Name:____________________________________________
1
Honors Unit 3 Notes – Stoichiometry
Objectives:
1. Students will review the derivation and meaning of atomic mass.
2. Students will be able to calculate molecular and formula masses using data from the periodic
table.
3. Students will understand the term and usage of the concept of a mole and be able to interconvert
between moles, mass, and number of particles (molecules, formula units, atoms, or ions).
4. Students will be able to calculate mass percent composition from formulas of molecular and ionic
compounds, including hydrates. They will also be able to determine the empirical and molecular
formulas, given appropriate percent composition data.
5. Students will be able to write and balance chemical equations, and based on classifying the
equations, they will be able to predict products for simple reactions.
6. Based on balanced equations, students will be able to interconvert between mass and moles of
reactants and products.
7. Students will be able to solve simple limiting reagent problems and determine percentage yield
for a chemical reaction.
3.1 Atomic Masses
1. Atomic Mass –
a. Atomic mass units = amu
i. Standard scale based on the _____________________ isotope.
ii. Mass of C – 12 atom = _____________________ (exactly)
b. Most elements exist in nature as a ______________________ of two
or more ________________________________.
c. A ___________________________ is used to determine atomic mass.
d. To calculate the atomic mass of such elements, you need to know:
i. ______________________ of individual isotopes.
ii. Percent abundance (_________________________________)
-22. Avogadro’s Number
a. A sample of any element with a mass equal to its atomic mass
contains the ______________________________________, NA,
regardless of the identity of the element.
b. Avogadro’s Number, NA = ________________________________
c. It represents the _________________________ of an element in a
sample whose ________________in grams is numerically equal to the
__________________________ of the element.
Example:
6.022 x 1023 H atoms in 1.008 g
atomic mass H = 1.01 amu
6.022 x 1023 S atoms in 32.065 g
atomic mass S = 32.07 amu
3.2 Mass and The Mole
1. Molecular mass –
2. Formula mass –
a. Example: Determine the molecular mass for each of the following:
CaI2
(NH4)2S
Al(NO3)3
C6H12O6
Miss Adams
Honors Chemistry 1
2013-14
-33. Mole –
a. Analogies
4. Molar Mass
a. Molar mass (MM), in __________________________________, is
numerically equivalent to the sum of the _________________ (in amu)
of the atoms in the formula.
b. MM is the mass of one ________________ of a particular substance.
5. Equivalencies (conversion factors)
****Diatomic Elements: Remember BrINClHOF!
Miss Adams
Honors Chemistry 1
2013-14
-4Examples: Use dimensional analysis!!!
1. Calcium carbonate, CaCO3, is the principal mineral found in marble and limestone.
How many moles are in 188.0 g of CaCO3?
2. What is the mass, in grams, of 0.329 mol of spearmint oil, C10H14O?
3. Find the mass of a single lead atom.
4. How many individual lead atoms are in a 1.000 g sample of this metal?
5. Calculate the number of moles of aluminum in a solid cube that measures 3.40 cm
on a side. (d=2.70 g/cm3).
b. How many atoms of aluminum are in the same sample?
Miss Adams
Honors Chemistry 1
2013-14
-56. How many molecules of oxygen, O2, are in 0.00100 grams of this gas?
b. How many atoms?
3.3 Percent Composition by Mass
1. Percent Composition by Mass –
2. Example: Sodium carbonate is a compound used in the manufacture of soap and
glass. Determine the percent composition by mass of each element in this compound.
3. Example: Determine the percent by mass of water in Al2(SO4)3∙18H2O
Miss Adams
Honors Chemistry 1
2013-14
-64. Example: Magnetite, Fe2O3, is one of the principal iron containing ores. How
much elemental iron can be obtained from a metric ton (103 kg) of this ore,
assuming 100% recovery? (Hint: first find % iron in Fe2O3)
3.3 (cont.) Empirical Formula Calculations [EBOOK Section 5.4]
A. In chemical analysis we determine the _____________________of each element in
a given amount of pure compound and derive the _______________or
________________ formula.
Percent to Mass
Mass to Moles
Divide by Small
(Multiply ‘til Whole)
Sample Problem: A compound of B and H is 81.10% B. What is its empirical formula?
1. Steps to solve problem:
1. Because it contains only B and H, it must contain 18.90% H.
2. ALWAYS ASSUME 100.0 g SAMPLE!
3. In 100.0 g of the compound there are 81.10 g of B and 18.90 g of H. (Just drop
percents and make them into grams!)
4. CONVERT TO MOLES: Calculate the number of moles of each element.
81.10 g B •
1 mol
= 7.502 mol B
10.81 g
18.90 g H •
1 mol
= 18.75 mol H
1.008 g
5. Now, recognize that atoms combine in the ratio of small whole numbers.
6. Take the moles of B and H. Always divide EACH number of moles by the
smallest mole value.
Miss Adams
Honors Chemistry 1
2013-14
18.75 mol H
= 2.499 mol H  2.5 mol H
7.502 mol B
7.502 mol B
 1.000 mol B
7.502 mol B
-77. But we need a whole number ratio! Must multiply EACH mole value by the SAME
smallest integer available to obtain whole numbers.
8. 2.5 mol H • 2 = 5 mol H
1.0 mol B • 2 = 2 mol B
Common Possible Endings
9. EMPIRICAL FORMULA = B2H5
.33
.25
.67
.50
x
x
x
x
3
4
3
2
Example: Bicarbonate of soda is used in products like Alka-Seltzer and generally relieves
an upset stomach. Determine the empirical formula of this compound based on the
following percent composition: 27.36% Na, 1.200% H, 14.30% C, 57.14% O.
Example: A 25.00 gram sample of an orange compound contains 6.64 g of potassium,
8.84 g of chromium, and 9.52 g of oxygen. Find its empirical formula.
Miss Adams
Honors Chemistry 1
2013-14
-8-
Molecular Formula from the Empirical Formula Calculations
SAMPLE PROBLEM: A compound of B and H is 81.10% B. What is its molecular formula?
a. Is the molecular formula B2H5, B4H10, B6H15, B8H20, etc.?
b. We need to do an EXPERIMENT to find the MOLAR MASS.
c. Here, the experiment gives 53.3 g/mol.
d. Compare with the mass of the empirical formula [B2H5 = 26.66 g/unit]
e. Find the ratio of these masses.
2 units of B2 H 5
53.3 g/mol
=
26.66 g/unit of B2 H 5
1 mol
f.
Multiply all of the subscripts by the ratio and obtain the molecular formula, B4H10
Example: A certain compound has the empirical formula C2H4O. Its molar mass is about
90 g/mole. What is the molecular formula?
Example: A hydrate of magnesium iodide has the formula MgI2 ∙ X H2O. To determine the
value of X, a student heats a sample of the hydrate until all the water is gone. A 1.628 g
sample of hydrate is heated to constant mass of 1.072 g. What is the value of X?
Miss Adams
Honors Chemistry 1
2013-14
-9-
3.4 Writing & Balancing Chemical Equations
1. All chemical reactions have two parts:

Reactants -

Products-
2. In a chemical reaction:

The way atoms are ______________________is ________________________.

Atoms aren’t created or destroyed; they just ________________________________
____________________________________________________________________.

Can be described using sentences, symbols or word equations
3. Parts of the equation:

The ______________separates the reactants from the products; means ____________

The _________________________= ________________

____________________are used to describe the number of _______________in a
_____________________________.

_______________________are used to describe the number of
________________________________________
in the _____________________.
They are the only things changed when _________________________a reaction.
4. States of matter





(s) after the formula –_________________________
A _____________________________is a solid formed in a reaction
(g) after the formula -_______________
(l) after the formula –________________
(aq) after the formula - ________________________________________
Miss Adams
Honors Chemistry 1
2013-14
- 10 

 used after a product indicates a _____________________
 used after a product indicates a _____________________
5. Other Symbols Used in Equations

indicates a reversible reaction (More later).

show that heat is supplied to the reaction.

is used to indicate a catalyst used supplied, in this case, platinum.
6. Balancing Chemical Equations
a. Law of Conservation of Mass must be obeyed…therefore,
__________________________________________________.
b. _______________________must be added so the number of ________________
atoms equals the number of ______________________atoms!
7. Hints for Balancing Chemical Equations
____ Al(s) + _____ Br2 (l)  _____ Al2Br6 (s)
____ Na3PO4 + ____ Fe2O3  ____ Na2O + ____ FePO4
Miss Adams
Honors Chemistry 1
2013-14
- 11 -
Types of Chemical Reactions
***See Dancers Handout for examples!
1.
2.
3.
4.
5.
Mass Relations for Equations
1. Mole ratios or stoichiometric factors:
a. Equations are balanced in order to obey the Law of Conservation of Mass.
As a result, there are mathematical relationships between the substances in a
balanced chemical which are the foundation of stoichiometry.
CS2 + 3O2  CO2 + 2 SO2
1. Interpretation in terms of moles
2. Conversion factors (mole ratios)
Example: Using the equation CS2 + 3O2  CO2 + 2SO2 determine:
a) the number of moles of oxygen required to react with 1.38 mol of carbon disulfide
b) the number of moles of SO2 produced from 1.38 moles of carbon disulfide.
Miss Adams
Honors Chemistry 1
2013-14
- 12 Example: For 2 NH3 + H2SO4  (NH4)2SO4
determine:
a) the mass of product possible when 1.43 mol of NH3 are reacted with an excess
of sulfuric acid.
b) the mass of NH3 required to react completely with 35.00 g of sulfuric acid.
c) the mass of sulfuric acid required to form 1000 grams of product.
Example: How many milliliters of liquid water can be produced by the combustion
of 775 mL of octane with oxygen? Assume that the volumes of the octane and the
water are measured at 20oC where the densities are 0.7025 g/mL for octane and
0.9982 g/mL for water. Remember to balance the reaction.
____ C8H18(l) +
Miss Adams
Honors Chemistry 1
2013-14
____ O2(g) 
____ CO2(g) +
____ H2O(l)
- 13 -
Limiting Reactants
a. Theoretical yield –
b. Experimental (Actual) yield –
c. Limiting reactant –
d. Excess Reactant –
e. Steps for determining limiting reactant:
Miss Adams
Honors Chemistry 1
2013-14
- 14 Example: For the following reaction,
CS2(l) + 3 O2(g)  CO2(g) + 2 SO2(g)
a) Determine the theoretical yield of product if one starts with 1.20 mol of CS 2 and
3.83 mol of O2.
b) Determine the theoretical yield of product if one starts with 105 g of CS2 and
145 g of O2.
Percent Yield
Miss Adams
Honors Chemistry 1
2013-14
- 15 Example:
a. For the following reaction, determine the theoretical yield if one starts with1.20 g
of antimony and 2.40 g of iodine.
2 Sb(s) + 3 I2(s)  2 SbI3(s)
Determine theoretical yield first:
b. If 3.00 g of product are actually formed, what is the percent yield?
Miss Adams
Honors Chemistry 1
2013-14