1

Unit 7:

Chemical

Quantities (Chapter 10)

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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 10

23 _________________ 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

1. 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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The MOLE is at the

Heart

of Chemistry

gram mole particle

periodic table 6.02 x 1023

- atoms - ions - formula units - molecules

1

rep.

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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 = 32.0 g/mo l Molar Mass = ___________ Molar Mass = ___________

Mass = 32.0 g 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 = 17.0 g/mol 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

1. Using the factor-label method, determine the mass of the following elements.

a. 3.00 moles of helium

b. 7.50 moles of boron

2. Using the factor-label method, determine the number of moles in each of the following

a. 25.0 grams of Li

b. 50.0 grams of Na

c. 8.50 grams of oxygen gas, which exists not as an atom but as a diatomic molecule, O2.

3. Using the factor-label method, determine the number of atoms in each of the following.

a. 2.00 moles of iron

b. 27 g of Be

4. 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?

5. 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.

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6. Complete these statements by supplying the correct 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.

7. Ascorbic acid, more commonly known as vitamin C (C6H8O6) is an essential vitamin. Since it cannot be

stored in the body, it must be present in the diet or taken as a dietary supplement. If a typical tablet

contains 500.0 mg of Vitamin C, how many molecules of the vitamin does it contain?

8. When dissolved in water, aspartame, marketed as NutraSweet, is 160 times sweeter than sucrose (table

sugar). The molecular formula for this artificial sweetener is C14H18N2O5.

a. How many moles of aspartame are contained in 10.0 grams of the sweetener?

b. How many atoms of nitrogen are in 1.2 g of aspartame?

9. Dimethylnitrosomine (C2H6N2O) 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 1014

molecules of dimethylnitrosomine?

10. Chloral hydrate (C2H3Cl3O2) is used as a sedative and a hypnotic. It is the compound used to make “Mickey Finns” in detective stories. How many moles of chloral hydrate are contained in 5.0 g of the drug?

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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 a separate sheet.

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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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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The Chemistry Mole Pre-Quiz

DIRECTIONS: Answer the following questions. They are very similar to the questions you will see on the quiz. The questions require a math problem. The others do not. 1. What is the total number of ions in one formula unit of ammonium sulfate? ____________

2. One mole of diatomic oxygen has a mass of ____________.

3. What is the mass of 0.5 moles of carbon dioxide? ____________

4. The representative particle of an ionic compound is called a ____________________________.

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5. How many moles of carbon can be obtained from 3 moles of aluminum carbonate? _______ 6. How many moles of nitrite ions are in one mole of barium nitrite? ____________

7. The molar mass of barium chloride is ____________.

8. How many atoms are in 52.00 g of chromium? ____________

9. How many grams is equivalent to 3.91 x 1024 molecules of diatomic bromine? ____________

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Name __________________________________________________________________ Period ______ OBJECTIVE: After writing your FULL name 15 times on the chalkboard, determine:

a. the mass of chalk on the chalk board b. the number of moles of chalk on the chalkboard c. the number of formula units of chalk on the chalkboard d. the number of atoms of oxygen on the chalkboard

MATERIALS: chalk (_______________________________), chalkboard, and an electronic balance Procedure:

1.

2.

3. conclusion: On the back of this paper, construct a data table and a calculations table. Below the tables, show all your work for each of the four calculations. Pay attention to significant figures and units.

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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 DIRECTIONS: Write the formula for the compound on the space provided. Then, calculate the percent composition of each element in the following compounds.

1. plumbic chloride __________________________

2. aluminum sulfite __________________________

3. barium nitrate __________________________

4. ammonium phosphate __________________________

5. Magnesium hydroxide is 54.87% oxygen by mass.

a. How many grams of oxygen would be contained in 175 g of the compound?

b. How many moles of oxygen does this represent?

6. An 8.20 g sample of magnesium combines completely with 5.40 g of oxygen to form a compound. What is the percent composition of the elements in this compound?

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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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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. 63.50% silver, 8.25% nitrogen, 28.26% 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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Empirical Formulas Continued

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.

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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. 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 its molecular formula?

4. A compound composed of hydrogen and oxygen is analyzed and a 10.00 g sample of the

compound yields 0.59 g of hydrogen. The actual molar mass of this compound is 34 g. Find the empirical formula and the molecular formula for this compound. Name the compound.

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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. Calculate the % phosphorus in each of the following substances.

a. sodium phosphate b. tetraphosphorus decoxide c. phosphorus trihydride d. calcium phosphite Arrange the above substances in order of increasing % phosphorus. ____________________________________________________________________________________________ 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. 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 ___________________

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 =

40

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 =

41

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 =

42

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.)

43

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!

44

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?

45

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

46

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

47

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!!

48

NAMES __________________________________________________ PERIOD ____________ You Burn Me Up!!

Data Table Mass of crucible and cover

Mass of crucible, cover, and Mg ribbon

Mass of crucible, cover, and product

CALCULATIONS: SHOW ALL WORK! Report all answers in the Calculations Table found on the next page. 1. Write the balanced chemical equation for the reaction that occurred in this lab.

________________________________________________________________________

2. Calculate the mass of magnesium used in the reaction. 3. Calculate the mass of product formed in the reaction. 4. Calculate the mass of oxygen consumed in the reaction. 5. Using the masses for magnesium and oxygen that you calculated, follow the poem to calculate the

empirical formula for magnesium oxide. NOTE: When determining the moles of magnesium and oxygen, calculate the answer to four decimal places.

49

6. What is the experimental empirical formula that you found in #5: ____________________

a. Based on this empirical formula what is the experimental mole ratio of Mg to O?

___________ mol Mg: ____________ mol O

7. What is the actual empirical formula for magnesium oxide that you found in #1. (The actual empirical formula for magnesium oxide can be found by following the rules of naming and writing formulas—the “criss-cross” method).

_________________________________

8. Based on the actual empirical formula for magnesium oxide that you found in #7, what is the actual

mole ratio of Mg to O?

___________ mol Mg: ____________ mol O

9. Ratio’s can also be written as fractions. For example, a 1:2 ratio can be written as ½ = 0.5. Re-write the actual and experimental mole ratios as fractions then decimal values. Actual mole ratio - ______________________ Experimental mole ratio - ______________________

10. Using your decimal values from #9, calculate the percent error of the mole ratios.

% error = actual – experimental x 100 actual

Calculations Table

1. Mass of magnesium

2. Mass of product, magnesium oxide

3. Mass of oxygen

4. Experimental Empirical Formula of magnesium oxide

5. Actual Empirical Formula of magnesium oxide

6. Percent error of mole ratio