STOICHIOMETRY
Ch. 3
Lab A Partners
• Group 1: Taylor and Ashley
• Group 2: Jackson and Nelson
• Group 3: Garrett and Kenney
Check for Understanding
• Name the following:
Cr2O3
HI
FeO NiCl2
HNO3
S3Cl6
• Write the chemical formula for the following:
Iron (III) Oxide
Calcium Nitride
Diphosphorus hexaiodide
Carbonic acid
Hydrochloric acid
Questions?
Heads Up
• Lab on Silver Oxide due Tomorrow!
• Lab on Alum on Wednesday, January 14th .
• Exam 1 Thursday, January 15th
• Covers Chapters 1, 2, and 3
Chapter 3 - Stoichiometry
• Atomic Mass
• The Mole
• Molar Mass
• Percent Composition
• Empirical and Molecular Formulas
• Balancing Chemical Equations
• Stoichiometry—Limiting Reactants
• Percent Yield
Practice Problems
• Complete the Review Packet
• Unit 1 Packet Practice
• Book problems starting on page 117:
• #28, 34, 36, 39, 41, 59, 81, 89, 99, 103
• **Please note that this is a small list to complete at minimum in
preparation for the exam.**
Learning Objectives – Part 1
• I can calculate average atomic mass of an atom from
mass spectrometry and isotopic data.
• I can calculate the molar mass of a compound and the
percent composition of an element in a compound.
• I can convert between mass, moles, and numbers of
representative particles (atoms, molecules, formula
units) of a substance.
Atomic Masses
• Instituted in 1961, based on Carbon-12 as the
standard.
• Carbon is assigned a mass of exactly 12 atomic mass units (amu), and
the masses of all other atoms are given relative to this standard.
• Mass Spectrometer
• Most accurate method for comparing masses of atoms.
• The amount of path deflection for each ion depends on its mass (the most
massive ions are deflected the smallest amount)
• The areas of the “peaks” or the heights of the bars indicate the relative
abundances of the different isotopes.
Example – Atomic Masses
• An element consists of 1.40% of an isotope with mass
203.973amu, 24.10% of an isotope with mass
205.9745amu, 22.10% of an isotope with mass 206.9759
amu, and 52.40% of an isotope with mass 207.9766
amu. Calculate the average atomic mass and identify the
element.
(Hint: add products and divide by 100)
Check for Understanding
• Page 117 Question 28
How does this apply?
• Most elements occur in nature as mixtures of
isotopes; thus atomic masses are usually
average values. (take into account relative
abundance of isotopes 98.89% C-12 and
1.11% C-13.
• The mass for each element listed on the
periodic table is an average value based on
the isotopic composition of the naturally
occurring element. (NO atom of hydrogen
actually has the mass 1.008)
The Mole
• Definition:
• The number equal to the number of carbon atoms in exactly 12 grams of
pure C-12. This number has been determined (via Mass Spectrometry) to be
6.022 *10 23 (Avogadro’s Number).
• One mole of something consists of 6.022 *10
23
units or particles of that substance. (Like a dozen
or ream)
• Video
• The mass of 1 mole of an element is equal to its
atomic mass in grams.
Example – The Mole
• Calculate the mass of 500. atoms of iron
(Fe).
• Diamond is a natural form of pure carbon.
What number of atoms of carbon are in a
1.00-carat diamond (1.00carat = 0.200g)?
Molar Mass
• The mass in grams of one mole of the
compound.
• Obtained by summing the masses of the
component atoms.
• Example
• Problem 40 and 42 on page 118.
• Review Packet
Percent Composition
• Mass percent of the elements can be
determined from the formula of the compound
by comparing the mass of each element
present in 1 mole of the compound to the total
mass of 1 mole of compound.
• Hint: Always check that percentages add up to 100.
• Example
• Number 60 page 119
Empirical/Molecular Formulas
• Empirical – Molecular Formula Determination
• 1. Mass Percent to Grams (100 g sample)
• 2. Convert each element from grams to moles
• 3. Divide all mole values by the smallest.
• 4. Subscripts obtained if whole number. =
Empirical Formula
• 5. Divide molar mass by empirical formula
mass to obtain integer to multiple all
subscripts in formula. “Method 1”
Calculating Empirical Formulas
Example Problem:
A compound of nitrogen and oxygen is
analyzed and a sample weighing
1.587g is found to contain 0.483 g N
and 1.104g O. What is the empirical
formula?
Check for Understanding
A sample compound weighing 83.5 g
contains 33.4 g of sulfur. The rest is
oxygen. What is the empirical formula
and chemical name?
Empirical/Molecular Formulas
• Empirical – Molecular Formula Determination
• 1. Mass Percent to Grams (100 g sample)
• 2. Convert each element from grams to moles
• 3. Divide all mole values by the smallest.
• 4. Subscripts obtained if whole number. =
Empirical Formula
• 5. Divide molar mass by empirical formula
mass to obtain integer to multiple all
subscripts in formula. “Method 1”
Example
• A sample of 100g of acetic acid contains
39.9 g C, 6.7 g H, and 53.4 g O. The molar
mass of the compound was determined by
experiment to be 60.0 g/mol. Determine the
molecular and empirical formula.
Check for Understanding
• Succinic acid is a substance produced by
lichens. Chemical analysis indicates it is
composed of 40.68% C, 5.08% H, 54.24%
O, and has a molar mass of 118.1 g/mol.
Determine the empirical and molecular
formula.
Learning Objectives –Part 2
• I can write a balanced chemical equation and predict the
amount of product formed from a given mass of reactant
or the amount of reactant required to produce a desired
amount of product.
• I can identify limiting reactants, and calculate the amount
of product formed when given the amounts of all the
reactants present.
• I can calculate the percent yield of a reaction.
Chemical Equations
• Chemical Reactions
• Reorganization of the atoms in one or more substances.
• Bonds are broken and new ones form.
• Atoms are neither created nor destroyed.
• All atoms present in the reactants must be accounted for among the products.
= Balancing a chemical equation!
• Reactants (initial substances) left of arrow
• Products (substances produced) right of arrow
• The relative number of reactants and products
in a reaction are indicated by the coefficients in
the balanced equation.
Balancing Chemical Equations
• ALWAYS check to make sure the equation you
are working with is balanced!!!
• Use element tables to help.
• Writing and Balancing the Equation for a
Chemical Reaction
• 1. Determine what reaction is occurring. What are the reactants, the
products, and the physical states involved?
• 2. Write the unbalanced equation that summarizes the reaction described
in step 1.
• 3. Balance the equation by inspection, starting with the most complicated
molecule(s). Determine what coefficients are necessary so that the same
number of each type of atom appears on both reactant and product sides.
Do NOT change the identities (formulas/subscripts) of any of the
reactants or products.
Example
• Nitrogen and oxygen combine to form dinitrogen
pentaoxide.
• Calcium nitrate reacts with lithium sulfide forming calcium
sulfide and lithium nitrate.
Practice
• Solutions of lead (II) chloride and sodium chromate react to
produce a precipitate of
lead (II) chromate and a solution of sodium chloride.
• Solid calcium reacts with solid sulfur to produce solid
calcium sulfide.
• Hydrogen gas reacts with fluorine gas to produce hydrogen
fluoride gas.
Exit Ticket!
• Page 120 #82
Lab Questions?
• Typed Lab Report Due Today!!
Heads UP!
• Exam Thursday!
• Lab A Report Due TODAY!
• Lab B Pre-lab Due TOMORROW
• Closed-toed shoes tomorrow!
• Practice Problems Ch. 3
Warm-up Question
Write the following balanced chemical reactions.
• Liquid carbon disulfide reacts with oxygen gas to
produce carbon dioxide gas and sulfur dioxide gas.
• Aqueous solutions of sodium chloride and silver
nitrate react to produce aqueous sodium nitrate and
a precipitate of silver chloride.
• Bubbling chlorine gas through a solution of
potassium iodide gives elemental iodine and a
solution of potassium chloride.
Stoichiometry 
Flowchart on page 104
• 1. Balance the equation for the reaction
• 2. Convert the known mass of the reactant or
product to moles of that substance.
• 3. Use the balanced equation to set up the
appropriate mole ratios.
• 4. Use the appropriate mole ratios to calculate the
number of moles of the desired reactant or product.
• 5. Convert from moles back to grams if required by
the problem.
• Example
• #90 page 121
Example
•
Calculate the mass of
hydrochloric acid needed to react
with 10.0 g zinc (Hint: Single
Replacement).
• (11.2 g HCl)
Practice
•
Calculate the mass of oxygen produced if 2.50 g of
potassium chlorate are completely decomposed by
heating.
•
(0.976 g O2)
Practice
•
If 20.0 g of magnesium react with excess
hydrochloric acid, how many grams of magnesium
chloride are produced?
•
(78.4 g MgCl2)
Practice
•
How many grams of chlorine gas must be reacted
with excess sodium iodide if 10.0 g of sodium
chloride are needed?
•
(6.07 g Cl2)
Practice
•
What mass of copper is required to replace silver
from 4.00 g of silver nitrate? (Hint: Copper (II)
product is formed!)
•
(0.747 g Cu)
Limiting Reactants
• Think of this as just a multi-part stoichiometry problem.
• Example:
• The reaction between solid sodium and iron (III) oxide
is one in a series of reactions that inflates an
automobile airbag.
•
Na (s) +
Fe2O3 (s) Na2O (s) +
Fe(s)
•
• If 100.0 g of Na and 100.0 g Fe2O3 are used in this
reaction, determine
• The limiting reactant
• The reactant in excess
• The mass of solid iron produced
• The mass of excess reactant that remains after the reaction is complete
Practice
• Zinc metal reacts with hydrochloric acid by the following
reaction:
Zn(s) + 2HCl(aq)
ZnCl2(aq) + H2(g)
• If 0.30 moles of zinc is added to hydrochloric acid
containing 0.52 mol HCl, how many moles of H2 are
produced?
Practice
• In a process for producing acetic acid, oxygen gas is bubbled
into acetaldehyde, CH3CHO, containing manganese (II)
acetate under pressure at 60C.
• 2CH3CHO(l) +
O2(g)
2HC2H3O2(l)
• In a laboratory test of this reaction, 20.0 g of CH3CHO and
10.0 g of O2 were put into a reaction vessel.
• How many grams of acetic acid can be produced by this
reaction from these amounts of reactants?
• How many grams of excess reactant remain after the reaction
is complete?
Percent Yield
• Theoretical Yield
• The amount of a product formed when the
limiting reactant is completely consumed.
• Percent Yield
• The actual yield of product is often given as a
percentage of the theoretical yield.
• (Actual yield ÷ Theoretical Yield) x 100% =
Percent Yield
HCl(aq) + NaOH(aq)  NaCl(s) + H2O(l)
Start with 5.0 grams of NaOH and after the reaction I have
6.3 grams of NaCl.
What is the percent yield of the reaction?
With A Partner
• A student runs a reaction to prepare 40.0g of aspirin and
yet recovers only 15.5g. What is the percent yield (please
show/be able to explain your work)?
CH4 + 2O2  CO2 + 2H2O
If you react 5.00 g of methane, CH4, with 12.00g of oxygen,
O2, what is the theoretical yield of H2O?
After completing the reaction there was 4.60g of H2O, what
is the percent yield?
Quiz Unit 1
• Complete Individually!
• Pre-lab for Lab B DUE TOMORROW!
• Read through lab handout (in black binder)
• Complete the 2 pre-lab questions
• Write an abstract of Part 2 (omit Part 1 –
sample data provided)
Lab Safety
• Lab Safety Video
• Lab Safety Contract
• Lab Expectations
Labs over Ch. 3
• Lab A: Empirical Formula
• Complete Lab Report
• Due _______
• Lab B: ALUM Metal Hydrate Formula
• Complete Pre-lab
• Due _________
Practice Problems
• Complete the Review Packet
• Book problems starting on page 117:
• #28, 34, 36, 39, 41, 59, 81, 89, 99, 103
• **Please note that this is a small list to complete at minimum in
preparation for the exam.**
Quiz Analysis
• Locate a DIFFERENT colored pen/pencil.
• 3 minutes reflect individually.
• 3 minutes communicate with a peer.
• Questions?!
Lab Report Revision Process
• Different colored pen/pencil or separate sheet ATTACHED
to draft 1.
• Review
• Individually
• With Peer
• With Mrs. Bechtum
• DUE: Tuesday, January 20th!
• Questions?
Lab B: Analysis of ALUM
• Pre-Lab Questions
• Safety
• Part 1: Sample Data
Trial #1
Trial #2
Measured melting point
92.1°C
92.8°C
Literature melting point
92.5°C
92.5°C
Alum Lab Pre-lab Questions
• 1. When measuring a melting point, why is it necessary to
raise the temperature very slowly in the vicinity of the
melting temperature?
• 2. Washing soda is a hydrated compound whose formula
can be written Na2CO3x H2O, where x is the number of
moles of H2O per mole of Na2CO3. When a 2.123 g
sample of washing soda was heated at 130 C, all of the
water of hydration was lost, leaving 0.787 g of anhydrous
sodium carbonate. Calculate the value of x.
Lab B: Analysis of ALUM
• Part 2:
• It is important to heat slowly at first so that the evolving water of
hydration does not carry the alum crystals with it!
• Multiple heatings (record number and duration of heatings recommend 5 minute increments)
• The remainder of the class period is yours to
review/prepare for tomorrow’s exam and/or begin
working on Lab B report.
Let me know if you have questions!
Lab B Partners
• Group 1:
• Group 2:
• Group 3:
Tomorrow…
• Exam 1!
• 2 Parts
• Multiple Choice – 40 minutes (no calculator)
• Free Response – 40 minutes (calculator)
• NO ELECTRONIC DEVICES ALLOWED IN CLASSROOM!
• Begin working on Lab B Report (Due Friday!!)
UNIT 1 Exam!
• 2 Parts
• Multiple Choice – 40 minutes (no calculator)
• Free Response – 40 minutes (calculator)
• NO ELECTRONIC DEVICES ALLOWED IN CLASSROOM!
• NO Talking!
• You CAN do this!!