STOICHIOMETRYCh. 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:

Cr2O3FeO NiCl2 S3Cl6 HIHNO3

• 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 Na2CO3x 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!!