. An atom of sulfur 6.02 X 10 23
Slide 10
1 mole = 6.02 x 10 23 particles Used to convert Moles to
Particles Particles to Moles
Slide 11
How many moles of methanol, CH3OH, are there in 6.53 x 10 23
molecules of methanol?
Slide 12
6.02 x 10 23 particles 1 mole or 1 mole 6.02 x 10 23 particles
Note that a particle could be an atom OR a molecule!
Slide 13
A sample containing 0.75 moles of CO 2 would contain how many
molecules?
Slide 14
How many Fe atoms would be present in 1.27 moles of Fe?
Slide 15
4.47 x 10 23 molecules of C 6 H 12 O 6 would be how many
moles?
Slide 16
Calculate the number of moles contained in 4.50 x 10 24 atoms
of zinc.
Slide 17
How many molecules in 2.8 moles of water?
Slide 18
Draw a bowl with one dozen grapes in it. Draw a bowl with a
dozen oranges in it. Compare the masses. Why are they not the
same?
Slide 19
Draw a circle. Label it as 1 mole of copper atoms. Draw another
circle. Label as 1 mole of aluminum atoms. How many atoms is one
mole? Label. Which weighs more? Look up masses on periodic
table.
Slide 20
The mass of one mole of a substance is called "molar mass"
Units g/mol (grams per mole). Molar mass is the weight in grams of
one mole One mole contains 6.02 x 10 23 entities Therefore, a molar
mass is the mass in grams of 6.022 x 10 23 entities
Slide 21
Aluminum Zinc Copper Iron
Slide 22
Just as a dozen oranges would not weigh the same as a dozen
grapes A mole of copper atoms does not have the same mass as a mole
of aluminum atoms. We know that the relative scale for atomic mass
uses the carbon-12 isotope as a standard.
Slide 23
Each atom of carbon-12 has a mass of 12 amu. The atomic mass on
the periodic table are weighted averages of all isotopes. Since one
mole is defined as the number of carbon-12 atoms in exactly 12
grams of C-12. Therefore
Slide 24
The mass of One Mole of C-12 is 12.0 grams. Called Molar Mass:
the mass in grams of one mole of a pure substance Example: An atom
of Manganese Atomic Mass = 54.94 amu Molar Mass = 54.94 g/mol
Slide 25
1 mole Ag= 6.02 x 10 23 atoms Ag = 107.87 grams Ag 1 mole K =
6.02 x 10 23 atoms K = 39.10 grams K 1 mole H 2 gas = 6.02 x 10 23
molecules H 2 = 2.02 g H 2
Slide 26
Calculate the mass in grams of 0.0450 moles of chromium.
Slide 27
How many moles of calcium are contained in 525 grams of
calcium?
Slide 28
Similar mole-mass, mass-mole conversions can be made for
compounds We must know the Molar Mass for the Compound. How do we
calculate Molar Mass of a Compound?
Slide 29
Recall that for CCl 2 F 2 The subscripts tell us that one
molecule of freon contains One atom carbon Two atoms chlorine 2
atoms fluorine
Slide 30
Suppose you have one mole of freon molecules there would then
be One mole of carbon atoms Two moles chlorine atoms Two moles
fluorine atoms
Slide 31
6 moles of CCl 2 F 2 contains 6 moles carbon atoms 12 moles
chlorine atoms 12 moles fluorine atoms
Slide 32
Mass in grams of 1 mole equal numerically to the sum of the
atomic masses 1 mole of CaCl 2 1 mole Ca x 40.1 g/mol + 2 moles Cl
x 35.5 g/mol = 70.9 g/mol = 111.1 g/mol CaCl 2
Slide 33
Calculate the molar mass of K 2 CrO 4
Slide 34
Calculate the Molar Mass of K 2 O
Slide 35
Prozac, C 17 H 18 F 3 NO, is a widely used antidepressant that
inhibits the uptake of serotonin by the brain. Find its molar
mass.
Slide 36
How many moles are represented by 16.0 g of ethanol, C 2 H 5 OH
?
Slide 37
How many moles of NaCl are in 16.0 grams of NaCl?
Slide 38
How many moles of potassium hydroxide, KOH are in 40.6 g?
Slide 39
How many moles of glucose, C 6 H 12 O 6 are in 27.2 g of
glucose?
Slide 40
How many grams in 0.158 moles of KMnO 4 ?
Slide 41
How many grams in 1.2 moles of H 2 O?
Slide 42
How many grams in 0.87 moles of H 2 O 2 ?
Slide 43
How many grams in 0.43 moles of C 6 H 12 O 6 ?
Slide 44
X molar mass Grams Moles Moles Grams divide by molar mass
Slide 45
The artificial sweetener aspartame (Nutra-Sweet) formula C 14 H
18 N 2 O 5 is used to sweeten diet foods, coffee and soft drinks.
How many moles of aspartame are present in 225 g of aspartame?
Slide 46
6.02 X 10 23 particles = 1 mole AND 1 mole = molar mass (grams)
You can convert atoms/molecules to moles and then moles to grams!
(Two step process) You cant go directly from atoms to grams!!!! You
MUST go thru MOLES. Thats like asking 2 dozen cookies weigh how
many ounces if 1 cookie weighs 4 oz? You have to convert to dozen
first!
Slide 47
How many grams of glucose are in 6.63 x 10 23 molecules of
glucose, C 6 H 12 O 6 ?
Slide 48
Determine the number of molecules found in a 12.4 g sample of H
2 SO 4.
Slide 49
3.14 x 10 23 molecules of CO 2 are produced in a chemical
reaction. How much would the sample weigh in grams?
Slide 50
How many atoms are in a 39.8 g sample of Fe?
Slide 51
Mole Road Map Everything must go through Moles!!!!
Slide 52
A sample of AlCl 3 has a mass of 35.6 grams. How many aluminum
ions are present? How many chloride ions are present?
Slide 53
What is the mass in grams of one formula unit of aluminum
chloride?
Slide 54
How many atoms of Cu are present in 35.4 g of Cu? 35.4 g Cu 1
mol Cu 6.02 X 10 23 atoms Cu 63.5 g Cu 1 mol Cu = 3.4 X 10 23 atoms
Cu
Slide 55
How many atoms of K are present in 78.4 g of K?
Slide 56
What is the mass (in grams) of 1.20 X 10 24 molecules of
glucose (C 6 H 12 O 6 )?
Slide 57
How many atoms of O are present in 78.1 g of oxygen? 78.1 g O 2
1 mol O 2 6.02 X 10 23 molecules O 2 2 atoms O 32.0 g O 2 1 mol O 2
1 molecule O 2
Slide 58
At standard temperature and pressure (STP) a mole of any gas
will occupy 22.4 liters Molar Volume @ STP 1 mole of any gas @STP =
22.4 Liters What is STP? Standard temperature and Pressure 101.3
kpa and 273 K Arbitrarily chosen conditions that we all agree on so
we know the conditions for measuring the volume
Slide 59
Convert 427 Liters of CO 2 to moles.
Slide 60
37 liters of O2 to moles.
Slide 61
How about 3.4 moles of CO to liters?
Slide 62
122 moles of Methane to liters?
Slide 63
How many molecules are in 22.4 liters of methane?
Slide 64
3.58 x 10 23 molecules of propane C 8 H 8 would occupy how much
space at STP?
Slide 65
You collect 14.2 liters of CO gas from an experiment. How many
molecules would be in the sample?
Slide 66
28 grams of H 2 are produce from an experiment. How much volume
would they displace at STP?
Slide 67
Slide 68
Many times a chemist is called upon to determine the makeup of
chemical compound. This is the job of an analytical chemist How???
Calculating Percent Composition.
Slide 69
% by mass = mass of element x 100 mass of compound Suppose a
100 gram sample of compound is made up of 55 grams of X and 45
grams of Y 55 g X / 100 g compound x 100 = 55 % X 45 g Y / 100 g
compound x 100 = 45 % Y Percents by mass of all element s in a
compound must always add to 100 percent
Slide 70
The percent composition of a formula is always the same. You
can assume the sample size is one mole. This allows the use of
molar mass to calculate percent composition.
Slide 71
Slide 72
What is the percent carbon in C 5 H 8 NO 4 (the glutamic acid
used to make MSG monosodium glutamate), a compound used to flavor
foods and tenderize meats?
Slide 73
Suppose we know the elements in a sample of a new compound We
can use this data to determine the formula of the compound.
How?
Slide 74
Formulas give the relative numbers of atoms or moles of each
element in a formula unit - always a whole number ratio (the law of
definite proportions). NO 2 2 atoms of O for every 1 atom of N NO 2
2 atoms of O for every 1 atom of N 1 mole of NO 2 : 2 moles of O
atoms to every 1 mole of N atoms If we know or can determine the
relative number of moles of each element in a compound, we can
determine a formula for the compound.
Slide 75
Empirical Formula The formula of a compound that expresses the
smallest whole number ratio of the atoms present. Ionic formula are
always empirical formula Molecular Formula The formula that states
the actual number of each kind of atom found in one molecule of the
compound.
Slide 76
If the two formulas are different the molecular formula is a
simple multiple of the empirical formula. For example: Hydrogen
Peroxide Empirical Formula is HO. Molecular Formula is H 2 O 2
Slide 77
1.Determine the mass in grams of each element present, if
necessary. 2.Calculate the number of moles of each element.
3.Divide each by the smallest number of moles to obtain the
simplest whole number ratio. 4. If whole numbers are not obtained *
in step 3), multiply through by the smallest number that will give
all whole numbers * Be careful! Do not round off numbers
prematurely
Slide 78
Percent to Mass Mass to Moles Divide by small Multiply til
whole
Slide 79
A sample of a brown gas, a major air pollutant, is found to
contain 2.34 g N and 5.34g O. Determine a formula for this
substance. require mole ratios so convert grams to moles moles of N
= 2.34g of N = 0.167 moles of N 14.01 g/mole 14.01 g/mole moles of
O = 5.34 g = 0.334 moles of O 16.00 g/mole 16.00 g/mole Formula:
Formula:
Slide 80
A substance has the following composition by mass: 60.80 % Na ;
28.60 % B ; 10.60 % H What is the empirical formula of the
substance?
Slide 81
Since two compounds can have the same Empirical Formula we must
determine the actual formula. Called Molecular Formula.
Slide 82
1. Determine the molar mass of the actual compound through an
experiment. 2. Calculate the molar mass of the empirical formula.
3. Divide the actual mass by the mass of the empirical formula. 4.
The result shows how many times bigger the molecular formula is
than the empirical formula. Multiply all subscripts in the
empirical formula by this factor.
Slide 83
A compound has an empirical formula of NO 2. The colorless
liquid, used in rocket engines has a molar mass of 92.0 g/mole.
What is the molecular formula of this substance?
A hydrate can be analyzed by driving off the water with heat.
The remaining substance is called the Anhydrous salt meaning
without water Some hydrates are a color different than their
anhydrous salt Cobalt (II) chloride hexahydrate is pink Without
water it is blue.
Slide 87
MgSO 4 * ?H 2 O 1. Heat the sample to drive off all water. 2.
Mass the anhydrous compound. 3. mass of water = Mass of hydrate
mass of anhydrous salt 4. Convert these masses to moles. 5.
Calculate the mole ratio between the compound and the water
molecules.
Slide 88
Mainly used as dessicants: substances used to keep things
moisture free by absorbing excess water.
