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Unit 1: Stoichiometry
Chemistry 2202
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Stoichiometry

Stoichiometry deals with quantities
used in OR produced by a chemical
reaction
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3 Parts



Mole Calculations (Chp. 2 & 3)
Stoichiometry and Chemical Equations
(Chp. 4)
Solution Stoichiometry (Chp. 7/8)
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PART 1 - Mole Calculations



Isotopes and Atomic Mass (pp. 43 - 46)
Avogadro’s number (pp. 47 – 49)
Mole Conversions (pp. 50 - 74)

M, MV, NA, n, m, v, N
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PART 1 - Mole Calculations

Percent composition:
- given mass (p. 79 - 82)
 - given the chemical formula (p. 83 - 86)




Empirical Formulas (pp. 87 - 94)
Molecular Formulas (pp. 95 - 98)
Lab: Formula of a Hydrate
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Chapter 2
Isotopes and Atomic Mass
atomic number - the number of protons
in an atom or ion
mass number - the sum of the protons
and neutrons in an atom
isotope - atoms which have the same
number of protons and electrons but
different numbers of neutrons
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Isotopes and Atomic Mass
eg.
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Isotope Name
Carbon - 12
Isotope Notation
12
6
mass number
Isotope Name
Carbon - 14
atomic number
Isotope Notation
14
6
mass number
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C
C
atomic number
Isotopes and Atomic Mass
35
17
Cl
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37
17
Cl
Isotopes and Atomic Mass
24
12
Mg
79 %
25
12
Mg
10 %
26
12
11 %
not all isotopes are created equal
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Mg
Isotopes and Atomic Mass
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Isotopes and Atomic Mass
atomic mass unit (AMU - p.43)
- a unit used to describe the mass of
individual atoms
- the symbol for the AMU is u
- 1 u is 1/12 of the mass of a carbon-12
atom
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Isotopes and Atomic Mass
average atomic mass (AAM)
-
the AAM is the weighted average of all
the isotopes of an element (p. 45)
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Finding % Abundance
Example: Copper, a metal known since
ancient times, is used in electrical cables
and pennies. The atomic masses of its two
stable isotopes, copper- 63 (69.09%) and
copper- 65 (30.91%), are 62.93 amu and
64.9278 amu, respectively. Calculate the
average atomic mass of copper.
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Finding % Abundance
ex. Br has two naturally occurring
isotopes. Br-79 has a mass of 78.92 u
and Br-81 has a mass of 80.92 u. If the
AAM of Br is 79.90 u, determine the
percentage abundance of each isotope.
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Finding % Abundance
Let x = fraction of Br-79
Let y = fraction of Br-81
y=1-x
x+y=1
78.92x + 80.92y = 79.90
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x+y=1
78.92x + 80.92y = 79.90
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Avogadro’s Number (p. 47)
p. 48
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Avogadro’s Number



The MOLE is a number used by chemists
to count atoms
The MOLE is the number of atoms
contained in exactly 12 g of carbon-12.
In honor of Amedeo Avogadro, the
number of particles in 1 mol has been
called Avogadro’s number.
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How big is Avogadro's number?


An Avogadro's number of soft drink cans
would cover the surface of the earth to a
depth of over 200 miles.
Avogadro's number of unpopped popcorn
kernels spread across the USA, would
cover the entire country to a depth of
over 9 miles.
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How big is Avogadro's number?

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.
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Avogadro’s Number
1 mole = 6.02214199 x 1023 particles
1 mol = 6.022 x 1023 particles
NA = 6.022 x 1023 particles/mol
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Avogadro’s Number
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Avogadro’s Number
Number of moles
Number of atoms
5 mol
3.011 x 1024 atoms
0.01 mol
6.022 x 1021 atoms
7.72 mol
4.65 x 1024 atoms
0.0133 mol
8.01 x 1021 atoms
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Avogadro’s Number
Formulas:
N
n
NA
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N = n x NA
n = # of moles
N = # of particles
(atoms, ions, molecules,
or formula units)
NA = Avogadro’s #
Avogadro’s Number
a)
How many moles are contained in the
following?
2.56 x 1028 Pb atoms
b)
7.19 x 1021 CO2 molecules

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Avogadro’s Number
eg. Calculate the number of moles in
4.98 x 1025 atoms of Al.
eg. How many formula units of Na2SO4
are in 5.69 mol of Na2SO4?
# of Na ions?
# of Oxygen atoms?
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Avogadro’s Number
1. How many molecules of glucose are in
0.435 mol of C6H12O6?
How many carbon atoms?
2. Calculate the number of moles in a
sample of glucose that has 3.56 x 1022
hydrogen atoms.
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Molar Mass
The mass of one mole of a substance is
called the molar mass of the substance
eg.
1 mole of Pb has a mass of 207.19 g
1 mole of Ag has a mass of 107.87 g

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Molar Mass

The symbol for molar mass is M and the
unit is g/mol
MPb = 207.19 g/mol
MAg = 107.87 g/mol
eg.
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Molar Mass
The molar mass of a compound is the
sum of the molar masses of the
elements in the compound
eg. Calculate the molar mass of:
a) H2O
b) C6H12O6 c) Ca(OH)2

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Molar Mass
H2O
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Molar Mass
C6H12O6
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Molar Mass
Ca(OH)2
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Your calculator
may not show
the zeroes.
There should be 2
digits after the
decimal when
adding molar
masses
Molar Mass Calculations
Avogadro’s #
N
n
NA
N = n x NA
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m
n
M
mass
molar
mass
m=nxM
Molar Mass Calculations
N
n
NA
m
n
M
N = n x NA
m=nxM
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Molar Mass Calculations
eg. How many moles
are in 25.3 g of NO2?
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Molar Mass Calculations
eg. What is the mass of 4.69 mol of water?
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Particle–Mole-Mass Conversions
particles
(N)
Moles
(n)
Mass
(m)
5.98 x 1026 Cu atoms
4.50 g H2O
6.15 mol O3
Molar Mass Calculations
particles
(N)
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x NA
÷ NA
moles
(n)
xM
÷M
mass
(m)
Molar Mass Calculations
eg. How many molecules are in 26.9 g of
water?
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Molar Mass Calculations
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Molar Mass Calculations
eg. How many molecules are in 4.78 g of
glucose?
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Molar Mass Calculations
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Molar Mass Calculations
eg. A sample of Sn contains 4.69 x 1028
atoms. Calculate its mass.
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Molar Mass Calculations
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Molar Volume
•The volume of a gas increases when
temperature increases but decreases
when pressure increases .
•The volume of gases is measured
under conditions of Standard
Temperature and Pressure (STP)
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Molar Volume
Standard Pressure – 101.3 kPa
 Standard Temperature – 0 °C
 Avogadro hypothesized that equal
volumes of gases at the same
temperature and pressure contain
equal numbers of molecules.

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Molar Volume

Experimental evidence shows the
volume of one mole of ANY GAS at
STP is 22.4 L/mol
OR
VSTP = 22.4 L/mol
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given volume
in Litres
Molar Mass Calculations
N
n
NA
N = n x NA
m
n
M
m=nxM
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v
n
VSTP
v =n x VSTP
Molar Mass Calculations
particles
(N)
x NA
÷ NA
moles
(n)
÷ VSTP
÷M
x VSTP
volume
(v)
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xM
mass
(m)
Chapter 3
Percent Composition (p. 79)

The mass percent of a compound is the
mass of an element in a compound
expressed as a percent of the total mass
of the compound.
mass of element
mass percent 
X 100%
total mass of compound
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Percent Composition
eg. 8.50 g of a compound was analyzed
and found to contain 6.00 g of hydrogen
and 2.50 g of carbon. Calculate the mass
percent for each element.
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Percent Composition

mass percent may be found using the
formula & the molar mass of a compound.
eg.
Find the percentage composition for CH4
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Percent Composition
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Empirical Formulas
An empirical formula gives the simplest
ratio of elements in a compound.
A molecular formula shows the actual
number of atoms in a molecule of a
compound.
Ionic compounds are always written as
empirical formulas
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Empirical Formulas
Compound
butane
Molecular
Formula
C4H10
glucose
C6H12O6
water
H2O
H2O
benzene
C6H6
CH
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Empirical
Formula
C2H5
CH2O
Empirical Formulas (p.87)
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Empirical Formulas
The empirical formula of a compound
may be determined by using the %
composition of a given compound.
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Empirical Formulas



Method:
assume you have 100.0 g of the
compound (ie. change % to g)
calculate the moles (n) for each element
divide each n by the smallest n to get the
ratio for the empirical formula
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Empirical Formulas
eg. A compound was analyzed and found
to contain 87.4% N and 12.6 % H by
mass. Determine the empirical formula
of the compound.
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Empirical Formulas
When finding the EF, the mole ratio may
not be a whole number ratio.
eg. A compound contains 84.73% N and
15.27 % H by mass. Determine the
empirical formula of the compound.

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Empirical Formulas
eg. A compound contains 89.91% C and
10.08 % H by mass. Determine the
empirical formula of the compound.
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Molecular Formulas

The molecular formula of a compound is
a multiple of the empirical formula.
See p. 95
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Molecular Formulas
To find the molecular formula we need
the empirical formula and the molar
mass of the compound
eg. The empirical formula of hydrazine is
NH2. The molar mass of hydrazine is
32.06 g/mol. What is the molecular
formula for hydrazine?

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Molecular Formulas
Example : An unknown was subjected to a
combustion analysis and found to be
85.60% carbon and 14.40% hydrogen.
Mass spectrometry indicated that its molar
mass was 28.06 g/mol. Determine the
molecular formula of the compound.
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CHC analyzer (p. 99 – 101)
1. Describe the operation of a carbon
hydrogen combustion analyzer.
2. 22.0 g of carbon dioxide and 10.8 g of
water is collected in a CHC analysis.
Determine the empirical formula of the
hydrocarbon.
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Formula of a hydrate
To determine the formula of a hydrate:
- calculate the moles of water
- calculate the moles of anhydrous
compound
- determine the simplest ratio
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Formula of a hydrate
eg. Use the data below to determine the
value of x in LiCl• xH2O.
mass of crucible = 26.35 g
crucible + hydrate = 42.15 g
crucible + anhydrous compound= 34.94 g
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mwater = 42.15 – 34.94 = 7.21 g H2O
mLiCl = 34.94 – 26.35 = 8.59 g LiCl
eg. Na2CO3. xH2O
crucible = 15.96 g
crucible + hydrate = 22.19 g
crucible + anhydrous compound = 19.67 g
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eg. CoCl2.xH2O
crucible = 151.96 g
crucible + hydrate = 164.35 g
crucible + anhydrous compound = 158.23 g
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Chapter 4
Stoichiometry
Stoichiometry is the determination of
quantities needed for, or produced by,
chemical reactions.
 Can be gravimetric (mass), gas (Vstp), or
solution stoichiometry (solutions)
 Ratios from balanced chemical equations
are used to predict quantities.

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Stoichiometry – p. 111
Clubhouse sandwich recipe
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Clubhouse sandwich recipe
Fill in the missing quantities:
Slices of
Toast
Slices of
Turkey
Strips of
Bacon
27
66
100
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# of
Sandwiches
12
Mole Ratios

A mole ratio is a mathematical expression
that shows the relative amounts of two
species involved in a chemical change.
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Mole Ratios

A mole ratio
Come from a balanced chemical equation
 Shows the relative amounts of the
reactants/products in moles
 Is the coefficient for the required species
in the numerator and the coefficient for
the given species in the denominator.
(What
you WANT over what you GOT)
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
N2(g) + 3 H2 (g) → 2 NH3 (g)
20
66
140
81
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C3H8 + 5 O2 → 3 CO2 + 4 H2O
# required  # given x
coefficient required
coefficient given
How many moles of CO2 are produced
when 31.5 mol of O2 react?
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C3H8 + 5 O2 → 3 CO2 + 4 H2O
# required  # given x
coefficient required
coefficient given
How many moles of H2O are produced
when 1.35 mol of O2 react?
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C3H8 + 5 O2 → 3 CO2 + 4 H2O
# required  # given x
coefficient required
coefficient given
How many moles of O2 are needed to
produce 31.5 mol of CO2?
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C3H8 + 5 O2 → 3 CO2 + 4 H2O
# required  # given x
coefficient required
coefficient given
How many moles of C3H8 are needed to
react with 0.369 mol of O2?
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C5H12
+
O2 →
# required  # given x
CO2 +
H2O
coefficient required
coefficient given
How many moles of CO2 are produced
when 6.35 mol of O2 react?
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Al(s) + Br2(l) → AlBr3(s)
# required  # given x
coefficient required
coefficient given
How many moles of Br2 are needed to
produce 0.315 mol of AlBr3?
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Mole to Mole Stoichiometry
1.
2.
3.
How many moles of nitrogen gas are needed
to produce 6.75 mol of NH3 in a reaction
with hydrogen gas? (3.38mol)
How many moles of silver would be
produced if 10.0 mol of silver nitrate reacts
with copper metal? (10.0 mol)
How many moles of water are produced
when 20.6 mol of CH4 burns? (41.2mol)
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Mole to Mole Stoichiometry
1.
How many moles of copper would be
produced if 20.5 mol of copper (II) oxide
decomposes? (20.5mol)
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Mass to Mole Stoichiometry
1.
How many moles of water are produced
when 20.6 g of CH4 burns? (2.56mol)
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Mass to Mole Stoichiometry
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Mass to Mass Stoichiometry
eg. Calculate the mass of HCl needed to
react with 3.56 g of Fe to produce FeCl2.
(4.63g)
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Mole to Mass Stoichiometry
1. What mass of CaCl2 is produced when
4.38mol of Ca(NO3)2 reacts with NaCl? (486g)
2. How many moles of copper would be
produced if 20.5 g of copper (II) oxide
decomposes? (0.258 mol)
3. How many moles of water are produced when
5.45 g of C3H8 burns? (0.496mol)
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Ca(NO3)2 + 2 NaCl → CaCl2 + 2 NaNO3
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2 CuO → 2 Cu + O2
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C3H8 + 5 O2 → 3 CO2 + 4 H2O
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Stoichiometry (Chp.4)
Four step stoichiometry:
1. Write a balanced chemical equation
2. Calculate moles given from mass
3. Mole ratio – find moles required
4. Calculate required quantity (mass)
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Stoichiometry (Chp.4)
eg. What mass of CO2 gas is produced
when 45.9 g of CH4 burns ? (126g)
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eg. How many moles of HCl needed to
react with 3.56 g of Fe to produce FeCl2.
(0.127mol)
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Mole Calculations (p. 121 #13)
Fe
3.56 g
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+ 2 HCl
→ FeCl2 + H2
? mol
How many moles of aluminum chloride can be
produced from the reaction of chlorine and
10.8 mol of aluminum ? (10.8mol)
Cl2(g) + Al(s) → AlCl3(s)
How many moles of magnesium are needed to
react with 27 g of iodine to form magnesium
iodide? (0.11mol)
How many grams of nitrogen are needed to
react with 14.0 mol of oxygen to produce
nitrogen dioxide ? (196g)
1 N2(g) + 2 O2(g) → 2 NO2(g)
Mole Calculations (p. 121 #14)
2.34 g
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?L
Limiting Reactant (p. 128)
1.
2.
3.
4.
10.0
15.0
10.0
15.0
g
g
g
g
of
of
of
of
Li requires _______ of Br2
Br2 requires _______ of Li
Li produces _______ of LiBr
Br2 produces ______ of LiBr
This problem has _____ g of excess _____
and will produce ______ g of LiBr
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Limiting Reactant (p. 128)
10.0 g of Li reacts with 15.0 g of Br2.
Calculate the mass of LiBr produced.
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Limiting Reactant (p. 128)


The Limiting Reactant (LR) OR Limiting
Reagent (LR) is the substance that is
completely used in a chemical reaction.
The Excess Reactant is the reactant
that is left over after a reaction is
complete.
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Limiting Reactant (p. 128)
eg. 2.00 g of NaI reacts with 2.00 g of
Pb(NO3)2. Determine the LR and
calculate the amount of PbI2 produced.
 write a balanced equation
 find n for each reactant (Step #2)
 find moles produced by each reactant
(Step #3)
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Pb(NO3)2 + 2 NaI → 2 NaNO3 + PbI2
2.00 g
2.00 g
What mass of calcium carbonate will be
produced when 20.0 g of calcium phosphate
reacts with 15.0 g of sodium carbonate?
(14.2 g)
What mass of barium hydroxide will be
produced when 10.0 g of barium nitrate reacts
with 30.0 g of sodium hydroxide? (6.56 g)
Ba(NO3)2 + 2 NaOH → 2 NaNO3 +Ba(OH)2
What volume of hydrogen gas at STP will be
produced when 10.0 g of zinc metal reacts
with 20.0 g of hydrogen chloride?
Zn + 2 HCl → H2 + ZnCl2
Law of Conservation of Mass (p. 118)
In a chemical reaction, the total mass
of reactants always equals the total
mass of products.
eg. 2 Na3N → 6 Na + N2
When 500.00 g of Na3N decomposes
323.20 g of N2 is produced. How much
Na is produced in this decomposition?
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Law of Conservation of Mass
( p. 118)
eg. To produce 90.1 g of water, what mass
of hydrogen gas is needed to react with
80.0 g of oxygen?
eg. If 3.55 g of chlorine reacts with exactly
2.29 g of sodium, what mass of NaCl will
be produced?
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Percent yield (p. 137)


The theoretical yield is the amount of
product that we calculate using
stoichiometry
The actual yield is the amount of
product obtained from a chemical
reaction
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Percent yield (p. 137)
percent yield 
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actual yield
x 100
theoretical yield