honors chemistry...1 honors chemistry unit 7 (chapter 8) chemical composition oh moleo! the greatest...
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Honors Chemistry
Unit 7
(Chapter 8)
Chemical Composition
Oh Moleo!
The greatest love story ever moled.
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Textbook Problem Sets
Problem Set #1: Read section 8.4 (221 – 226);
pgs. 239-240: answer questions 25, 26, 27abc, 31ab, 35cd, 39ab
Problem Set #2: Read section 8.7 (230 – 235); answer Self Check Exercises 8.8, 8.9, 8.10
Problem Set #3: Read section 8.8 (235 – 237); answer Self Check Exercises 8.11; pg. 242: answer questions 76, 77
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Chemistry Particles
1. Atom: the smallest particle of an _______________________.
Examples: ______________________________________________________________
______________________________________________________________
2. Molecule: The smallest particle of a ______________________________ ________________________________. In general,
covalent molecular compounds are composed of ____________________________ _____________________ that are covalently bonded together.
Examples: ______________________________________________________________ ______________________________________________________________
3. Ion: The term used by chemists to describe a _______________ ________________. More specifically, positive ions are
called __________________ and negative ions are referred to as _________________________.
Examples: ______________________________________________________________ ______________________________________________________________
4. Formula Unit: The name given to a particle of an ______________________________ ________________________________, a
compound composed of _________________ held together by an ___________________________________ _________________________________. In this class, we will assume that all compounds that begin with a ____________________
are ionic.
Examples: ______________________________________________________________ ______________________________________________________________
a) name or write the formula for the substance and
b) match the representative particle with the appropriate substance
Substance Name/Formula Letter
1. SO3 a. atom
2. Na2CrO4
b. molecule c. ion
3. radium
d. formula unit
4. N2
5. Ferric hydroxide
6. Sn+4
7. tribromine heptachloride
8. P
9. K3N
10. ClO-1
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Why did chemists create their own counting unit called the MOLE?
1. _____________________________________________________________________________ _____________________________________________________________________________
_____________________________________________________________________________
Mole (mol): the SI unit of measure used to count numbers of representative particles, such as, ________________, ________________, ______________, or __________________.
1 mole = ____________________ particles
6.02 x 1023 particles = _______________________
Avogadro’s Number = ______________________ = ___________________
The mole is simply a unit used to count the smallest unit of a substance!!! 1 mole of C = 6.02 x 1023 _________________ of C 1 mole of O2 = 6.02 x 1023 _________________ of O2 1 mole of MgO = 6.02 x 1023 _________________ of MgO 1 mole of Zn+2 = 6.02 x 1023
_________________ of Zn+2
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1 mole of NaCl = ___________________ ____________________ of NaCl What are formula units composed of?______________________________ 1 f.u. of NaCl = ________________ Na+ ion(s) 1 f.u. of NaCl = ________________ Cl- ion(s)
1. Water (H2O) - What is the representative particle in H2O? ___________________
2. Copper (Cu) - What is the representative particle in Cu? __________________
3. Sodium Chloride (NaCl) – What is the representative particle in NaCl?
1 mole of H2O = ___________________ ____________________ of H2O What are molecules composed of?______________________________ 1 molecule of H2O = ________________ H atom(s) 1 molecule of H2O = ________________ O atom(s)
_______________
+
+
+
+ +
+
+ + +
1 mole of Cu = ___________________ ____________________ of Cu What is copper wire composed of? ____________________________
- +
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Complete the Following Chart:
SUBSTANCE CHEMICAL
FORMULA
REPRESENTATIVE
PARTICLE
REPRESENTATIVE PARTICLES IN 1 MOLE
atomic nitrogen
nitrogen gas
calcium ion
beryllium fluoride
sucrose
C12H22O11
What are the masses of the elements in each of the following beakers?
Carbon (C) = _____ g Copper (Cu) = _____ g
Silicon (Si) = _____ g
Sulfur (S) = _____ g
Zinc (Zn) = _____ g
Lead (Pb) = _____ g
There is _________of each element
in each beaker!
2 different substances of equal __________________
have an equal # of ____________________________,
BUT differ in ________________________________.
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Where have you seen these masses before? ____________________________________ What did we call these masses? ______________________________________________ What was the unit of measurement? _________________________________________
1 carbon atom = ________________ The mass on the period table is also known as the _______________________________
1 mole C atoms = ________________ Molar Mass: ____________________________________________________________
Units = __________________________ SO…. 1 mole of C = ____________________ atoms of carbon 1 mole of Cu = ____________________ atoms of Cu 1 mole of Si = ___________________ atoms of Si BUT….. 1 mole of C = _________ grams of carbon 1mole of Cu = _________ grams of Cu 1 mole of Si = _________ grams of Si
INTERPRETING FORMULAS
SUBSTANCE PARTICLES MOLES
Al2O3
1 _____________ Al2O3
2 ______ ______
3 ______ ______
1 mole of Al2O3
____ mol(s) Al+3 ions
____ mol(s) O-2 ions
CO2
1 ___________ of CO2
1 ______ of ______
2 ______ of ______
1 mole of CO2
____ mol of C atoms
____ mol of O atoms
C
12.0115
Cu
63.55
Si
28.09
S
32.07
Zn
65.39
Pb
207.2
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THE MOLAR MASS OF COMPOUNDS
a.) Calculate the molar mass of ethanol, C2H6O.
b.) How many moles of ethanol are in 105 grams? 2.
a.) Calculate the molar mass of calcium chloride.
b.) What is the mass, in grams, of a 3.5 mol sample of calcium chloride?
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3. a.) Calculate the molar mass of aluminum sulfate.
b.) How many formula units are contained in 50. grams of aluminum sulfate?
Let’s Take a Look Inside
4. What is the mass of all the fluorine atoms in 0.15 moles of sulfur hexafluoride? 5. How many phosphide ions are in 97.4 grams of magnesium phosphide?
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Mole-Mass Relationships
Even though we know that a mole contains 6.02 x 1023 particles, for simplicity and practical purposes, we
will use the 12 particles to represent one mole.
1. Each one of the following jars contains a different number of moles of oxygen gas. So Jar A contains
one mole of oxygen gas. In jars B and C, draw the correct number of molecules in the jars.
A B C
1 mole O2 0.5 mole O2 0.25 mole O2
Molar Mass = ___________ Molar Mass = ___________ Molar Mass = ___________
Mass = ___________ Mass = ___________ Mass = ___________
2. For jars D, E, and F, use the number of molecules in the jars to answer the following questions.
D E F
_____________ mole O2 _______________mole O2 _______________mole O2
Mass = ______________ Mass = ________________ Mass = ________________
Number of O atoms_______ Number of O atoms_______ Number of O atoms______
3. Jars G, H, and I contains ammonia gas. Draw the correct number of molecules in the jars and answer
the following questions. When drawing the molecules, be sure to take into account the total number
of atoms in ammonia.
G H I
1 mol NH3 2 mol NH3 0.666 mol NH3
Molar Mass = ___________ Molar Mass = ___________ Molar Mass = ___________
Mass = __________ Mass = ___________ Mass = ___________
Number of N atoms _____ Number of N atoms ______ Number of N atoms _________
Number of H atoms _____ Number of H atoms ______ Number of H atoms _________
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The Mole – The Heart of Chemistry
1 mole = 6.02 x 1023
particles 1 mole = molar mass of the substance
To help answer these questions, read pages 214 – 221 in your textbook.
1. What is one amu equivalent to in grams? ________________________________________
2. In 26.98 grams of Aluminum, there are __________________________ aluminum atoms present.
3. Assume you have a sample of sodium weighing 11.50 grams.
a. How many atoms of sodium are present in the sample? _________________________
b. What mass of potassium contains the same number of atoms as there are in the sodium sample?
4. Which has the smaller mass, 1 mole of helium atoms or 4 moles of hydrogen atoms?
5. Using the factor-label method, determine the number of mass of the following elements.
a. 3.00 moles of uranium
b. 7.50 moles of boron
c. 0.025 moles of fluorine gas, which exists not as an atom but as a diatomic molecule, F2.
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6. Using the factor-label method, determine the number of moles in each of the following.
a. 25.0 grams of lithium
b. 8.50 grams of oxygen gas, O2
7. Using the factor-label method, determine the number of atoms in each of the following.
a. 2 moles of iron
b. 28 g of nitrogen gas
8. With the price of gold skyrocketing, you decide to trade your gold class ring at the Pawn Shop for
some quick cash. The pawn broker, also a retired chemistry teacher, says he will give you $500/mole.
If the ring weighs 4.51 g, how much money will you receive for your ring?
9. Diamonds are a very pure form of carbon.
a. Calculate the number of moles of carbon in a 1 carat diamond. (1 carat = 0.200 g)
b. Calculate the number of atoms of carbon in the 1 carat diamond.
10. Calculate the average mass, in grams, of 1 atom of oxygen.
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Names: ____________________________ ________________________ Period: __________
Practice Makes Perfect
1. Your uncle, a rare coin collector, possesses a U.S. penny from 1814 – when the pennies minted in the United States were 100% copper. If this penny weighs 3.13 g, how many atoms of copper are present in the penny?
2. Acetylsalicylic acid (C9H8O4) – more commonly known as Aspirin – is an important pain reliever and blood thinning medication taken by millions of people every day. If you take one 325 mg tablet of Aspirin, how many molecules of acetylsalicylic acid are you ingesting? How many hydrogen atoms would this involve?
3. With a strange request, your father asks you for a glass of 4.53 x 1024 molecules of water. How
much would this water weigh?
4. The Hope diamond is one of the world’s largest, at 45.52 carats. If one carat = 0.200 g, and the diamond is composed entirely of carbon, how many carbon atoms would be present?
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the mole REVIEW
Solve the following problems on the next page.
1. How many moles are in 333 grams of stannous fluoride? 2. How many nitrate ions are in 333 grams of calcium nitrate?
3. How many oxygen atoms are needed to makeup 0.5 moles of carbon dioxide?
4. Find the mass in grams of 0.720 moles of hydrogen gas.
5. Compare the number of atoms in a mole of neon to the number of atoms in a mole of
calcium.
6. How many molecules are in 0.59 moles of carbon tetrachloride?
7. Find the mass in grams of 3.5 x 1022 formula units of sodium sulfate.
8. List the four types of representative particles and give an example of each.
9. The statue of liberty in New York harbor is made of 2.00 x 105 pounds of copper sheets bolted to an iron framework. How many moles of copper does this represent? (hint: 1 lb = 454 g)
10. Ibuprofen, the active ingredient in many nonprescription pain relievers, has the formula,
C13H18O2.
a. If the tablets in a bottle contain a total of 33 grams of ibuprofen, how many moles of ibuprofen are in the bottle?
b. How many molecules are in the bottle?
c. What is the total mass in grams of carbon in this much ibuprofen?
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The Mole
1. Complete these statements by supplying the missing quantity.
a. One mole of O atoms contains ______________________________ atoms. b. One mole of O2 molecules contains ______________________________ molecules. c. One mole of O2 molecules contains ______________________________ O atoms. d. One mole of O atoms has a mass of ______________________________ grams. e. One mole of O2 molecules has a mass of ______________________________ grams.
Solve the following problems on a separate sheet of paper. Use the factor label method.
2. Arrange the following in order of increasing mass: a. 1.06 moles of sulfur tetrafluoride b. 8.7 x 10
25 molecules of dichlorine heptoxide
c. 4.17 x 1025
atoms of argon d. the oxygen contained in 59 grams of water
3. Aspartame, marketed under the name NutraSweet, is an artificial sweetener that is 160 times sweeter
than table sugar when dissolved in water. The molecular formula for the compound is C14H18N2O5. How many moles of aspartame are contained in 10.0 grams of the sweetener?
4. Dimethylnitrosamine, (CH3)2N2O, is a carcinogen that may be formed in foods, beverages, or gastric juices from the reaction of the nitrite ion (used as a food preservative) with other substances. What is the mass of 1.0 x 10
14 molecules of dimethylnitrosamine?
5. Chloral hydrate (C2H3Cl3O2) is used as a sedative and hypnotic drug. It is the compound used to make
“Mickey Finns” in detective stories. a. How many moles of chloral hydrate are contained in 500.0 g of the drug?
b. How many chlorine atoms are contained in 5.0 grams of chloral hydrate? 6. One atom of an unknown element has a mass of 10.55 x 10
-23 grams. What is the molar mass of the
element? Name the element.
7. How many atoms of oxygen are in each of the following: a. 5.0 moles of manganese dioxide b. 255 grams of magnesium carbonate c. 5 x 10
18 molecules of water
8. The United States Environmental Protection Agency (EPA) sets the maximum safe level for lead in the
blood at 240 µg/liter. A 1.00 mL sample of a patient’s blood contains 1.50 x 10-8
moles of lead. Express the patient’s lead level in µg/liter to see if there is danger of lead poisoning.
9. A ream of NA school paper (500 sheets) is 5 cm high. What is the height, in meters, of a stack of Avogadro’s number of sheets of paper?
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Deadly Problems
Name _________________________________________________________ Period ______________
1. Sarin, C4H10FO2P, is an extremely toxic substance whose sole application is as a nerve agent. It is a colorless, odorless gas with a lethal dose of 0.01 mg/kg of body mass. If the average high school student has a mass of 75 kg, how many moles of sarin would be considered a lethal dose?
2. Strychnine, C21H22N2O2, produces some of the most dramatic and painful symptoms of any toxic
reaction. For this reason, it is often used in literature and film. For example, Norman Bates used it in the classic thriller, “Psycho” to poison his mother. If strychnine’s lethal dose is 3.747 x 10-5 mol/kg body mass, how many grams are needed to poison a 135 lb teen? (HINT: 1 lb = 454 g)
How many atoms of nitrogen are present in this lethal dose?
3. Cicutoxin, C17H22O2, is found in the water hemlock. Historically, hemlock berries were reportedly used to poison Socrates. How many molecules of this poison are contained in a 230 gram sample, the lethal dose for a horse?
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Name ______________________________________________________________________________________
The Little Mole Lab OBJECTIVE: After writing your first, middle, and last name 15 times on the chalkboard, determine
1. the mass of chalk on the chalkboard. 2. the number of moles of chalk on the chalkboard. 3. the number of formula units of chalk on the chalkboard. 4. the number of atoms of oxygen on the chalkboard.
MATERIALS: chalk chalkboard electronic balance PROCEDURE: That’s for you to figure out. But remember, the procedure contains only the steps necessary to collect the data needed to meet your objectives. It does not include the steps needed to calculate your answers to the objectives. CONCLUSION: On a separate sheet of lined paper,
1. List the steps in your procedure. Be specific. (3 points) 2. Use a ruler to construct a data table. (2 points) 3. Calculate the answers to the four objectives. Show all your work. (8 points) 4. Use a ruler to construct a calculations table. (2 points)
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Percent Composition
It is frequently useful to know the percent composition by mass of a chemical compound. The percent of iron in iron III oxide may be used to calculate the mass of iron in an iron ore. Or the percent of oxygen in potassium chlorate may be useful if it is to be used as a source of oxygen. The percent composition of a compound is the percent by mass of each element in a compound. The percent composition is the same, no matter what the size of the sample. The percent by mass of an element in a compound can be determined by the following equation:
mass of element molar mass of compound
**A good check is to see if the results add up to 100%.
Because of rounding off or experimental error, the total may not always be exactly 100%.** EXAMPLE: Calculate the percent composition of copper I sulfide. 1. Write the formula or give its name. 2. Determine the mass of each element present in the compound. 3. Calculate the molar mass of the compound. 4. Use the formula above to calculate the % of each element. 5. Check to be sure the results add up to 100%.
x 100 = % mass of element
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Percent Composition of a Hydrate
As some compounds crystallize from a water solution, they trap water molecules. Sodium carbonate forms such a hydrate, in which 10 water molecules are present for every formula unit of sodium carbonate. Define:
a. hydrate - _____________________________________________________________________________ _____________________________________________________________________________________
b. anhydrous - ___________________________________________________________________________
_____________________________________________________________________________________
EXAMPLE: Find the percent of water in sodium carbonate decahydrate.
1. Write the formula. 2. Determine the mass of the water present in the compound (don’t forget the coefficient!) 3. Find the molar mass of the compound (don’t forget the water!) 4. Plug your answers into the following formula:
mass of water molar mass of hydrate Solve the following problems in the space provided. 1. Find the percent of water in zinc sulfate decahydrate. 2. Find the percent of water in cupric sulfate pentahydrate.
x 100 = % water
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Percent Composition
Solve the following problems on a separate sheet of paper. 1. Glucose (C6H12O6) is the most important nutrient in the cell for general and chemical
potential energy. a. What is the mass percent of each element in glucose? b. How many grams of carbon are found in 16.55 g of glucose?
2. Ammonium carbonate is a white crystalline powder that decomposes with warming. It has many uses, including being a component of baking powder, fire extinguishers, and smelling salts.
a. How many moles are equivalent to 41.6 grams of ammonium carbonate? b. What percent of the compound is composed of carbonate ions?
3. Ammonium nitrate is used to manufacture the important dental anesthetic dinitrogen monoxide, also called nitrous oxide and laughing gas.
a. Calculate the mass percent of nitrogen in ammonium nitrate. b. What mass of nitrogen is present in 35.8 mg of nitrous oxide?
4. Plaster of Paris, calcium sulfate half hydrate, is often used in making surgical casts. It is a hydrated compound in which two formula units of calcium sulfate are combined with one molecule of water.
a. How many moles of calcium are contained in 58.5 g of Plaster of Paris? b. What is the mass percent of water in the plaster?
5. Lead can be prepared from the galena (lead II sulfide) by roasting the mineral in oxygen gas. How much lead can be obtained from a metric ton (1ooo kg) of galena?
6. Zinc and sulfur react to form zinc sulfide. If we mix 19.5 grams of zinc and 9.40 grams of sulfur, have we added sufficient sulfur to fully react with all the zinc? Show evidence for your answer.
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Empirical Formulas
Once a new compound has been made in the laboratory, we can usually determine its percent composition experimentally. From the percent composition data, we can calculate the empirical formula of the compound. Empirical means based on experiment. Therefore, an empirical formula is one that is obtained from experimental data and represents the smallest whole-number ratio of atoms in a compound. It is also known as the simplest formula. STEPS: 1. Change % to grams. 2. Change grams to moles. 3. Pick smallest number of moles and divide all values by it. 4. If you get a decimal value, 0.5, multiply everything to make the numbers whole. **If the number ends in .5, multiply everything by 2. Or, simply follow the steps in the poem:
Percent to mass Mass to mole
Divide by smallest Multiply ‘til whole
EXAMPLE: Analysis shows a compound to contain 56.58% potassium, 8.68% carbon, and 34.73% oxygen. Find the empirical (simplest) formula of this compound. NAME this compound!
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Empirical Formulas
DIRECTIONS: Solve the following problems.
1. 43.64% phosphorus, 5.36% oxygen — calculate the empirical formula and name the compound.
2. 21.31% iron, 12.23% sulfur, 18.32% oxygen, 48.14% water — calculate the empirical formula and name the compound.
3. A binary compound that contains oxygen and arsenic is 75.7% arsenic by mass. What is the empirical formula?
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4. In an experiment, a 2.514-g sample of calcium was heated in a stream of pure oxygen, and was found to increase in mass by 1.004 g. Calculate the empirical formula of the compound and name it.
5. If cobalt metal is mixed with excess sulfur and heated strongly, a sulfide is produced that contains 55.06% cobalt by mass. Calculate the empirical formula and name the compound.
6. Analysis of a certain compound yielded the following percentages of the elements by mass: nitrogen, 29.16%; hydrogen, 8.392%; carbon, 12.50%; oxygen, 49.95%. Calculate the empirical formula and name the compound.
7. If 10.0 grams of an unknown carbon-nitrogen-hydrogen compound contains 17.7% nitrogen and 3.8 x 1023 atoms of hydrogen, what is the empirical formula of the compound?
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vs.
Objective: To determine the % sugar in a piece of gum (either juicy fruit or double bubble)
Materials: electronic balance, gum, yourself
Conclusion: 1. Write a detailed procedure of each step
you will take to achieve the objective of
the lab. At least three steps.
2. Construct a data table and calculations
table.
3. Below the tables, show all your work
for each calculation. Pay attention to
significant figures and units.
4. Report your % composition to Miss
Uhernik.
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KEYS TO A SUCCESSFUL LAB WRITE-UP
1. Use a ruler to make the tables!
2. Consult the, “Little Mole Lab,” write-up
for examples of procedure, data
tables, and calculations table. (YOUR
TABLES MUST LOOK EXACTLY like the
ones on the Little Mole Lab)
3. Show all of your work and label it
with units and descriptions.
4. Title your lab. For example, “Percent
Sugar in Bazooka Gum.”
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% Composition and Empirical Formula Review DIRECTIONS: Solve the following problems.
1. Calculate the percent composition of each element in calcium perchlorate. 2. Calculate the percent of water in nickel (II) chloride hexahydrate. 3. 1,6-diaminohexane is used to make nylon. What is the empirical formula of this
compound if it is 62.1% carbon, 13.8% hydrogen, and 24.1% nitrogen? 4. Analysis of a 1.34g sample is known to contain 0.365g Na, 0.221g N, and oxygen. What
the empirical formula and name of this compound?
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5. Calculate the percent composition of each compound: a. H2S b. (NH4)2C2O4 c. Mg(OH)2 d. Na3PO4
6. Using your answers from #5, calculate the number of grams of these elements:
a. sulfur in 3.54 g H2S b. nitrogen in 25.0 g (NH4)2C2O4 c. magnesium in 97.4 g Mg(OH)2 d. phosphorus in 804 g Na3PO4
7. Calculate the percent of water in potassium aluminum sulfate dodecahydrate,
KAl(SO4)2 · 12 H2O.
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The molecular formula of a compound is either the same as its experimentally determined empirical formula, or it is some whole-number multiple of it. In other words, an empirical formula may or may not be the same as a molecular formula. For example, dinitrogen tetrahydride (molecular formula, N2H4) has an empirical formula of NH2 because this is the simplest ratio of nitrogen to hydrogen. Its molecular formula is not the same as its empirical formula, but it is a whole-number multiple of it. However, for carbon dioxide, its empirical and molecular formulas are the same, CO2, because this is the simplest ratio of carbon to oxygen and the actual formula for the compound. We can determine the molecular formula of a compound if we are given 2 pieces of information:
1. The compound’s empirical formula. With this information, we can determine the empirical formula mass (efm).
2. The molecular formula mass (mfm).
A relationship exists between the empirical formula and the molecular formula of a compound: (empirical formula)x = molecular formula
where x is a whole-number multiple of the empirical formula.
Therefore, the formula masses have the same relationship: (empirical formula mass)x = molecular formula mass
To determine x: x = mfm = WHOLE NUMBER!! (If not, check your work)
efm EXAMPLE: When 10.0 g of white phosphorus is exposed to air, it will ignite instantaneously, consuming 12.9 g of oxygen. The molecular formula mass is actually 284 g/mol. What is the molecular formula? Name both compounds.
1. Write down all given information. (Determine the empirical formula if not given.) 2. Calculate the empirical formula mass from the empirical formula. 3. Solve for x using the equation x = mfm/efm. 4. Plug x into the equation (empirical formula)x = molecular formula to determine the molecular
formula.
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Molecular Formula Problems
1. A compound with the empirical formula CH was found by experiment to have a molar mass of approximately 78. What is the molecular formula of the compound?
2. Fructose is a very sweet natural sugar that is present in honey, fruits, and fruit juices. It has a molar
mass of 180 g/mol and a composition of 40.0 % carbon, 6.5% hydrogen, and 53.5% oxygen. What is its molecular formula?
3. A 10.0 g sample is analyzed and found to contain 9.41 g oxygen and 0.59 g hydrogen. The actual
molar mass of the compound is 34 g/mol. Calculate the compound’s empirical formula and molecular formula. Name the substance.
4. A substance whose formula is X2O is 93.1% X and 6.9% O. Identify the element X.
5. Aspirin is well-known as a pain reliever (analgesic) and as a fever reducer (antipyretic). It has a molar mass of 180.2 g/mol and a composition of 60.0% carbon, 4.48% hydrogen, and 35.5% oxygen. What is the molecular formula of aspirin?
6. A student weighs 18 grams of aluminum and needs twice as many magnesium atoms as she has of
aluminum. How many grams of magnesium does she need?
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Name ____________________________________________________ Period ___________________
CHEMICAL QUANTITIES - Vocabulary Review
DIRECTIONS: Select the term from the following list that best matches each description. Each term may only be used once.
a. percent composition b. Avogadro’s number c. empirical formula d. molar mass e. 1 mole f. hydrate g. anhydrous h. molecular formula i. catalyst j. precipitate
1. the percent by mass of each element in a compound ____________
2. the lowest whole-number ratio of the atoms of the elements in a compound ____________
3. a simple whole number multiple of an empirical formula ____________
4. the mass (in grams) of one mole of a compound ____________
5. the number of representative particles contained in one mole of a substance ____________
6. 6.02 x 1023 particles ____________
7. a compound that has a specific number of water molecules bound to each
formula unit ____________
8. a substance that speeds up a chemical reaction but is not considered part
of the balanced chemical equation ____________
9. a compound in which the bound water molecules have been removed ____________
10. a solid that forms when two or more aqueous solutions are mixed ____________
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The Mole Review Solve the following problems. 1. Penicillin, the first of a now large number of antibiotics, was discovered accidentally by the Scottish
bacteriologist Alexander Fleming in 1928, but he was never able to isolate it as a pure compound. This and similar antibiotics have saved millions of lives that otherwise would have been lost to infections. Penicillin has the following composition:
53.82% C, 6.47% H, 8.97% N, 10.26% S, & 20.48% O.
a. What is penicillin's empirical formula?
b. If the antibiotic's molecular formula mass is 312 g/mol, what is its molecular formula?
c. How many carbon atoms are contained in 15 g of the substance? 2. Epsom salts, first isolated from mineral springs at Epsom in England in the seventeenth century, were
once used as a mild laxative. a. Calculate the empirical formula of the compound if a 25.2-g sample decomposes to produce
magnesium sulfate and 12.9 g of water.
b. Name the hydrate. _________________________________________________________________
c. How many molecules of water are contained in the 25.2-g sample?
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3. Combustion analysis of 1.00 grams of the male sex hormone testosterone yields 2.90 g CO2 and 0.875 g H2O. What are the mass percents of carbon, hydrogen, and oxygen in testosterone?
4. The actual molar mass of a compound is 92 g/mol. Analysis of a sample of the compound indicates it
contains 0.606 g of nitrogen and 1.382 g of oxygen. What is the compound's molecular formula?
5. The formula of a compound containing only copper and sulfur can be determined by heating weighted quantities of the two elements in a crucible. The following experimental data was collected in the laboratory.
mass of crucible 19.732 g mass of crucible and Cu 27.304 g mass of crucible and product 29.214 g Using this data, determine the empirical formula of the compound. Name the ionic compound. 6. Green Paris is a copper containing ionic compound that is used as a fungicide and herbicide on
grapes. The compound is 34.99% Cu, 26.45% C, 3.3% H, and 35.26% O. What is the empirical formula of Green Paris? Give the two chemical names for the compound.
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Chemical Quantities Review
1. How many moles are contained in 1.5 x 1023
molecules of ammonia? _____________________________ 2. What is the molar mass of Ba(NO3)2? _____________________________ 3. How many moles of chloride ions are in 1.5 moles of calcium chloride? _____________________________ 4. Calculate the mass of 2.50 moles of iron II hydroxide. _____________________________ 5. How many sulfide ions are in one mole of aluminum sulfide? _____________________________ 6. What is the total number of oxygen atoms in one formula unit of Ba(NO3)2? _____________________________ 7. In the following pairs of elements, circle the one that contains more atoms.
a. 1 mole of calcium or 1 mole of zinc b. 10 g of lithium or 10 g of bromine
c. 5.0 g of Al or 0.50 moles of boron d. 1 x 1023 atoms of lead or 1 mole of chromium 8. What is the empirical formula of hydrogen peroxide? _____________________________ 9. The representative particle of a covalent compound is called a _____________________________
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10. Five grams of beryllium reacts with 9 grams of oxygen. Calculate the empirical formula of the compound.
_____________________________ 11. Calculate the percent composition of Mg(OH)2. _____________________________ 12. A compound is composed of 50.7 % carbon, 4.2 % hydrogen, and 45.1 % oxygen. Its molecular mass
is 142 grams. Calculate the empirical and molecular formulas of the compound. Empirical _____________________________ Molecular ____________________________ 13. What is the percentage of water in sodium sulfate decahydrate? _____________________________ 14. Circle the following formula(s) that is a molecular formula ONLY?
a. NH4OH b. Fe(C2H3O2)3 c. C6H12O6 d. Na2SO3
15. The seven diatomic elements are ________________________________________________________.
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Name ________________________________________ Period __________ DIRECTIONS: Use the clues on the back of this page to complete the crossword.
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ACROSS 5. A substance that speeds up a chemical
reaction without being permanently altered.
8. NH3 10. The representative particle for
dichlorine heptoxide. 11. A substance that is consumed in a
chemical reaction. 13. The prefix for “ten.” 15. A compound in which the bound
water molecules have been removed. 18. To balance a chemical equation,
______________are added to the equation. 23. Represents the simplest whole
number ratio of atoms in a compound. 24. Percent ______________is calculated in the
laboratory to determine how close experimental results are to the actual values.
DOWN 1. Natural gas. 2. Two different substances of equal
moles have an equal number of particles, but differ in ______________
3. A substance that is produced in a chemical reaction.
4. During chemical reactions, ______________ is not created or destroyed.
6. To convert from moles of a substance to particles, one must use ______________’s number.
7. An atom or group of atoms with a negative charge.
9. The ______________ is at the heart of chemistry.
10. The mass in grams of one mole of a substance.
12. An atom or group of atoms with a positive charge.
14. Reactions that absorb heat and cool the immediate environment.
16. Solid formed in a double replacement reaction.
17. The number of protons in the nucleus of an atom.
18. The percent ______________of a compound is the same, no matter what size of the sample.
19. The representative particle for magnesium chloride.
20. The common name for calcium carbonate.
21. Chemical reactions that give off heat. 22. The composition reaction between a
nonmetallic oxide and water produces an ______________.
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Name ____________________________________________________ Period ___________________
Measuring Mass: A Means of Counting
objectives: Measure masses of common compounds, objects, and minerals
Calculate moles and atoms from experimental masses
introduction: You can often measure how much of something you have by counting individual objects. For example, you
can count the number of pennies you have in your pocket or the number of pencils you have in your locket.
You learned in Chapter 10 that in chemistry there is a name for a number of atoms, ions, or molecules. One
mole of a substance is equal to 6.02 x 1023 atoms, ions, or molecules of that substance. You also learned
that you can “count” the number of moles in a substance by obtaining the mass of the substance.
purpose: In this experiment you will measure the masses of samples of various common compounds such as water,
salt, and sugar. You will use your results as a means of counting the atoms, ions, and molecules in your
samples. You will extend your technique to common objects that you can consider to be pure substances,
such as glass marbles, pieces of chalk, and polystyrene peanuts. Finally, you will measure the masses of
various mineral samples and use your results to find the number of atoms in each.
materials: sodium chloride polystyrene peanuts sucrose (C12H22O11)
sulfur glass slides fluorite (CaF2)
chalk hematite (Fe2O3) other common minerals
electronic balance
procedure: 1. Mass one level teaspoon or one piece of each substance in the data table. Record the masses in their
respective data tables, Table 13.1, 13.2, or 13.3.
2. Mass one level teaspoon of gypsum. Record in question #1 in the, “Now It’s Your Turn” section.
3. Mass a nickel coin. Record in question #2 in the, “Now It’s Your Turn,” section.
things to remember: 1. Do not contaminate! Please use only the designated plastic cup and spoon for each chemical. READ
LABELS!!
2. After weighing each substance, place the sample back into the proper container.
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Name ____________________________________________________________________ Period ________
Table 13.1 Counting Particles in Common Substances
Table 13.2 Counting Particles in Common Items
Formula
name mass in grams
molar mass
moles in 1 teaspoon
moles of each
element
atoms of each
element
SiO2
glass slide
CaCO3
polystyrene peanut
104 g/mol
(per unit molecule)
Table 13.3 Counting Particles in Minerals
Formula
name mass in grams
molar mass
moles in 1 teaspoon
moles of each
element
atoms of each
element
S
CaF2
fluorite
Fe2O3
or hematite
Formula
Name mass in grams
molar mass
moles in 1 teaspoon
moles of each
element
atoms of each
element
NaCl
H2O
C12H22O11
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Name ___________________________________________________________ Period ____________ Questions for Analyses 1. Calculate the number of moles of one level teaspoon of each substance in Tables 13.1, 13.2,
and 13.3. Show ALL WORK and place your answer in the appropriate place in the data table.
Example: 23.46 g MgCl2 1 mol MgCl2 0.2464 mol MgCl2
95.21 g MgCl2
x =
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2. Calculate the moles of each element contained in each substance in Tables 13.1, 13.2, and 13.3. Show ALL WORK and place your answer in the appropriate place in data table.
Example: 0.2464 mol MgCl2 1 mol Mg 0.2464 mol Mg
1 mol MgCl2
0.2464 mol MgCl2 2 mol Cl 0.4928 mol Cl
1 mol MgCl2
x =
x =
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3. Calculate the atoms of each element contained in each substance in Tables 13.1, 13.2, and 13.3. Show ALL WORK and place your answer in the appropriate place in the data table.
Example: 0.2464 mol Mg 6.02 x 10
23 atoms Mg 1.483 x 10
23 atoms Mg
1 mol Mg
0.4928 mol Cl 6.02 x 1023 atoms Cl 2.967 x 1023 atoms Cl 1 mol Cl
x =
x =
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4. In step 1, you measured equal volumes of three different compounds. Which of the three
compounds has the greatest number of moles in one teaspoon?
_________________________________________________________________________________
5. Which of the three compounds in step 1 has the greatest total number of atoms?
_________________________________________________________________________________
6. Why can we use the technique of measuring volume as a means of counting?
_________________________________________________________________________________
Now it’s Your Turn
1. A nickel coin is a mixture of metals called an alloy. It consists of 75 percent copper and 25 percent nickel. How many nickel atoms are in one 5-cent piece? (HINT: you will need to measure the mass of a nickel.)
2. A common mineral used in wallboard and plaster of Paris is gypsum, CaSO4 ∙ 2H2O. Gypsum is an example of a hydrate. A hydrate is a compound that has water molecules incorporated into its crystal structure. The chemical formula of gypsum indicates that there are two water molecules for every calcium and sulfate ion within the crystal structure of gypsum. These water molecules are called water of hydration. Determine the number of water molecules in a small sample of gypsum. (HINT: you will need to measure the mass of a sample of gypsum.)
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Determination of the Empirical Formula of a Hydrate
INTRODUCTION:
Many salts that have been crystallized from water
solutions appear to be perfectly dry, yet when heated
yield large quantities of water. The crystals change
form, and sometimes color, as the water is driven off.
This suggests that water was present as part of the
crystal structure. Such compounds are called
hydrates. A hydrate that has lost its water is called
anhydrous salt. For a hydrate, the number of moles
of water present per mole of salt is usually some
simple, whole number.
Because salts consist of cations and anions bonded
together (and also because all metals are cations and
all nonmetals are anions), an anhydrous salt if often
symbolized MN, where the M stands for "metal" and
the N stands for "nonmetal." Similarly, a hydrate -
which consists of an anhydrous salt and water - is
often symbolized MN • ?H2O, where the question
mark indicates the integer number of water
molecules for each formula unit of salt. The dot
between the MN and the ? H2O means that the
water molecules are rather loosely attached to the
anhydrous salt. When referring to an unknown
hydrate, chemists use the notation described above.
One example of a hydrate is copper (II) chloride
dihydrate. Its blue crystals look and feel dry, but each
mole of the anhydrous salt is actually bonded to two
moles of water. The compound's formula is CuCl2 •
2H2O. The molar mass of CuCl2 • 2H2O is:
63.5 g + 2(35.4 g) + [2(18.0 g)] = 170.3 g
If a 170.3 g sample of CuCl2 • 2H2O were heated to
drive off all the water, the anhydrous salt CuCl2
would weigh
63.5 g + 2(35.4 g) = 134.5 g,
which is the mass of one mole of CuCl2. The mass of
water that has been boiled off into the air is [2(18.0
g)] = 36.0 g, which is the mass of two moles of water.
The formula of the hydrate shows the ratio of the
moles of anhydrous salt to the moles of water; in the
above case, that ratio is 1:2.
In this experiment, you will be given a sample of a
hydrate. You will determine the mass of the water
driven off by heating, as well as the amount of
anhydrous salt that remains behind. Then, given the
mas of one mole of the anhydrous salt, you will
determine the empirical formula of the hydrate.
MATERIALS:
hydrate 150 mL beaker
wire gauze balance
Bunsen burner glass stirring rod
hot hands
ring stand with iron ring
PROCEDURE:
1. Set up a ring stand apparatus. Place the wire gauze on the iron ring.
2. Place a clean, dry 150 mL beaker on the wire
gauze.
3. With the Bunsen burner on low heat, warm the
beaker for two minutes.
4. Allow the beaker to cool for five minutes and
then use the "hot hands" to carry the beaker over
to a balance. Weigh the beaker and record this
mass in the Data Table.
5. Without using the tare button and while the
beaker is still on the balance, place 1 spoonful of
the unknown hydrate crystals into the beaker.
Record this mass in the Data Table.
6. Place the beaker back on the wire gauze and
heat with the Bunsen burner on low heat until all
of the blue color is gone.
7. Allow the beaker to cool for five minutes, then
use the "Hot hands" to carry it back to the same
balance you used before. Record this mass in the
Data Table.
8. Place the beaker back on the wire gauze and
reheat with the Bunsen burner for an additional
4 minutes on low heat.
9. Once again, allow the beaker to cool for five
minutes and then use the hot hands to carry it
back to the same balance you used before.
Record this value in the Data Table.
10. Dump the solid in the garbage can and clean the
beaker. Put all equipment neatly back where you
found it.
11. Obtain the molar mass of the anhydrous salt from
your teacher. Record this value in the Data Table.
All Dried Up!
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Pre-Laboratory Assignment Use the following information to answer the questions. Show work, include units, and put your answers in the blanks. Tess Tube weighs an empty beaker and finds it to have a mass of 95.85 g. After putting a spoonful of an unknown hydrate into the beaker, she finds that the mass has increased slightly to 99.87 g. The chemist heats the beaker and its contents twice, and finds that the mass has dropped to 97.22 g. Tess is told by her teacher that the molar mass of the anhydrous salt is 74.10 grams. 1. What mass of hydrate did Tess start with? _______________________ 2. How much water was driven off from the hydrate during the heating process in units of...
A. grams? _______________________ B. moles?
_______________________ 3. How much anhydrous salt remained in the beaker in units of...
A. grams?
_______________________ B. moles?
_______________________ 4. A. Write down the mole ratio as decimal numbers: _____ moles anhydrous salt : _______ moles water B. Write down the mole ratio as whole numbers: _____ moles anhydrous salt : _______ moles water 5. What is the formula of the hydrate? (use MN to symbolize the anhydrous salt) _______________________ 6. Based on Tess' data, calculate the percentage of water in the sample of hydrate.
_______________________ 7. Why must Tess heat the sample twice instead of just once?
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Name _____________________________________________________________ Period __________
All Dried Up!
Data Table Quantity Measured Mass
dry beaker
beaker and contents before heating
beaker and contents after first heating
beaker and contents after second heating
molar mass of anhydrous salt (from teacher)
Calculations Table Calculation Answer
1. mass of hydrate
2a. mass of water
2b. moles of water
3a. mass of anhydrous salt
3b. moles of anhydrous salt
4b. mole ratio of anhydrous salt : moles of water
5. formula of hydrate
6. percentage of water in hydrate
7. percent error
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Calculations: Show your work & write your answers in the blanks to the right & in the Calculations table. 1. What mass of hydrate did you start with? _______________________ 2. How much water was driven off from the hydrate during the heating process in units of...
A. grams? _______________________ B. moles?
_______________________ 3. How much anhydrous salt remained in the beaker in units of...
A. grams?
_______________________ B. moles?
_______________________ 4. A. Write down the mole ratio as decimal numbers: _____ moles anhydrous salt : _______ moles water B. Write down the mole ratio as whole numbers: _____ moles anhydrous salt : _______ moles water 5. What is the formula of the hydrate? (use MN to symbolize the anhydrous salt) _______________________ 6. Based on Tess' data, calculate the percentage of water in the sample of hydrate.
_______________________ 7. The actual percent of water in the hydrate is ________________________. Using the formula below,
calculate the percent error in your experiment by comparing the actual percentage of water with the percentage you obtained in your experiment. Show your work. % error = |actual – experimental| actual
_______________________
x 100
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Determination of the Empirical Formula of a Magnesium Oxide
The empirical formula of any compound can easily be calculated from its percentage composition or composition by mass. The mole ratio is obtained by dividing the mass or percentage of each element in the compound by its atomic weight. This ratio is converted to whole numbers which become the subscripts in the empirical formula of the compound. **This ratio, when converted to whole numbers, represents the subscripts in the empirical formula of the compound.** In this lab a known amount of magnesium metal will be heated in the presence of oxygen to form magnesium oxide. The weight of product will be determined. From the experimental data, the weight of oxygen that reacted can then be calculated. Using these known masses of reactants, the empirical formula of magnesium oxide can then be determined. OBJECTIVES: 1. To determine the composition of a compound
in terms of masses and mass percentages of the elements of which it is composed.
2. To determine the experimental and actual empirical formulas of a compound.
MATERIALS: magnesium ribbon nickel crucible and cover clay triangle ring stand Bunsen burner balance tongs glass stirring rod wire gauze metric ruler distilled water PROCEDURE: 1. Obtain a nickel crucible and cover from your
teacher. Set up the crucible support apparatus. 2. To clean the crucible and cover, heat them for
3 - 4 minutes. This initial heating is necessary to drive of any volatile materials from the crucible. From this point on, handle the crucible with the crucible tongs only. Hot crucibles look the same as cold ones. Don't burn your fingers!!
3. Remove the crucible from the ring stand and let cool on the wire gauze. Proceed with the next step while waiting allowing it to it to cool for 5 minutes.
4. Cut a strip of magnesium ribbon approximately 35 cm long. Sand the ribbon with a piece of sandpaper.
5. Wind the Mg ribbon into a hollow ball that will fit in the bottom of the crucible.
6. Weigh the cooled crucible and cover. Record in data table.
7. Place the ball of Mg in the crucible; weigh the crucible, cover, and ribbon.
8. Place the crucible (no cover!) and contents on the clay triangle.
9. For optimum results, it is necessary that the magnesium burn very slowly, and that the finely divided, white magnesium oxide smoke be kept from escaping. This can be accomplished by heating the bottom of the crucible and lifting the cover only momentarily to allow a fresh supply of air to enter the combustion chamber. Begin by holding the crucible cover with the tongs and heating the bottom of the crucible rather strongly until the magnesium ignites. At this moment, place the cover on the crucible and turn down the flame a bit. After a short interval, lift the cover and allow enough air to enter to again ignite the Mg. A puff of white smoke is sufficient evidence that the Mg has ignited. Immediately replace the cover.
10. Repeat this procedure until the magnesium no longer ignites when the cover is raised.
11. At this point, adjust the cover so that there is a small gap to allow a steady flow of air into the crucible. Heat with a hot flame for three minutes. Allow to cool.
12. Pulverize the crucible contents with a glass stirring rod. Be sure that no powder sticks to the rod. Add 5 to 10 drops of distilled water to the product with your dropper; replace the cover and heat gently for 3 - 4 minutes. Carefully note any odor coming from the crucible.
13. Turn up the flame and heat strongly for another 3 - 4 minutes. This procedure will convert a chemical by-product, magnesium nitride, to magnesium oxide.
14. Allow the crucible, cover, and contents to cool. 15. Weigh the crucible, cover, and contents.
Record in the data table. 16. Clean the crucible and cover as well as possible
by wiping it out with a dry paper towel. Do not use water to clean the crucible. Return the clean crucible and cover to the central distribution table.
You Burn Me Up!!
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NAMES __________________________________________________ PERIOD ____________
CONCLUSION:1. On the back of this paper, perform the calculations necessary to determine the empirical formula for
magnesium oxide. Show your work, label all numbers, and follow significant figure rules. No credit will be given for an answer that is not supported by work. Place your answers in the Calculations Table.
2. Write a paragraph explaining the main source of error in the procedure and how it affects your results.
Use complete sentences, proper grammar and correct spelling. Make sure the paragraph is neat, organized, and understandable to a nonchemist.
Mass of crucible and cover
Mass of crucible, cover, and Mg ribbon
Mass of crucible, cover, and product
Calculations Table
1. Mass of magnesium
2. Mass of product, magnesium oxide
3. Mass of oxygen in the product
4. Experimental Empirical Formula of magnesium oxide
5. Experimental mole ratio as a decimal value
6. Actual Empirical Formula of magnesium oxide
7. Actual mole ratio as a decimal value
8. Percent error of mole ratio
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1. Calculate the mass of magnesium consumed in the reaction.
2. Calculate the mass of product formed in the reaction.
3. Calculate the mass of oxygen consumed in the reaction.
4. Using the masses of magnesium and oxygen consumed in the reaction, follow the poem to calculate the experimental empirical formula for magnesium oxide. When calculating the moles of magnesium and oxygen, report the answers to four decimal places.
5. Convert the experimental mole ratio of Mg to O into a decimal value. Example: Mg5O3 The mole ratio is 5:3 or 5/3, which as a decimal value is1.67.
6. Determine the actual empirical formula of magnesium oxide by following the rules of writing formulas.
7. Convert the actual mole ratio into a decimal value.
8. Using the decimal values from #5 and #7, calculate the percent error of the mole ratio.
% error = actual – experimental x 100
actual